Atomic Molecular Physics

# Atomic structure and bonding This study guide summarizes the fundamental principles of atomic structure and chemical bonding, covering the subatomic particles of atoms, quantum mechanical descriptions of electrons, various types of chemical bonds, and theories explaining their formation. ## 1. Atomic structure ### 1.1 Components of an atom An atom consists of a central nucleus and surrounding electrons [1](#page=1). * **Nucleus:** Contains positively charged protons and neutral neutrons. The nucleus has an overall positive charge determined by the number of protons [1](#page=1). * **Protons:** Positively charged (+1) [1](#page=1). * **Neutrons:** Neutral [1](#page=1). * Protons and neutrons have nearly equal masses [1](#page=1). * **Electrons:** Negatively charged (-1) and have a mass approximately 1/2000th of a proton's mass, contributing minimally to the atomic mass [1](#page=1). ### 1.2 Atomic number and mass number * **Atomic number (Z):** The number of protons in an atom's nucleus. This number determines the element and the number of electrons in a neutral atom [1](#page=1). * **Atomic mass number (A):** The total number of protons and neutrons in the nucleus [1](#page=1). ### 1.3 Atomic mass unit (Da or U) The unified atomic mass unit (Da or U) is defined as 1/12th of the mass of a Carbon-12 atom. It is approximately equal to the mass of a proton or neutron (around $1.66 \times 10^{-27}$ kg) [1](#page=1). * Hydrogen-1 ($^1$H): Approximately 1.008 Da [1](#page=1). * Carbon-12 ($^{12}$C): Exactly 12.00 Da [1](#page=1). * Oxygen-16 ($^{16}$O): Approximately 16.00 Da [1](#page=1). * A typical protein weighs around 30 kDa, which is equivalent to 30,000 Da or $30,000 \times 1.66 \times 10^{-23}$ kg [1](#page=1). ### 1.4 The periodic table * Historically organized by atomic mass, it is now ordered by atomic number (Z) [1](#page=1). * **Groups (Columns):** Elements in the same group have the same number of valence electrons and thus similar chemical properties [1](#page=1). * **Periods (Rows):** Elements in the same period have the same number of electron shells [1](#page=1). ### 1.5 Atomic properties * **Ionization energy:** The energy required to eject an electron from a neutral atom [1](#page=1). * **Electron affinity:** The energy change when an electron is added to a neutral atom [1](#page=1). * **Atomic radius:** The distance from the center of the nucleus to the outermost electron shell [1](#page=1). * **Electronegativity:** The ability of an atom to attract electrons to itself in a chemical bond. It is often estimated as the average of ionization energy and electron affinity. Electronegativity differences help determine bond type [1](#page=1). ### 1.6 Isotopes Isotopes are atoms of the same element (same atomic number) with different numbers of neutrons, resulting in different atomic mass numbers. They exhibit similar chemistry due to having the same number of electrons. The periodic table lists the average atomic mass, which is a weighted average of the isotopic masses [1](#page=1). ### 1.7 Atomic models * **Bohr model:** Describes the hydrogen atom with electrons in specific orbits around the nucleus, but is limited for atoms with more than one electron [1](#page=1). * **Quantum mechanical model:** Treats electrons as having wave-like properties and describes their behavior using the Schrödinger wave equation. Solutions to this equation are wave functions ($\psi$), which, when squared ($\psi^2$), represent the probability density of finding an electron in a particular region around the atom [1](#page=1). * **Atomic orbital:** A region in an atom where there is a 90% probability of finding an electron [1](#page=1). ### 1.8 Quantum numbers Atomic orbitals are described by four quantum numbers that define the properties of an electron within an atom: * **Principal quantum number ($n$):** Determines the size and energy level of the orbital. Larger $n$ values correspond to larger orbitals further from the nucleus with higher energy. $n$ can be any positive integer (1, 2, 3,...) [1](#page=1). * **Orbital (azimuthal) quantum number ($l$):** Determines the shape of the orbital and the subshell. The values of $l$ range from 0 to $n-1$ [1](#page=1). * $l=0$ corresponds to an 's' subshell (spherical orbitals) [1](#page=1). * $l=1$ corresponds to a 'p' subshell (dumbbell-shaped orbitals) [1](#page=1). * $l=2$ corresponds to a 'd' subshell (more complex shapes) [1](#page=1). * $l=3$ corresponds to an 'f' subshell (even more complex shapes) [1](#page=1). * **Magnetic quantum number ($m_l$):** Specifies the orientation of the orbital in space. The values of $m_l$ range from $-l$ to $+l$, including 0 [1](#page=1). * For $l=0$ (s subshell), $m_l=0$ (1 orbital) [1](#page=1). * For $l=1$ (p subshell), $m_l=-1, 0, +1$ (3 orbitals) [1](#page=1). * For $l=2$ (d subshell), $m_l=-2, -1, 0, +1, +2$ (5 orbitals) [1](#page=1). * **Degenerate orbitals:** Orbitals within the same subshell but with different orientations have the same energy (e.g., $p_x$, $p_y$, $p_z$) [1](#page=1). * **Spin quantum number ($m_s$):** Describes the intrinsic angular momentum of an electron, which can be visualized as spin. An electron can have one of two spin states, usually represented as spin up ($+1/2$) or spin down ($-1/2$) [1](#page=1). ### 1.9 Pauli Exclusion Principle and Orbital Occupancy * **Pauli Exclusion Principle:** No two electrons in the same atom can have the same set of four quantum numbers. This means an atomic orbital can hold a maximum of two electrons, and they must have opposite spins [1](#page=1). * **Aufbau Principle:** Electrons fill atomic orbitals of lower energy first before occupying higher energy orbitals [1](#page=1). * **Hund's Rule:** Within a subshell, electrons will individually occupy degenerate orbitals before pairing up in any one orbital. Unpaired electrons in degenerate orbitals will have the same spin [1](#page=1). ## 2. Chemical Bonding ### 2.1 Types of chemical bonds Chemical bonds form when atoms share or transfer electrons to achieve a more stable electron configuration (#page=1, page=2) [1](#page=1) [2](#page=2). * **Ionic Bonds:** Formed by the electrostatic attraction between oppositely charged ions, which are created when atoms with a large electronegativity difference (typically > 2) transfer electrons (#page=1, page=2). The atom that loses electrons becomes a positively charged cation, and the atom that gains electrons becomes a negatively charged anion. Ionic compounds form crystal lattices [1](#page=1) [2](#page=2). * Cations are generally smaller than their parent atoms because they have fewer electrons to repel each other [2](#page=2). * Anions are generally larger than their parent atoms because the increased electron-electron repulsion expands the electron cloud [2](#page=2). * **Covalent Bonds:** Formed by the sharing of electron pairs between atoms, typically non-metals (#page=1, page=2) [1](#page=1) [2](#page=2). * **Bonding pairs:** Shared electrons [2](#page=2). * **Lone pairs:** Unshared electrons [2](#page=2). * When atoms have similar electronegativity, they share electrons more equally, forming nonpolar covalent bonds [2](#page=2). * If there's a difference in electronegativity, the shared electrons are pulled closer to the more electronegative atom, creating a polar covalent bond [2](#page=2). ### 2.2 Molecular Formulae and Representation * **Molecular Formula:** Lists the types and number of atoms in a molecule (e.g., $CO_2$, $C_6H_{12}O_6$) [2](#page=2). * **Empirical Formula:** Shows the simplest whole-number ratio of atoms in a molecule (e.g., for $C_6H_{12}O_6$, it is $CH_2O$) [2](#page=2). * **Structural Formula:** Shows the order of connectivity and arrangement of atoms in a molecule [2](#page=2). * **Skeletal Formula:** A shorthand for organic molecules where lines represent bonds and vertices/ends represent carbon atoms, with implicit hydrogens and heteroatoms shown explicitly [2](#page=2). ### 2.3 Valence Bond Theory and Hybridization Valence Bond Theory explains bonding by the overlap of atomic orbitals [2](#page=2). * **Overlap:** The sharing of electrons occurs in overlapping atomic orbitals. Greater overlap leads to stronger bonds [2](#page=2). * **Hybridization:** Atomic orbitals within an atom can combine to form new hybrid orbitals that are better suited for bonding, with specific shapes and orientations [2](#page=2). * **Sigma ($\sigma$) bonds:** Formed by head-to-head overlap of atomic orbitals along the internuclear axis [2](#page=2). * **Pi ($\pi$) bonds:** Formed by the sideways overlap of parallel p orbitals, occurring above and below the sigma bond axis. Pi bonds are weaker than sigma bonds and restrict rotation around the bond axis [2](#page=2). * **Examples of Hybridization:** * **$sp^3$ hybridization:** One s orbital and three p orbitals combine to form four identical $sp^3$ hybrid orbitals, arranged tetrahedrally. This is seen in methane ($CH_4$) [2](#page=2). * **$sp^2$ hybridization:** One s orbital and two p orbitals combine to form three $sp^2$ hybrid orbitals in a trigonal planar arrangement. The remaining unhybridized p orbital forms a $\pi$ bond. This is seen in molecules with double bonds, like ethene ($C_2H_4$) [2](#page=2). * **$sp$ hybridization:** One s orbital and one p orbital combine to form two $sp$ hybrid orbitals in a linear arrangement. The two remaining unhybridized p orbitals form two $\pi$ bonds. This is seen in molecules with triple bonds, like ethyne ($C_2H_2$) [2](#page=2). ### 2.4 Molecular Orbital Theory (MO Theory) An alternative to Valence Bond Theory, MO Theory combines atomic orbitals to form molecular orbitals that span the entire molecule. It can accurately predict bond lengths and energies but becomes complex for larger molecules [2](#page=2). ### 2.5 Fajans' Rules Fajans' rules predict the degree of ionic or covalent character in a bond based on the polarizing power of the cation and the polarizability of the anion [2](#page=2). * **Covalent character increases when:** * The cation is small and has a high charge [2](#page=2). * The anion is large and has a high charge [2](#page=2). * The cation has a high charge density [2](#page=2). ### 2.6 Bond Length and Strength As bond length increases, bond strength decreases. Bonds with more 's' character in hybridization are generally shorter and stronger than those with more 'p' character [2](#page=2). ## 3. Key Terminology * **Atom:** The smallest particle of an element that retains its chemical identity [3](#page=3). * **Element:** A substance consisting of atoms that all have the same number of protons [3](#page=3). * **Isotope:** Atoms of the same element with different numbers of neutrons [3](#page=3). * **Mole:** A unit representing a specific number of particles, equal to Avogadro's number ($6.022 \times 10^{23}$) [3](#page=3). * **Avogadro's Constant ($N_A$):** $6.022 \times 10^{23}$ $mol^{-1}$ [3](#page=3). * **Molar Mass:** The mass in grams of one mole of a substance, typically expressed in g/mol [3](#page=3). * **Molarity (Molar Concentration):** The amount of solute in moles per liter (or $dm^{-3}$) of solution [3](#page=3). * **Electronegativity:** An atom's intrinsic ability to attract electrons in a bond [3](#page=3). * **Dipole:** A separation of electric charge within a molecule [3](#page=3). * **Dipole Moment:** The magnitude of a dipole [3](#page=3). ## 4. Quantum Numbers and Atomic Orbitals Recap * **Principle quantum number ($n$):** Indicates the energy level and size of the electron shell [1](#page=1). * **Orbital quantum number ($l$):** Determines the shape of the subshell and can range from 0 to $n-1$ [1](#page=1). * For $n=3$, $l$ can be 0, 1, or 2, corresponding to 3s, 3p, and 3d subshells [4](#page=4). * **Magnetic quantum number ($m_l$):** Specifies the orientation of orbitals in space [1](#page=1). * **Spin quantum number ($m_s$):** Describes the intrinsic spin of an electron, which can be spin up or spin down [1](#page=1). ## 5. Bonding Examples * **Sodium Chloride (NaCl) formation:** Sodium (Na) has one valence electron, and Chlorine (Cl) has seven. Chlorine is more electronegative and can pull sodium's valence electron, forming a positive sodium ion ($Na^+$) and a negative chloride ion ($Cl^-$). The electrostatic attraction between these ions forms an ionic bond [4](#page=4). ## 6. Electron Configuration The arrangement of electrons in atomic orbitals follows specific rules. * **Aufbau Principle:** Fill lower energy orbitals first [1](#page=1). * **Hund's Rule:** Fill degenerate orbitals singly before pairing electrons [1](#page=1). * **Pauli Exclusion Principle:** No two electrons in an atom can have the same set of four quantum numbers (#page=1, page=3) [1](#page=1) [3](#page=3). * **Example (Nitrogen, Z=7):** Electron configuration is $1s^2 2s^2 2p^3$. The three electrons in the 2p orbitals occupy separate orbitals with parallel spins according to Hund's rule [4](#page=4). * **Shorthand Notation:** Uses the preceding noble gas to represent core electrons [4](#page=4). * Potassium (K, Z=19): $[Ar 4s^1$ [4](#page=4). * Iodine (I, Z=53): $[Kr 5s^2 4d^{10} 5p^5$ [4](#page=4). --- # Chemical calculations and acid-base chemistry This section covers fundamental chemical calculations and the principles of acid-base chemistry, including the pH scale and buffer systems. ### 2.1 Chemical calculations #### 2.1.1 Moles and Avogadro's Constant * **Mole (mol):** A unit used to measure the amount of a substance, representing a specific number of particles [1](#page=1) [2](#page=2) [3](#page=3). * **Avogadro's Number/Constant ($N_A$ or $L$):** Defined as $6.022 \times 10^{23}$ particles per mole ($mol^{-1}$) [1](#page=1) [2](#page=2) [3](#page=3). **Formula:** Amount (moles) = $\frac{\text{mass (grams)}}{\text{molar mass (g/mol)}}$ [1](#page=1) [2](#page=2) [3](#page=3). #### 2.1.2 Molar Mass * **Molar Mass:** The mass in grams of one mole of a substance. For a compound, it is calculated by summing the relative atomic masses of all its constituent elements [1](#page=1) [2](#page=2) [3](#page=3). * **Dalton (Da) / Unified Atomic Mass Unit (U):** A unit of mass equal to 1/12 the mass of an atom of Carbon-12. It is approximately the mass of a proton or neutron [1](#page=1) [2](#page=2) [3](#page=3). * **Relationship:** 1 gram per mole (g mol$^{-1}$) is equivalent to 1 Dalton (Da) [1](#page=1) [2](#page=2) [3](#page=3). #### 2.1.3 Concentration * **Solution:** A mixture where a solute is dissolved in a solvent [1](#page=1) [2](#page=2) [3](#page=3). * **Concentration:** Can be expressed in three ways: * **Molar Concentration (Molarity):** Amount of solute in moles per unit volume of solution [1](#page=1) [2](#page=2) [3](#page=3). * **Units:** mol dm$^{-3}$ (or M) [1](#page=1) [2](#page=2) [3](#page=3). * **Formula:** Molar Concentration = $\frac{\text{amount of solute (moles)}}{\text{volume of solution (litres or dm}^{-3}\text{)}}$ [1](#page=1) [2](#page=2) [3](#page=3). * A solution with one mole of a substance in one litre is a one molar (1 M) solution [1](#page=1) [2](#page=2) [3](#page=3). * **Mass Concentration:** Mass of solute per unit volume of solution [1](#page=1) [2](#page=2) [3](#page=3). * **Units:** g dm$^{-3}$ [1](#page=1) [2](#page=2) [3](#page=3). * **Formula:** Mass Concentration = $\frac{\text{mass of solute (grams)}}{\text{volume of solution (litres or dm}^{-3}\text{)}}$ [1](#page=1) [2](#page=2) [3](#page=3). * **Volume Concentration:** Volume of solute as a percentage of the total volume of the solution [1](#page=1) [2](#page=2) [3](#page=3). #### 2.1.4 Dilutions * **Dilution:** The process of reducing the concentration of a solute in a solution, usually by adding more solvent [1](#page=1) [2](#page=2) [3](#page=3). * **Formula for single dilution:** $C_1V_1 = C_2V_2$, where $C$ is concentration and $V$ is volume [1](#page=1) [2](#page=2) [3](#page=3). * $V_1$ = volume of stock solution needed [1](#page=1) [2](#page=2) [3](#page=3). * $V_2$ = final desired volume of the diluted solution [1](#page=1) [2](#page=2) [3](#page=3). * **Volume of diluent to add:** $V_2 - V_1$ [1](#page=1) [2](#page=2) [3](#page=3). * **Dilution Factor (DF):** A numerical representation of how many times the original solution has been diluted [1](#page=1) [2](#page=2) [3](#page=3). * DF = $\frac{\text{Stock Concentration}}{\text{Diluted Final Concentration}}$ [1](#page=1) [2](#page=2) [3](#page=3). * DF = $\frac{\text{Final Volume}}{\text{Volume of Stock}}$ [1](#page=1) [2](#page=2) [3](#page=3). * **Serial Dilution:** A rapid method to generate solutions of low concentrations by repeatedly diluting a stock solution [1](#page=1) [2](#page=2) [3](#page=3). > **Tip:** A dilution factor of 1:10 means the original solution has been diluted 10 times. For example, 1 part sample mixed with 9 parts solvent results in a 1:10 dilution ratio and a dilution factor of 10. ### 2.2 Acid-Base Chemistry #### 2.2.1 Acids, Bases, and Water * **Water Dissociation:** Water can dissociate into a hydrogen ion (H$^{+}$) and a hydroxide ion (OH$^{-}$). In reality, protons are not free but associate with water molecules to form hydronium ions (H$_3$O$^{+}$) [1](#page=1) [2](#page=2) [3](#page=3). * $2\text{H}_2\text{O} \rightleftharpoons \text{H}_3\text{O}^+ + \text{OH}^-$ [1](#page=1) [2](#page=2) [3](#page=3). * **Equilibrium Constant ($K_w$):** The ionic product of water is a measure of the dissociation of water [1](#page=1) [2](#page=2) [3](#page=3). * $K_w = [\text{H}^+][\text{OH}^-]$ [1](#page=1) [2](#page=2) [3](#page=3). * At 25°C, $K_w = 1.0 \times 10^{-14} \text{ M}^2$ [1](#page=1) [2](#page=2) [3](#page=3). * **Concentration of H$^{+}$ and OH$^{-}$:** The concentration of hydrogen and hydroxide ions in pure water can be found by taking the square root of $K_w$. * $[\text{H}^+] = [\text{OH}^-] = \sqrt{K_w} = 1.0 \times 10^{-7} \text{ M}$ [1](#page=1) [2](#page=2) [3](#page=3). #### 2.2.2 Brønsted-Lowry Theory * **Acid:** A proton (H$^{+}$) donor [1](#page=1) [2](#page=2) [3](#page=3). * **Base:** A proton (H$^{+}$) acceptor [1](#page=1) [2](#page=2) [3](#page=3). #### 2.2.3 Acidity and pKa * **Acid Dissociation Constant ($K_a$):** A quantitative measure of an acid's strength in solution, representing the equilibrium constant for its dissociation [1](#page=1) [2](#page=2) [3](#page=3). * For a generic acid HA: $\text{HA} \rightleftharpoons \text{H}^+ + \text{A}^-$ [1](#page=1) [2](#page=2) [3](#page=3). * $K_a = \frac{[\text{H}^+][\text{A}^-]}{[\text{HA}]}$ [1](#page=1) [2](#page=2) [3](#page=3). * **pKa:** The negative logarithm (base 10) of the $K_a$ value [1](#page=1) [2](#page=2) [3](#page=3). * $pK_a = -\log_{10}(K_a)$ [1](#page=1) [2](#page=2) [3](#page=3). * **Stronger acids** have higher $K_a$ values and lower (or negative) $pK_a$ values, indicating greater dissociation [1](#page=1) [2](#page=2) [3](#page=3). * **Weaker acids** have lower $K_a$ values and higher $pK_a$ values, indicating less dissociation [1](#page=1) [2](#page=2) [3](#page=3). #### 2.2.4 pH Scale * **pH:** A measure of the acidity or alkalinity of a solution, based on the concentration of hydrogen ions [1](#page=1) [2](#page=2) [3](#page=3). * **pH < 7:** Acidic (higher [H$^{+}$] than [OH$^{-}$]) [1](#page=1) [2](#page=2) [3](#page=3). * **pH = 7:** Neutral (equal [H$^{+}$] and [OH$^{-}$]) [1](#page=1) [2](#page=2) [3](#page=3). * **pH > 7:** Basic/Alkaline (higher [OH$^{-}$] than [H$^{+}$]) [1](#page=1) [2](#page=2) [3](#page=3). * **Logarithmic Scale:** A one-unit change in pH represents a tenfold change in [H$^{+}$] [1](#page=1) [2](#page=2) [3](#page=3). * **Formula:** $\text{pH} = -\log_{10}[\text{H}^+]$ [1](#page=1) [2](#page=2) [3](#page=3). * **Calculating pH of Strong Acids:** The molar concentration of the strong acid is directly used as the [H$^{+}$] concentration, as strong acids fully dissociate [1](#page=1) [2](#page=2) [3](#page=3). * Example: For a 0.1 M HCl solution, $[\text{H}^+] = 0.1$ M, so pH = $-\log_{10}(0.1) = 1$. * **Calculating pH of Weak Acids:** Use the $K_a$ value and assume the concentration of undissociated acid remains relatively constant at equilibrium [1](#page=1) [2](#page=2) [3](#page=3). * **Formula:** $[\text{H}^+] = (\text{K}_a \times [\text{HA}])^{0.5}$ [1](#page=1) [2](#page=2) [3](#page=3). * **pH of Strong Bases:** Calculate the hydroxide ion concentration [OH$^{-}$] from the molarity of the strong base, then use $K_w$ to find [H$^{+}$] and subsequently the pH [1](#page=1) [2](#page=2) [3](#page=3). * **Relationship:** $\text{pH} + \text{pOH} = 14$ [1](#page=1) [2](#page=2) [3](#page=3). * $\text{pOH} = -\log_{10}[\text{OH}^-]$ [1](#page=1) [2](#page=2) [3](#page=3). #### 2.2.5 Buffers and the Henderson-Hasselbalch Equation * **pH Buffer:** An aqueous solution containing a weak acid and its conjugate base, or a weak base and its conjugate acid, that resists changes in pH upon addition of small amounts of acid or base [1](#page=1) [2](#page=2) [3](#page=3). * **Buffering Capacity:** Weak acids and bases buffer most effectively within a pH range of approximately $\pm 1$ unit around their $pK_a$ value [1](#page=1) [2](#page=2) [3](#page=3). * **Preparation:** A buffer solution is typically prepared by mixing a solution of a weak acid with a solution of a salt of that same acid [1](#page=1) [2](#page=2) [3](#page=3). * **Henderson-Hasselbalch Equation:** Used to estimate the pH of a buffer solution [1](#page=1) [2](#page=2) [3](#page=3). * $\text{pH} = pK_a + \log_{10}\left(\frac{[\text{A}^-]}{[\text{HA}]}\right)$ where $[\text{A}^-]$ is the concentration of the conjugate base and $[\text{HA}]$ is the concentration of the weak acid [1](#page=1) [2](#page=2) [3](#page=3). * **Ionization State:** The equation helps determine the ionization state of a functional group based on the pH relative to its $pK_a$ [1](#page=1) [2](#page=2) [3](#page=3). * If $\text{pH} < pK_a$, the group is mostly protonated (acid form) [1](#page=1) [2](#page=2) [3](#page=3). * If $\text{pH} > pK_a$, the group is mostly deprotonated (base form) [1](#page=1) [2](#page=2) [3](#page=3). * If $\text{pH} = pK_a$, the group is 50% ionized [1](#page=1) [2](#page=2) [3](#page=3). --- # Biomolecules: Structure and function Biomolecules are the essential organic molecules that form the basis of life, categorized into four major classes: carbohydrates, lipids, nucleic acids, and proteins, each with unique structures and functions crucial for cellular processes [2](#page=2) [3](#page=3). ### 3.1 Carbohydrates Carbohydrates, commonly known as sugars, are essential for energy storage, fuel, metabolic processes, and structural support. Their general formula is $(CH_2O)_n$ [3](#page=3). #### 3.1.1 Monosaccharides Monosaccharides are the simplest sugars, serving as the building blocks for larger carbohydrates. They are classified into two main subclasses based on the position of their carbonyl group (C=O) [3](#page=3): * **Aldoses**: Contain an aldehyde functional group (R-CH=O) at the end of the carbon chain. Examples include glucose and glyceraldehyde [3](#page=3). * **Ketoses**: Contain a ketone functional group (R-C(=O)-R') within the carbon chain. Examples include fructose and dihydroxyacetone [3](#page=3). Monosaccharides can exist in both open-chain and cyclic forms, with cyclization occurring when a hydroxyl group attacks the carbonyl group, forming a hemiacetal (from aldoses) or hemiketal (from ketoses). This cyclization creates a new stereocenter at the anomeric carbon, leading to two possible configurations: alpha ($\alpha$) and beta ($\beta$) anomers, which differ in the orientation of the hydroxyl group relative to the ring structure. These cyclic forms are often represented using Haworth projections [3](#page=3). #### 3.1.2 Disaccharides Disaccharides are formed when two monosaccharides are linked together by a glycosidic bond through a condensation reaction, releasing water. The nomenclature of disaccharides specifies the monosaccharides involved and the type of glycosidic bond [3](#page=3). * **Glycosidic bond**: Formed between the anomeric carbon of one monosaccharide and a hydroxyl group of another. * **Sucrose**: A disaccharide where both anomeric carbons (C1) of glucose and fructose are involved in an $\alpha$,$\beta$-1,2 glycosidic bond [3](#page=3). * **Lactose**: Composed of glucose and galactose, linked by a $\beta$-1,4 glycosidic bond, with one free anomeric carbon [3](#page=3). #### 3.1.3 Polysaccharides Polysaccharides are long chains of covalently bonded monosaccharides. Their structure and properties are determined by the type of monosaccharide units and the glycosidic linkages [3](#page=3): * **Cellulose**: Formed by $\beta$-D-glucose units linked via $\beta$-1,4 glycosidic bonds, providing structural support in plants [3](#page=3). * **Starch and Glycogen**: Composed of $\alpha$-D-glucose units linked via $\alpha$-1,4 glycosidic bonds. Starch consists of amylose (a linear helix) and amylopectin (branched via $\alpha$-1,6 linkages). Glycogen is more highly branched than amylopectin and serves as an energy store in animals [3](#page=3). Carbohydrates also function in post-translational modifications, such as in the ABO blood groups [3](#page=3). ### 3.2 Lipids Lipids are a diverse group of molecules characterized by their insolubility in water, serving as energy stores, structural components, and signaling molecules [3](#page=3). #### 3.2.1 Fatty Acids Fatty acids are long hydrocarbon chains with a terminal carboxylic acid group. They are classified based on saturation [3](#page=3): * **Saturated fatty acids**: Contain only single bonds between carbon atoms [3](#page=3). * **Unsaturated fatty acids**: Contain one or more double bonds. The stereochemistry of these double bonds (cis or trans, often denoted as Z or E) significantly affects their properties and packing. Cis (Z) configuration introduces a kink in the chain, while trans (E) is straighter [3](#page=3). Fatty acids are further categorized by chain length: short-chain (1-5 carbons), medium-chain (6-12 carbons), long-chain (13-21 carbons), and very long-chain (22+ carbons). Saturation is often indicated by notation like "16:0" for a 16-carbon saturated fatty acid, or "20:2" for a 20-carbon fatty acid with two double bonds. The position of double bonds can be specified from either the methyl end (n-7) or the carboxyl end (C-9) [3](#page=3). #### 3.2.2 Triacylglycerols Triacylglycerols are the primary form of stored energy in animals. They consist of a glycerol molecule esterified to three fatty acids [3](#page=3). #### 3.2.3 Glycerophospholipids Glycerophospholipids are major components of cellular membranes. They have a glycerol backbone esterified to two fatty acid chains and a phosphate group, which is further linked to a head group (e.g., choline, serine, inositol). These molecules are amphipathic, possessing both hydrophobic tails and hydrophilic heads, allowing them to form lipid bilayers, vesicles, and micelles in aqueous environments [3](#page=3). #### 3.2.4 Sphingolipids Sphingolipids are a class of phospholipids that utilize a sphingoid backbone, derived from serine and acetyl-CoA. They can be classified based on their head groups [3](#page=3): * **Sphingomyelins**: Contain a phosphate linked to choline or ethanolamine and are found in myelin sheaths [3](#page=3). * **Cerebrosides**: Have a monosaccharide as a head group and are found in neuron membranes [3](#page=3). * **Gangliosides**: Possess an oligosaccharide head group, including sialic acid units, and are present in the brain [3](#page=3). Sphingolipids are crucial for stabilizing lipid membranes [3](#page=3). #### 3.2.5 Sterol Lipids Sterol lipids, such as cholesterol, are predominantly found in eukaryotes. They feature a rigid, hydrophobic ring system and a polar hydroxyl group, making them amphipathic. Cholesterol is a precursor for steroid hormones and plays a vital role in maintaining membrane fluidity and stability [3](#page=3). #### 3.2.6 Lipid Aggregation Due to their amphipathic nature, lipids spontaneously aggregate in aqueous solutions to minimize contact between their hydrophobic tails and water. This phenomenon, known as the hydrophobic effect, leads to the formation of structures like micelles (from single fatty acids) and lipid bilayers or vesicles (from phospholipids and sphingolipids). The hydrophobic interior of these structures dictates membrane permeability [3](#page=3). * **Melting Temperature ($T_m$)**: The temperature at which a lipid bilayer transitions from a solid to a liquid phase. Saturated fatty acids pack tightly, increasing $T_m$ and membrane rigidity. Unsaturated fatty acids, with their kinks, pack less tightly, decreasing $T_m$ and increasing membrane fluidity [3](#page=3). ### 3.3 Nucleotides Nucleotides are the building blocks of nucleic acids (DNA and RNA) and also play crucial roles in cellular metabolism (e.g., ATP, GTP), second messenger systems (e.g., cAMP), and as enzyme substrates. A nucleotide consists of a ribose sugar, a phosphate group, and a nitrogenous base [3](#page=3). In nucleic acids, nucleotides are linked via phosphodiester bonds between the phosphate group of one nucleotide and the 3' carbon of the sugar of another, forming a sugar-phosphate backbone. DNA typically exists as a double helix, stabilized by hydrogen bonds between complementary base pairs (A-T with two hydrogen bonds, G-C with three hydrogen bonds). The higher number of hydrogen bonds in G-C pairs contributes to a higher melting temperature of DNA. The DNA helix has a diameter of approximately 2 nm, with bases stacked about 0.34 nm apart, and a full turn occurring every 3.4 nm, typically accommodating 10 base pairs [3](#page=3). ### 3.4 Proteins Proteins are polymers of amino acids linked by peptide bonds formed between the carboxyl group of one amino acid and the amino group of another through a condensation reaction. The sequence of amino acids (primary structure) determines the protein's higher-order structures and ultimate function [3](#page=3). * **Primary Structure**: The linear sequence of amino acids [3](#page=3). * **Secondary Structure**: Local folding patterns, such as $\alpha$-helices and $\beta$-sheets, stabilized by hydrogen bonds between backbone atoms [3](#page=3). * **Tertiary Structure**: The overall three-dimensional shape of a single polypeptide chain, stabilized by various non-covalent interactions (hydrogen bonds, ionic interactions, van der Waals forces, hydrophobic interactions) and sometimes covalent disulfide bridges between cysteine residues [3](#page=3). * **Quaternary Structure**: The association of multiple polypeptide subunits to form a functional protein complex [3](#page=3). Protein folding is a complex process often aided by chaperone proteins and driven by the minimization of Gibbs free energy. The structure and function of proteins can be investigated using various bioanalytical techniques, including protein purification via chromatography (size exclusion, ion exchange, affinity) and electrophoresis (SDS-PAGE). Enzymes, which are biological catalysts, are typically proteins whose three-dimensional structure is critical for their function. They speed up reactions by lowering activation energy without being consumed. The kinetics of enzyme-catalyzed reactions are often described by the Michaelis-Menten equation, characterized by parameters like $V_{max}$ (maximum velocity) and $K_m$ (substrate concentration at half $V_{max}$). Enzyme activity is influenced by factors such as pH, temperature, and the presence of inhibitors or cofactors [3](#page=3). --- # Metabolic pathways: Glycolysis, Krebs cycle, and oxidative phosphorylation Metabolic pathways are a series of interconnected biochemical reactions that occur within living organisms to sustain life, primarily focusing on energy production and utilization. ## 4. Metabolic pathways: Glycolysis, Krebs cycle, and oxidative phosphorylation Metabolic pathways are crucial for generating the energy required for cellular functions. This section focuses on three key pathways: glycolysis, the Krebs cycle (also known as the citric acid cycle or TCA cycle), and oxidative phosphorylation. These pathways are central to cellular respiration, the process by which cells convert glucose and other fuel molecules into ATP, the cell's primary energy currency. ### 4.1 Glycolysis Glycolysis is the initial metabolic pathway that breaks down glucose into pyruvate, generating a small amount of ATP and electron carriers (NADH) in the process. This pathway occurs in the cytosol and does not require oxygen, making it an anaerobic process [2](#page=2). #### 4.1.1 Overview of Glycolysis Glycolysis involves a series of ten enzyme-catalyzed reactions that convert one molecule of glucose (a six-carbon sugar) into two molecules of pyruvate (a three-carbon molecule). The overall reaction is [2](#page=2): Glucose + 2 NAD$^+$ + 2 ADP + 2 Pi $\rightarrow$ 2 Pyruvate + 2 NADH + 2 H$^+$ + 2 ATP + 2 H$_2$O [2](#page=2). #### 4.1.2 Stages of Glycolysis Glycolysis can be divided into two main phases: 1. **Investment Phase:** This phase requires an input of energy to prepare glucose for cleavage. * **Reaction 1: Phosphorylation:** Glucose is phosphorylated to glucose-6-phosphate by hexokinase using ATP. This step traps glucose within the cell by making it negatively charged and also prevents it from diffusing out of the cell [2](#page=2). * Reaction: Glucose + ATP $\rightarrow$ Glucose-6-phosphate + ADP [2](#page=2). * **Reaction 2: Isomerization:** Glucose-6-phosphate is converted to fructose-6-phosphate by phosphoglucose isomerase. This step rearranges the molecule to prepare it for symmetrical cleavage [2](#page=2). * Reaction: Glucose-6-phosphate $\rightleftharpoons$ Fructose-6-phosphate [2](#page=2). * **Reaction 3: Phosphorylation:** Fructose-6-phosphate is phosphorylated to fructose-1,6-bisphosphate by phosphofructokinase again using ATP. This is a highly regulated, irreversible step [2](#page=2). * Reaction: Fructose-6-phosphate + ATP $\rightarrow$ Fructose-1,6-bisphosphate + ADP [2](#page=2). 2. **Payoff Phase:** This phase generates ATP and NADH through substrate-level phosphorylation. * **Reaction 4: Cleavage:** Fructose-1,6-bisphosphate is split into two three-carbon molecules: dihydroxyacetone phosphate and glyceraldehyde-3-phosphate by aldolase [2](#page=2). * Reaction: Fructose-1,6-bisphosphate $\rightleftharpoons$ Dihydroxyacetone phosphate + Glyceraldehyde-3-phosphate [2](#page=2). * **Reaction 5: Isomerization:** Dihydroxyacetone phosphate is converted into glyceraldehyde-3-phosphate by triosephosphate isomerase. This ensures both molecules entering the payoff phase are glyceraldehyde-3-phosphate [2](#page=2). * Reaction: Dihydroxyacetone phosphate $\rightleftharpoons$ Glyceraldehyde-3-phosphate [2](#page=2). * **Reactions 6-10:** These reactions convert glyceraldehyde-3-phosphate into pyruvate, producing ATP and NADH. Key reactions include: * Oxidation and phosphorylation of glyceraldehyde-3-phosphate to 1,3-bisphosphoglycerate, producing NADH [2](#page=2). * Substrate-level phosphorylation: 1,3-bisphosphoglycerate transfers a phosphate group to ADP, forming ATP, catalyzed by phosphoglycerate kinase [2](#page=2). * Phosphate group rearrangement by phosphoglycerate mutase [2](#page=2). * Dehydration of 2-phosphoglycerate to phosphoenolpyruvate by enolase [2](#page=2). * Final substrate-level phosphorylation: Phosphoenolpyruvate transfers a phosphate to ADP, forming ATP and pyruvate, catalyzed by pyruvate kinase. This is another irreversible step [2](#page=2). #### 4.1.3 Net Products of Glycolysis For each molecule of glucose, glycolysis yields: * 2 molecules of pyruvate [2](#page=2). * 2 molecules of ATP (net production: 4 ATP produced - 2 ATP consumed) [2](#page=2). * 2 molecules of NADH [2](#page=2). ### 4.2 Krebs Cycle (Citric Acid Cycle) The Krebs cycle is a central metabolic pathway that further oxidizes the products of glycolysis, acetyl-CoA, to generate electron carriers (NADH and FADH$_2$) and a small amount of ATP or GTP. It takes place in the mitochondrial matrix [2](#page=2). #### 4.2.1 Overview of the Krebs Cycle Before entering the Krebs cycle, pyruvate is converted to acetyl-CoA through oxidative decarboxylation, a process catalyzed by the pyruvate dehydrogenase complex [2](#page=2). **Pyruvate + NAD$^+$ + CoA $\rightarrow$ Acetyl-CoA + NADH + CO$_2$ + H$^+$** [2](#page=2). The Krebs cycle then begins with acetyl-CoA combining with oxaloacetate to form citrate. Through a series of reactions, citrate is oxidized, releasing carbon dioxide, generating ATP (or GTP), NADH, and FADH$_2$, and regenerating oxaloacetate to continue the cycle [2](#page=2). The overall stoichiometry for one turn of the Krebs cycle (starting from acetyl-CoA) is: Acetyl-CoA + 3 NAD$^+$ + FAD + ADP + Pi + 2 H$_2$O $\rightarrow$ CoA + 3 NADH + 3 H$^+$ + FADH$_2$ + ATP + 2 CO$_2$ [2](#page=2). **Key Steps:** 1. **Citrate Synthesis:** Acetyl-CoA (2 carbons) condenses with oxaloacetate (4 carbons) to form citrate (6 carbons) [2](#page=2). 2. **Isomerization:** Citrate is isomerized to isocitrate [2](#page=2). 3. **Oxidative Decarboxylation:** Isocitrate is oxidized and decarboxylated, producing $\alpha$-ketoglutarate, NADH, and CO$_2$ [2](#page=2). 4. **Oxidative Decarboxylation:** $\alpha$-ketoglutarate is oxidized and decarboxylated to succinyl-CoA, generating NADH and CO$_2$ [2](#page=2). 5. **Substrate-Level Phosphorylation:** The energy released from breaking the thioester bond in succinyl-CoA drives the synthesis of ATP (or GTP) from ADP (or GDP) and Pi, forming succinate [2](#page=2). 6. **Oxidation:** Succinate is oxidized to fumarate, producing FADH$_2$ [2](#page=2). 7. **Hydration:** Fumarate is hydrated to L-malate [2](#page=2). 8. **Oxidation:** L-malate is oxidized to regenerate oxaloacetate, producing NADH [2](#page=2). #### 4.2.1 Electron Carriers The Krebs cycle produces significant amounts of NADH and FADH$_2$, which are crucial for the next stage of energy production. These molecules are reduced electron carriers that will donate electrons to the electron transport chain [2](#page=2). ### 4.3 Oxidative Phosphorylation Oxidative phosphorylation is the primary process for ATP synthesis during aerobic respiration, occurring on the inner mitochondrial membrane. It involves two coupled processes: the electron transport chain (ETC) and chemiosmosis [2](#page=2). #### 4.3.1 Electron Transport Chain (ETC) The ETC is a series of protein complexes embedded in the inner mitochondrial membrane that accept electrons from NADH and FADH$_2$. As electrons move through the chain, they are passed sequentially to electron acceptors with progressively higher reduction potentials [2](#page=2). * **Complexes:** The ETC consists of four main enzyme complexes (Complex I-IV) and mobile electron carriers like ubiquinone (Coenzyme Q) and cytochrome c [2](#page=2). * **Complex I (NADH-Q oxidoreductase):** Oxidizes NADH, transfers electrons to ubiquinone, and pumps protons from the mitochondrial matrix to the intermembrane space [2](#page=2). * **Complex II (Succinate-Q reductase):** Oxidizes FADH$_2$ (which enters the ETC at this complex), transfers electrons to ubiquinone [2](#page=2). * **Complex III (Q-cytochrome c oxidoreductase):** Receives electrons from ubiquinone, transfers them to cytochrome c, and pumps protons across the membrane [2](#page=2). * **Complex IV (Cytochrome c oxidase):** Receives electrons from cytochrome c and transfers them to oxygen, the final electron acceptor, reducing it to water. This complex also pumps protons [2](#page=2). * **Proton Pumping:** The energy released from electron transport is used to pump protons (H$^+$) from the mitochondrial matrix into the intermembrane space, creating an electrochemical gradient known as the proton-motive force [2](#page=2). #### 4.3.2 Chemiosmosis and ATP Synthesis The proton gradient generated by the ETC represents a form of stored potential energy. This energy is harnessed by ATP synthase, a large enzyme complex also embedded in the inner mitochondrial membrane [2](#page=2). * **ATP Synthase:** This enzyme acts as a molecular motor. Protons flow back down their concentration gradient from the intermembrane space into the matrix through a channel in the F$_0$ component of ATP synthase. This flow of protons drives the rotation of a central stalk, which in turn causes conformational changes in the F$_1$ catalytic component [2](#page=2). * **ATP Production:** The conformational changes in the F$_1$ component facilitate the phosphorylation of ADP to ATP, using the energy from the proton flow [2](#page=2). #### 4.3.3 ATP Yield The number of ATP molecules produced per molecule of glucose varies depending on the shuttle system used to transport NADH from glycolysis into the mitochondria. * Each NADH entering the ETC from the mitochondrial matrix yields approximately 2.5 ATP [2](#page=2). * Each FADH$_2$ yields approximately 1.5 ATP [2](#page=2). In total, aerobic respiration can yield a maximum of about 30-32 ATP molecules per molecule of glucose, with the majority produced through oxidative phosphorylation [2](#page=2). ### 4.4 Regulation of Metabolic Pathways These pathways are tightly regulated to meet the cell's energy demands and maintain homeostasis. Regulation often occurs at key irreversible steps, involving allosteric effectors (like ATP, ADP, AMP, citrate, and NADH), covalent modification of enzymes, and hormonal control (e.g., insulin and glucagon). The coordinated regulation of glycolysis and gluconeogenesis ensures that glucose is either produced or consumed appropriately to maintain blood glucose levels [2](#page=2). --- # Enzyme kinetics and regulation Enzyme kinetics and regulation are fundamental to understanding how biological catalysts function and how their activity is controlled within living organisms. ### 5.1 Enzyme kinetics Enzyme kinetics studies the rates of enzyme-catalyzed reactions and the factors that influence them. This field is crucial for understanding drug development, as many drugs target specific enzymes [3](#page=3). #### 5.1.1 Enzyme activity and substrate concentration * The rate of an enzyme-catalyzed reaction is influenced by substrate concentration ([S]) [3](#page=3). * At low substrate concentrations, the reaction rate is directly proportional to the substrate concentration, exhibiting **first-order kinetics** [3](#page=3). * At high substrate concentrations, the enzyme becomes saturated with substrate, and the reaction rate reaches a plateau, becoming independent of further increases in substrate concentration. This is known as **zero-order kinetics** [3](#page=3). * The maximum velocity of a reaction, denoted as $V_{max}$, is reached when the enzyme is saturated with substrate [3](#page=3). #### 5.1.2 Michaelis-Menten kinetics The Michaelis-Menten equation describes the relationship between the initial reaction velocity ($v$) and substrate concentration ([S]) [3](#page=3): $$ v = \frac{V_{max}[S]}{K_m + [S]} $$ * $V_{max}$ represents the maximum reaction velocity, achieved when the enzyme is fully saturated with substrate [3](#page=3). * $K_m$ (Michaelis constant) is the substrate concentration at which the reaction rate is half of $V_{max}$ [3](#page=3). * A low $K_m$ indicates that the enzyme has a high affinity for its substrate, reaching half-maximal velocity at a low substrate concentration [3](#page=3). * A high $K_m$ indicates a low affinity, requiring a higher substrate concentration to reach half-maximal velocity [3](#page=3). * $K_m$ is unique for each enzyme and is independent of enzyme concentration [3](#page=3). #### 5.1.3 Lineweaver-Burk equation (Double reciprocal plot) The Lineweaver-Burk equation is a linearized form of the Michaelis-Menten equation, useful for graphically determining $V_{max}$ and $K_m$. It plots $1/v$ against $1/[S]$ [3](#page=3): $$ \frac{1}{v} = \frac{K_m}{V_{max}} \left( \frac{1}{[S]} \right) + \frac{1}{V_{max}} $$ * The y-intercept of the Lineweaver-Burk plot is $1/V_{max}$. * The x-intercept is $-1/K_m$. * The slope of the line is $K_m/V_{max}$. #### 5.1.4 Turnover number ($k_{cat}$) * $k_{cat}$, also known as the turnover number, represents the maximum number of substrate molecules an enzyme can convert into product per second when the enzyme is working at its maximum speed (saturated with substrate) [3](#page=3). * It is a measure of the catalytic efficiency of an enzyme [3](#page=3). * **Catalytic efficiency** is often expressed as the ratio $k_{cat}/K_m$, providing a measure of how efficiently an enzyme converts substrate to product. A higher $k_{cat}/K_m$ value indicates greater efficiency [3](#page=3). ### 5.2 Factors affecting enzyme activity Several environmental factors can significantly influence the rate of enzyme-catalyzed reactions. #### 5.2.1 Effect of pH * Enzyme activity is highly sensitive to pH [3](#page=3). * Each enzyme has an **optimum pH** at which its activity is maximal [3](#page=3). * Deviations from the optimum pH can alter the ionization state of amino acid residues in the active site or affect the overall enzyme structure, leading to reduced activity [3](#page=3). * Extreme pH values can cause irreversible denaturation of the enzyme [3](#page=3). #### 5.2.2 Effect of temperature * Like pH, temperature has an optimal range for enzyme activity [3](#page=3). * Increasing temperature generally increases reaction rates due to increased kinetic energy of molecules [3](#page=3). * However, beyond the optimum temperature, enzymes become **denatured**. This occurs as excessive heat causes the enzyme molecule to vibrate so rapidly that non-covalent bonds are broken, disrupting its secondary and tertiary structure, rendering it inactive [3](#page=3). ### 5.3 Enzyme regulation Enzyme activity is tightly controlled through various mechanisms to meet cellular needs and respond to environmental changes. #### 5.3.1 Allosteric regulation * **Allosteric regulation** involves molecules binding to an enzyme at a site other than the active site (an **allosteric site**) [3](#page=3). * Binding to an allosteric site can either increase (allosteric activation) or decrease (allosteric inhibition) enzyme activity by causing a conformational change in the enzyme [3](#page=3). * Allosteric inhibitors can decrease enzyme activity by stabilizing a less active conformation or by preventing substrate binding [3](#page=3). * **Allosteric enzymes**, often multi-subunit proteins with quaternary structure, do not typically follow Michaelis-Menten kinetics. Their activity versus substrate concentration plots often show a **sigmoidal shape**, indicating cooperativity [3](#page=3). #### 5.3.2 Reversible covalent modifications * Enzymes involved in signal transduction pathways are frequently regulated by reversible covalent modifications [3](#page=3). * Common modifications include **phosphorylation** (addition of a phosphate group), acetylation, methylation, and carboxylation [3](#page=3). * For example, phosphorylation can induce a conformational change in an enzyme, leading to activation or inactivation, and this modification can be reversed by other enzymes (e.g., phosphatases) [3](#page=3). #### 5.3.3 Enzyme inhibition Enzyme inhibitors are molecules that bind to enzymes and reduce their activity [3](#page=3). * **Reversible inhibitors** form temporary, non-covalent bonds with enzymes and can be further classified: * **Competitive inhibitors:** These molecules bind to the enzyme's active site, competing with the substrate. They increase the apparent $K_m$ but do not affect $V_{max}$, as their effect can be overcome by a sufficiently high substrate concentration [3](#page=3). * **Uncompetitive inhibitors:** These inhibitors bind only to the enzyme-substrate (ES) complex, preventing product release. They reduce both $V_{max}$ and $K_m$ proportionally [3](#page=3). * **Non-competitive inhibitors:** These bind to a site distinct from the active site and reduce enzyme activity by decreasing the effective concentration of active enzyme. They decrease $V_{max}$ but do not alter $K_m$ [3](#page=3). * **Mixed inhibitors:** These can bind to both the enzyme and the ES complex, affecting both $V_{max}$ and $K_m$ in a more complex manner [3](#page=3). * **Irreversible inhibitors** bind covalently to the enzyme, permanently inactivating it [3](#page=3). #### 5.3.4 Cofactors * Many enzymes require **cofactors**, which are non-protein chemical compounds that assist in enzymatic function [3](#page=3). * **Coenzymes** are a type of cofactor that are small, organic molecules, often derived from vitamins, that bind to the active site and participate in the reaction [3](#page=3). #### 5.3.5 Allosteric inhibition vs. other inhibition types * Allosteric inhibitors can produce effects similar to competitive, non-competitive, or uncompetitive inhibition depending on how they affect substrate binding and enzyme conformation [3](#page=3). * A key distinction is that allosteric inhibitors regulate enzyme activity by binding at a site separate from the active site [3](#page=3). --- ## Common mistakes to avoid - Review all topics thoroughly before exams - Pay attention to formulas and key definitions - Practice with examples provided in each section - Don't memorize without understanding the underlying concepts

| Term | Definition | |---|---| | Atom | The smallest unit of an element that retains the chemical properties of that element, consisting of a nucleus (protons and neutrons) and electrons. | | Molecule | An electrically neutral entity composed of two or more atoms chemically bonded together. | | Isotope | Atoms of the same element that have the same number of protons but different numbers of neutrons, resulting in different atomic masses. | | Quantum numbers | A set of numbers used to describe the energy, shape, orientation, and spin of an electron in an atom. The four main quantum numbers are the principal quantum number (n), orbital quantum number (l), magnetic quantum number (m_L), and spin quantum number (m_S). | | Atomic orbital | A region in space around the nucleus of an atom where there is a high probability (typically 90%) of finding an electron. | | Electronegativity | A measure of an atom's ability to attract electrons in a chemical bond. | | Covalent bond | A chemical bond formed by the sharing of electron pairs between atoms. | | Ionic bond | A chemical bond formed by the electrostatic attraction between oppositely charged ions, typically resulting from the transfer of electrons from one atom to another. | | Hybridization | The mixing of atomic orbitals within an atom to form new hybrid orbitals that are better suited for bonding, resulting in different shapes and orientations. | | Sigma bond | A type of covalent bond formed by the direct, head-on overlap of atomic orbitals along the internuclear axis. | | Pi bond | A type of covalent bond formed by the sideways overlap of atomic orbitals, typically p orbitals, above and below the internuclear axis. | | VSEPR theory | Valence Shell Electron Pair Repulsion theory, which predicts the molecular geometry of molecules based on the repulsion between electron pairs in the valence shell of the central atom. | | Mole | A unit of measurement representing a specific number of particles (Avogadro's number, $6.022 \times 10^{23}$) of a substance. | | Molar mass | The mass of one mole of a substance, expressed in grams per mole (g/mol). | | Molarity | A measure of concentration defined as the number of moles of solute per liter (or cubic decimeter) of solution, expressed in mol/L or M. | | Acid | A substance that donates protons (H+) in a solution, according to the Brønsted-Lowry definition. | | Base | A substance that accepts protons (H+) in a solution, according to the Brønsted-Lowry definition. | | pH buffer | An aqueous solution that resists changes in pH upon the addition of small amounts of acid or base, typically containing a weak acid and its conjugate base or a weak base and its conjugate acid. | | Strong acid | An acid that completely dissociates in aqueous solution, releasing all its protons. | | Weak acid | An acid that only partially dissociates in aqueous solution, establishing an equilibrium between the undissociated acid and its ions. | | Isomerism | The phenomenon where two or more compounds have the same molecular formula but different structural or spatial arrangements of atoms, leading to different physical and chemical properties. | | Enantiomers | Stereoisomers that are non-superimposable mirror images of each other, meaning they have the same connectivity but differ in their three-dimensional arrangement and rotate plane-polarized light in opposite directions. | | Diastereomers | Stereoisomers that have different configurations at one or more chiral centers but are not mirror images of each other. | | Anomer | A specific type of epimer that forms at the anomeric carbon of a cyclic saccharide after ring closure. | | Hydrogen bond | A weak intermolecular or intramolecular attraction between a hydrogen atom bonded to a highly electronegative atom (like oxygen, nitrogen, or fluorine) and another electronegative atom with a lone pair of electrons. | | Thermodynamics | The branch of physics concerned with heat and its relation to other forms of energy and work. It studies the transfer and transformation of energy in chemical and physical processes. | | Enthalpy (H) | A thermodynamic property representing the total heat content of a system, defined as the sum of its internal energy and the product of its pressure and volume. | | Exothermic reaction | A reaction that releases energy, usually in the form of heat, into its surroundings, resulting in a decrease in the enthalpy of the system. | | Endothermic reaction | A reaction that absorbs energy, usually in the form of heat, from its surroundings, resulting in an increase in the enthalpy of the system. | | Entropy (S) | A thermodynamic measure of the dispersal or randomness of energy within a system. The second law of thermodynamics states that the total entropy of an isolated system can only increase over time. | | Gibbs free energy (G) | A thermodynamic potential that measures the maximum amount of reversible work that may be performed by a thermodynamic system at a constant temperature and pressure. Its change ($ \Delta G $) determines the spontaneity of a reaction: $ \Delta G < 0 $ is spontaneous, $ \Delta G > 0 $ is non-spontaneous, and $ \Delta G = 0 $ is at equilibrium. | | Catalyst | A substance that increases the rate of a chemical reaction without itself undergoing any permanent chemical change. It functions by providing an alternative reaction pathway with a lower activation energy. | | Nucleophile | A chemical species that donates an electron pair to form a covalent bond. Nucleophiles are typically electron-rich and often carry a negative charge or have a lone pair of electrons. | | Electrophile | A chemical species that accepts an electron pair to form a covalent bond. Electrophiles are typically electron-poor and often carry a positive charge or have an incomplete electron shell. | | Carbohydrates | Organic compounds with the general formula $ (CH_2O)_n $, composed of carbon, hydrogen, and oxygen, serving as primary energy sources and structural components in cells. | | Fatty acids | Long hydrocarbon chains with a carboxyl group ($ -COOH $) at one end, which are fundamental building blocks of lipids. | | Lipids | A diverse group of hydrophobic molecules that includes fats, oils, waxes, steroids, and phospholipids, essential for energy storage, cell membrane structure, and signaling. | | Membranes | Biological structures, primarily composed of phospholipid bilayers, that enclose cells and cellular organelles, regulating the passage of substances and compartmentalizing cellular functions. | | Nucleotides | The monomeric units of nucleic acids (DNA and RNA), consisting of a nitrogenous base, a pentose sugar (ribose or deoxyribose), and one or more phosphate groups. | | Proteins | Macromolecules composed of amino acid subunits linked by peptide bonds, performing a vast array of functions within cells, including enzymatic activity, structural support, and transport. | | Enzyme | A biological catalyst, typically a protein, that accelerates the rate of specific biochemical reactions by lowering the activation energy. | | Michaelis constant ($K_m$) | The substrate concentration at which the reaction rate of an enzyme-catalyzed reaction is half of its maximum velocity ($V_{max}$). It is often used as an indicator of the enzyme's affinity for its substrate. | | Maximum velocity ($V_{max}$) | The maximum rate of an enzyme-catalyzed reaction, achieved when the enzyme is saturated with substrate. | | Turnover number ($k_{cat}$) | The maximum number of substrate molecules that an enzyme molecule can convert into product per unit time, when the enzyme is operating at its maximum rate. | | Glycolysis | The metabolic pathway that converts glucose into pyruvate, producing ATP and NADH in the process. It occurs in the cytoplasm and does not require oxygen. | | Gluconeogenesis | The metabolic pathway that synthesizes glucose from non-carbohydrate precursors, such as lactate, amino acids, and glycerol, primarily occurring in the liver. | | Substrate-level phosphorylation | The direct transfer of a phosphate group from a high-energy substrate molecule to ADP, forming ATP, typically occurring during glycolysis and the Krebs cycle. | | Normoglycaemia | The condition of having a normal blood glucose concentration, typically within the range of 4-8 mM. | | Hypoglycaemia | A condition characterized by abnormally low blood glucose levels. | | Oxidative decarboxylation | A reaction that involves both the removal of carbon dioxide and the oxidation of a molecule, often coupled with the reduction of electron carriers like NAD+. | | Krebs cycle (TCA cycle) | A series of metabolic reactions occurring in the mitochondrial matrix that oxidizes acetyl-CoA, producing ATP, NADH, FADH2, and releasing carbon dioxide. | | Electron transport chain (ETC) | A series of protein complexes embedded in the inner mitochondrial membrane that transfer electrons from NADH and FADH2 to oxygen, coupled with the pumping of protons across the membrane to create a proton gradient. | | Oxidative phosphorylation | The process by which ATP is synthesized using the energy derived from the electron transport chain and the subsequent flow of protons back across the inner mitochondrial membrane through ATP synthase. | | Sterocenter | An atom in a molecule that is bonded to four different atoms or groups, such that switching any two of these groups results in a stereoisomer. | | Stereoisomers | Molecules with the same molecular formula and connectivity but different spatial arrangements of atoms. | | Structural isomers | Molecules with the same molecular formula but different bonding arrangements of atoms. | | Optical isomer | Stereoisomers that are non-superimposable mirror images of each other. | | Racemate | An equimolar mixture of two enantiomers. | | Fischer projection | A two-dimensional representation of a three-dimensional molecule, typically used for sugars and amino acids, where horizontal lines represent bonds projecting out of the plane and vertical lines represent bonds projecting into the plane. | | Anomeric carbon | The carbon atom in a cyclic saccharide that was originally the carbonyl carbon (aldehyde or ketone) in the open-chain form and is now bonded to two oxygen atoms. | | Glycosidic bond | A type of covalent bond that links a carbohydrate molecule to another group, which may be another carbohydrate, a protein, or a lipid. | | Phospholipid | A type of lipid that is a major component of cell membranes, consisting of a hydrophilic head (a phosphate group linked to a variable group) and two hydrophobic fatty acid tails. | | Triacylglycerol | A lipid formed from glycerol esterified with three fatty acids, serving as a major form of energy storage. | | Sphingolipid | A class of lipids that have a backbone of sphingosine, a complex amino alcohol, and are important components of cell membranes, particularly in the nervous system. | | Amphiphilic | Having both hydrophilic (water-attracting) and hydrophobic (water-repelling) properties, characteristic of molecules like phospholipids that form cell membranes. | | Hydrophobic effect | The tendency of nonpolar molecules to aggregate in aqueous solution, driven by the exclusion of water molecules and the increase in entropy of the solvent. | | DNA double helix | The characteristic structure of DNA, consisting of two antiparallel polynucleotide strands wound around a central axis, held together by hydrogen bonds between complementary base pairs. | | Peptide bond | A covalent bond formed between the carboxyl group of one amino acid and the amino group of another, linking amino acids together to form proteins. | | Primary structure of protein | The linear sequence of amino acids in a polypeptide chain. | | Secondary structure of protein | Localized folding of the polypeptide chain into regular structures like alpha-helices and beta-sheets, stabilized by hydrogen bonds between backbone atoms. | | Tertiary structure of protein | The overall three-dimensional shape of a single polypeptide chain, formed by various interactions between amino acid side chains. | | Quaternary structure of protein | The arrangement of multiple polypeptide subunits in a multi-subunit protein complex. | | Bioanalysis | Techniques used to study the structure, function, and interactions of biological molecules. | | Chromatography | A laboratory technique used to separate mixtures of compounds based on their differing distributions between a stationary phase and a mobile phase. | | SDS-PAGE | Sodium dodecyl sulfate-polyacrylamide gel electrophoresis, a technique used to separate proteins based on their molecular weight. | | ATP synthase | An enzyme complex embedded in the inner mitochondrial membrane that uses the energy of a proton gradient to synthesize ATP from ADP and inorganic phosphate. | | Aerobic respiration | The metabolic process that occurs in the presence of oxygen to convert glucose into carbon dioxide and water, generating a large amount of ATP. | | Glycolysis | The metabolic pathway that breaks down glucose into two molecules of pyruvate, producing a net of two ATP molecules and two NADH molecules. | | Gluconeogenesis (GNG) | The synthesis of glucose from non-carbohydrate precursors, such as lactate, amino acids, and glycerol. | | Krebs cycle (TCA cycle) | A series of chemical reactions used to generate energy, in the form of ATP, NADH, and FADH2, by oxidizing acetyl-CoA. | | Oxidative phosphorylation | The process by which ATP is synthesized using the energy released from the transfer of electrons from NADH and FADH2 to oxygen, via the electron transport chain. | | Electron transport chain (ETC) | A series of protein complexes in the inner mitochondrial membrane that facilitate the transfer of electrons and pump protons, creating a gradient that drives ATP synthesis. | | Proton-motive force | The potential energy stored in the proton gradient across a membrane, generated by the electron transport chain, which powers ATP synthesis. | | Warburg effect | The observation that cancer cells often exhibit increased glycolysis even in the presence of oxygen, a phenomenon that supports rapid cell proliferation. | | Endothermic reaction | A reaction that absorbs heat from its surroundings, leading to a decrease in temperature. | | Exothermic reaction | A reaction that releases heat into its surroundings, leading to an increase in temperature. | | Gibbs energy | A thermodynamic potential that measures the maximum amount of non-expansion work that can be extracted from a thermodynamic system. Its change ($ \Delta G $) indicates the spontaneity of a process. | | Catalyst | A substance that increases the rate of a chemical reaction without being consumed in the process. | | Carbohydrates | Organic compounds that serve as a primary source of energy and play structural roles in living organisms, typically with the empirical formula $ (CH_2O)_n $. | | Fatty acids | Long hydrocarbon chains with a terminal carboxyl group, forming the basis of many lipids. | | Lipids | A broad group of naturally occurring molecules including fats, waxes, sterols, fat-soluble vitamins, monoglycerides, diglycerides, triglycerides, and phospholipids, which are characterized by their insolubility in water. | | Membranes | Semipermeable barriers that enclose cells and organelles, primarily composed of a lipid bilayer. | | Nucleotides | The building blocks of nucleic acids (DNA and RNA), consisting of a nitrogenous base, a pentose sugar, and a phosphate group. | | Proteins | Macromolecules made up of amino acid chains, responsible for a vast range of functions in living organisms. | | Enzyme | A biological catalyst that speeds up biochemical reactions by lowering the activation energy. | | $K_m$ | The Michaelis constant, representing the substrate concentration at which the enzyme's activity is half of its maximum. | | $V_{max}$ | The maximum rate of an enzyme-catalyzed reaction when the enzyme is saturated with substrate. | | $k_{cat}$ | The catalytic constant, representing the number of substrate molecules converted to product per enzyme molecule per second. | | Glycolysis | The metabolic pathway that converts glucose into pyruvate, producing ATP and NADH. | | Gluconeogenesis (GNG) | The synthesis of glucose from non-carbohydrate precursors. | | Investment phase (of glycolysis) | The initial steps of glycolysis where ATP is consumed to activate glucose and prepare it for cleavage. | | Payoff phase (of glycolysis) | The latter stages of glycolysis where ATP and NADH are produced through substrate-level phosphorylation and oxidation reactions. | | Substrate-level phosphorylation | The direct transfer of a phosphate group from a high-energy substrate molecule to ADP, forming ATP. | | Normoglycaemia | Normal blood glucose levels. | | Hypoglycaemia | Low blood glucose levels. | | Oxidative decarboxylation | A reaction that removes carbon dioxide and oxidizes a molecule, often coupled with NAD+ reduction. | | Krebs/TCA cycle | A central metabolic pathway that oxidizes acetyl-CoA, generating ATP, NADH, and FADH2. | | Electron transport chain (ETC) | A series of protein complexes in the inner mitochondrial membrane that transfer electrons, ultimately to oxygen, creating a proton gradient. | | Oxidative phosphorylation | The process of ATP synthesis driven by the proton gradient established by the ETC, utilizing ATP synthase. | | Uncouplers | Compounds that disrupt the coupling between the electron transport chain and ATP synthesis by dissipating the proton gradient, leading to increased oxygen consumption without ATP production. | | Warburg effect | The observation that cancer cells preferentially use glycolysis even in the presence of oxygen. | | Endothermic reaction | A reaction that absorbs heat from the surroundings. | | Exothermic reaction | A reaction that releases heat into the surroundings. | | Gibbs energy | A thermodynamic potential that determines the spontaneity of a process. | | Catalyst | A substance that increases the rate of a chemical reaction without being consumed. | | Nucleophile | An electron-pair donor. | | Electrophile | An electron-pair acceptor. | | Carbohydrates | Sugars, starches, and cellulose, important for energy and structure. | | Fatty acids | Long hydrocarbon chains with a carboxyl group, forming lipids. | | Lipids | Hydrophobic molecules essential for cell membranes and energy storage. | | Membranes | Lipid bilayers that enclose cells and organelles. | | Nucleotides | The monomers of nucleic acids (DNA and RNA). | | Proteins | Polymers of amino acids with diverse biological functions. | | Enzyme | A biological catalyst that accelerates reactions. | | $K_m$ | Michaelis constant, indicating substrate concentration at half maximal velocity. | | $V_{max}$ | Maximum reaction velocity of an enzyme. | | $k_{cat}$ | Turnover number, the rate at which an enzyme converts substrate to product. | | Glycolysis | The breakdown of glucose into pyruvate. | | Gluconeogenesis (GNG) | The synthesis of glucose from non-carbohydrate precursors. | | Investment phase (of glycolysis) | The initial energy-consuming steps of glycolysis. | | Payoff phase (of glycolysis) | The ATP and NADH producing steps of glycolysis. | | Substrate level phosphorylation | Direct ATP synthesis from a substrate. | | Normoglycaemia | Normal blood glucose levels. | | Hypoglycaemia | Low blood glucose levels. | | Oxidative decarboxylation | Removal of CO2 and oxidation, e.g., pyruvate to acetyl-CoA. | | Krebs/TCA cycle | A metabolic pathway that oxidizes acetyl-CoA, producing ATP, NADH, and FADH2. | | Electron transport chain (ETC) | A series of protein complexes that transfer electrons, generating a proton gradient. | | Oxidative phosphorylation | ATP synthesis coupled to the ETC and proton gradient. |

# Constitution et caractéristiques de l'atome L'atome, unité fondamentale de la matière, est composé d'un noyau central et d'électrons en orbite, chaque constituant possédant des caractéristiques bien définies qui déterminent l'identité et le comportement d'un élément chimique [1](#page=1). ### 1.1 Constituants de l'atome L'atome est la plus petite unité de matière qui conserve les propriétés chimiques d'un élément. Il est constitué de deux parties principales: le noyau et les électrons [1](#page=1). #### 1.1.1 Le noyau Le noyau est la partie centrale de l'atome et contient la quasi-totalité de sa masse. Il est composé de particules élémentaires appelées nucléons [1](#page=1). * **Neutrons (N)**: Particules de charge électrique nulle ($q_N = 0$ C) et de masse $m_N = 1,675 \times 10^{-27}$ kg [1](#page=1). * **Protons (p)**: Particules de charge électrique positive égale à $q_p = +e = 1,6 \times 10^{-19}$ C et de masse $m_p = 1,673 \times 10^{-27}$ kg [1](#page=1). #### 1.1.2 Les électrons (e) Les électrons sont des particules élémentaires chargées négativement, avec une charge $q_e = -1,6 \times 10^{-19}$ C. Ils gravitent autour du noyau dans des régions suivant des trajectoires bien définies. Leur masse est d'environ $m_e = 9,109 \times 10^{-31}$ kg, ce qui est beaucoup plus léger que celle des protons et des neutrons [1](#page=1). ### 1.2 Caractéristiques de l'atome La représentation générale d'un atome (élément chimique) X est la suivante : $$ ^A_Z X $$ #### 1.2.1 Numéro atomique (Z) Le numéro atomique, $Z$, représente le nombre de protons qui composent le noyau d'un atome [1](#page=1). * $Z$ caractérise un élément et détermine son identité [1](#page=1). * Pour un atome neutre, $Z$ est également égal au nombre d'électrons [1](#page=1). * Pour un atome ionisé, $Z$ reste inchangé, mais le nombre d'électrons varie [1](#page=1). * Toutes les espèces chimiques ayant le même numéro atomique $Z$ appartiennent au même élément chimique [1](#page=1). > **Tip:** Un atome chargé est appelé un ion. Un ion positif est un cation, tandis qu'un ion négatif est un anion. **Exemple**: Les espèces Fe ($Z=26$), Fe$^{2+}$ ($Z=26$), et Fe$^{3+}$ ($Z=26$) appartiennent toutes à l'élément Fer car elles partagent le même numéro atomique [1](#page=1). #### 1.2.2 Nombre de masse (A) Le nombre de masse, $A$, représente la somme du nombre de protons ($Z$) et du nombre de neutrons ($N$) dans le noyau d'un atome. Il est calculé par la formule [2](#page=2): $$ A = Z + N $$ * $A$ est utilisé pour calculer la masse atomique d'un élément. La masse d'un atome, exprimée en unités de masse atomique, est approximativement égale à la somme des masses des protons et des neutrons, la masse des électrons étant négligeable [2](#page=2). #### 1.2.3 Unité de masse atomique (u.m.a.) L'unité de masse atomique (u.m.a.) est définie comme 1/12 de la masse d'un atome de Carbone-12 ($^{12}$C). Par convention, une mole de $^{12}$C pèse 12 grammes et contient $N_A$ atomes de carbone (où $N_A$ est le nombre d'Avogadro) [2](#page=2). $$ 1 \text{ u.m.a.} = \frac{1}{12} m(\text{1 atome de } ^{12}\text{C}) $$ $$ 1 \text{ u.m.a.} = \frac{1}{N_A} \times \frac{12 \text{ g}}{12} = \frac{1 \text{ g}}{N_A} \approx 1,67 \times 10^{-24} \text{ g} = 1,67 \times 10^{-27} \text{ Kg} $$ Notez que 1 u.m.a. est approximativement égale à la masse d'un proton ($m_p$) ou d'un neutron ($m_N$) [2](#page=2). > **Tip:** La masse molaire d'un élément exprimée en grammes par mole (g/mol) est numériquement égale à sa masse atomique exprimée en unités de masse atomique (u.m.a.). Par exemple, la masse atomique d'un atome de $^{14}$N est d'environ 14 u.m.a., et la masse molaire d'une mole de $^{14}$N est d'environ 14 g/mol [2](#page=2). #### 1.2.4 Isotopes Les isotopes d'un élément donné sont des atomes qui possèdent le même nombre de protons (donc le même $Z$) mais un nombre différent de neutrons (donc un $A$ différent). Ils ne diffèrent que par la composition de leur noyau [2](#page=2). **Exemple** : Le chlore (Cl) possède deux isotopes : * $^{35}_{17}$Cl (17 protons, 18 neutrons) [2](#page=2). * $^{37}_{17}$Cl (17 protons, 20 neutrons) [2](#page=2). #### 1.2.5 Abondance isotopique L'abondance isotopique correspond au pourcentage relatif d'un isotope dans un échantillon naturel de cet élément. La masse molaire atomique d'un élément est déterminée en tenant compte de l'abondance naturelle de ses isotopes [2](#page=2) [3](#page=3). Pour calculer la masse molaire atomique ($M$) d'un élément avec plusieurs isotopes, on utilise la formule suivante, où $M_i$ est la masse molaire de l'isotope $i$ et $v_i$ est son abondance naturelle : $$ M = \sum_i (v_i \times M_i) $$ **Exemple**: Le chlore (Cl) a deux isotopes: $^{35}$Cl (abondance de 75,4%) et $^{37}$Cl (abondance de 24,6%). La masse molaire atomique du chlore est donc [3](#page=3): $M_{Cl} = (35 \times 75,4\%) + (37 \times 24,6\%) = 35,5$ g/mol [3](#page=3). --- # Structure électronique et nombre quantiques Ce chapitre explore la nature dualiste du rayonnement électromagnétique, les interactions énergétiques au sein des atomes, et introduit le système des nombres quantiques qui décrivent les orbitales atomiques. ### 2.1 Rayonnement électromagnétique Le rayonnement électromagnétique, qu'il soit visible ou invisible, possède une nature dualiste, se manifestant à la fois comme une onde et comme un flux de particules [3](#page=3). #### 2.1.1 L'aspect ondulatoire En tant qu'onde, le rayonnement électromagnétique est caractérisé par sa fréquence ($v$) en Hertz (s⁻¹), sa longueur d'onde ($\lambda$) en mètres (m), et sa vitesse de propagation dans le vide ($c$), qui est de $3 \times 10^8$ m/s. Ces grandeurs sont liées par la formule [3](#page=3): $$ \lambda = \frac{c}{v} $$ #### 2.1.2 L'aspect corpusculaire En tant que corpuscule, le rayonnement est considéré comme un flux de particules appelées photons ou quanta. L'énergie ($E$) d'un photon est directement proportionnelle à la fréquence de la radiation, selon la relation de Planck : $$ E = h v $$ où $h$ est la constante de Planck, dont la valeur est $6,62 \times 10^{-34}$ J·s [3](#page=3). #### 2.1.3 Spectre électromagnétique Le spectre électromagnétique s'étend sur une large gamme de longueurs d'onde, allant des rayons gamma aux ondes radio, incluant les rayons X, le rayonnement ultraviolet (UV), la lumière visible, l'infrarouge et les micro-ondes [4](#page=4). > **Tip:** Comprendre la relation entre longueur d'onde et fréquence est essentiel. Une longueur d'onde plus courte implique une fréquence plus élevée, et donc une énergie de photon plus grande. #### 2.1.4 Absorption et émission d'énergie Les électrons au sein d'un atome ne peuvent absorber ou émettre de l'énergie que lors de transitions entre différents niveaux énergétiques (orbitales) [4](#page=4). ##### 2.1.4.1 La relation de Planck pour les transitions La quantité d'énergie absorbée ou émise lors d'une transition électronique est exactement égale à la différence d'énergie entre les niveaux initial ($E_i$) et final ($E_f$). Cette relation est exprimée par : $$ \Delta E = |E_f - E_i| = h v $$ ##### 2.1.4.2 Transitions électroniques * **Absorption:** Lorsqu'un électron passe d'un niveau d'énergie inférieur vers un niveau supérieur, il absorbe un photon dont l'énergie correspond à la différence entre ces deux niveaux: $\Delta E = E_f - E_i$. Dans ce cas, $\Delta E > 0$ [4](#page=4). * **Émission:** Lorsqu'un électron retourne d'un niveau d'énergie supérieur à un niveau inférieur, il émet un photon dont l'énergie correspond à la différence entre ces deux niveaux: $\Delta E = |E_f - E_i|$. Dans ce cas, $\Delta E < 0$ [4](#page=4). > **Tip:** Visualisez ces transitions comme des "sauts" quantifiés d'un électron. L'atome ne peut gagner ou perdre que des paquets d'énergie discrets, correspondant aux photons. #### 2.1.5 États de l'électron Les états d'un électron dans un atome sont définis par le nombre quantique principal ($n$) [5](#page=5). * **État fondamental:** Correspond à $n=1$. C'est l'état de plus basse énergie [5](#page=5). * **État excité:** Correspond à $n>1$. L'électron possède plus d'énergie que dans l'état fondamental [5](#page=5). * **Ionisation:** Correspond à $n \rightarrow \infty$. L'électron est complètement détaché de l'atome [5](#page=5). L'énergie d'ionisation est l'énergie minimale requise pour extraire un électron de son état fondamental vers l'infini [5](#page=5). #### 2.1.6 Spectre d'émission de l'atome d'hydrogène L'excitation de l'atome d'hydrogène gazeux conduit à l'émission de lumière qui, lorsqu'elle est décomposée, révèle un spectre discret de raies spectrales. Ces raies correspondent à des longueurs d'onde spécifiques associées aux transitions électroniques [5](#page=5). Les principales séries de raies observées pour l'atome d'hydrogène sont : * **Série de Lyman:** Transitions vers $n=1$ (domaine UV) [5](#page=5). * **Série de Balmer:** Transitions vers $n=2$ (domaine visible) [5](#page=5). * **Série de Paschen:** Transitions vers $n=3$ (domaine infrarouge) [5](#page=5). * **Série de Brackett:** Transitions vers $n=4$ (domaine infrarouge) [5](#page=5). * **Série de Pfund:** Transitions vers $n=5$ (domaine infrarouge) [5](#page=5). La formule de Ritz permet de calculer les longueurs d'onde de ces raies : $$ \frac{1}{\lambda} = R_H Z^2 \left( \frac{1}{n^2} - \frac{1}{p^2} \right) $$ où $R_H$ est la constante de Rydberg ($1,1 \times 10^7$ m⁻¹), $Z$ est le numéro atomique, et $n$ et $p$ sont des entiers naturels avec $p > n$ [5](#page=5). > **Tip:** Le modèle de Bohr, bien qu'explicatif pour l'hydrogène, est limité. La mécanique quantique est nécessaire pour décrire les atomes polyélectroniques [6](#page=6). #### 2.1.7 Énergies et rayons atomiques Pour l'hydrogène ($Z=1$) et les espèces hydrogénoïdes ($Z>1$), les énergies des niveaux ($E_n$) et les rayons des orbitales ($r_n$) peuvent être calculés : * **Énergie :** * Hydrogène: $E_n = -13,6 \times \frac{1}{n^2}$ eV [6](#page=6). * Hydrogénoïde: $E_n = -13,6 \times \frac{Z^2}{n^2}$ eV [6](#page=6). * **Rayon :** * Hydrogène: $r_n = 0,53 \times n^2$ Å [6](#page=6). * Hydrogénoïde: $r_n = 0,53 \times \frac{n^2}{Z}$ Å [6](#page=6). Une unité d'énergie couramment utilisée est l'électron-volt (eV), où $1$ eV $= 1,6 \times 10^{-19}$ Joules [6](#page=6). ### 2.2 Nombres quantiques Les nombres quantiques sont un ensemble de valeurs qui décrivent l'état d'un électron dans un atome, notamment son énergie, la forme et l'orientation de son orbitale, ainsi que son spin [7](#page=7) [8](#page=8). #### 2.2.1 Nombre quantique principal ($n$) * **Définition:** $n \in \mathbb{N}^*$, donc $n = 1, 2, 3, \dots$ [7](#page=7). * **Rôle:** Définit la couche électronique et détermine principalement l'énergie de l'électron [7](#page=7). * **Désignation des couches :** * $n=1$: Couche K * $n=2$: Couche L * $n=3$: Couche M * $n=4$: Couche N * $n=5$: Couche O * $n=6$: Couche P * $n=7$: Couche Q * **Capacité maximale d'électrons par couche:** $2n^2$ [7](#page=7). * **Nombre d'orbitales par couche:** $n^2$ [7](#page=7). #### 2.2.2 Nombre quantique secondaire ou azimutal ($l$) * **Définition:** $l \in \{0, 1, 2, \dots, n-1\}$ [8](#page=8). * **Rôle:** Caractérise la forme de l'orbitale et définit une sous-couche électronique [8](#page=8). * **Désignation des sous-couches :** * $l=0$: sous-couche s * $l=1$: sous-couche p * $l=2$: sous-couche d * $l=3$: sous-couche f * $l=4$: sous-couche g (et ainsi de suite) * **Capacité maximale d'électrons par sous-couche:** $2(2l+1)$ [8](#page=8). > **Example:** Pour la couche $n=2$, $l$ peut prendre les valeurs 0 et 1, correspondant aux sous-couches 2s et 2p. #### 2.2.3 Nombre quantique magnétique ($m_l$) * **Définition:** $m_l \in \{-l, -l+1, \dots, 0, \dots, l-1, l\}$. Il y a donc $2l+1$ valeurs possibles pour $m_l$ pour un $l$ donné [8](#page=8). * **Rôle:** Détermine le nombre d'orbitales au sein d'une sous-couche et leur orientation spatiale [8](#page=8). * **Nombre d'orbitales par sous-couche :** $2l+1$. * $l=0$ (s) $\rightarrow m_l = 0$ : 1 orbitale s * $l=1$ (p) $\rightarrow m_l = -1, 0, +1$ : 3 orbitales p * $l=2$ (d) $\rightarrow m_l = -2, -1, 0, +1, +2$ : 5 orbitales d * $l=3$ (f) $\rightarrow m_l = -3, -2, -1, 0, +1, +2, +3$ : 7 orbitales f > **Tip:** Chaque orbitale atomique est définie par une combinaison unique des nombres quantiques $n$, $l$, et $m_l$. #### 2.2.4 Nombre quantique de spin ($m_s$ ou $s$) * **Définition:** $s = +1/2$ ou $s = -1/2$ [8](#page=8). * **Rôle:** Caractérise le moment cinétique intrinsèque de l'électron, souvent visualisé comme une rotation sur lui-même (spin) [8](#page=8). * **Représentation graphique:** Une flèche vers le haut ($\uparrow$) pour $s=+1/2$ et une flèche vers le bas ($\downarrow$) pour $s=-1/2$ sont couramment utilisées pour représenter le spin d'un électron dans une case quantique [8](#page=8). > **Tip:** Le principe d'exclusion de Pauli stipule que deux électrons dans un même atome ne peuvent pas avoir le même ensemble de quatre nombres quantiques ($n, l, m_l, m_s$). --- # Configuration électronique et tableaux périodiques Cette section aborde la manière dont les électrons sont répartis dans les orbitales atomiques, les règles qui gouvernent ce remplissage, les exceptions notables, et comment cette organisation électronique se reflète dans la structure et les propriétés du tableau périodique. ### 3.1 Configuration électronique des atomes La configuration électronique d'un atome décrit la disposition de ses électrons dans les orbitales atomiques lorsque l'atome est dans son état fondamental. Le remplissage des orbitales doit respecter trois principes fondamentaux [9](#page=9): 1. **Principe d'exclusion de Pauli:** Chaque orbitale ne peut contenir qu'un maximum de deux électrons, et ces électrons doivent avoir des spins opposés [9](#page=9). 2. **Principe de stabilité (règle de Klechkowski):** Les électrons remplissent d'abord les orbitales atomiques de plus basse énergie avant de passer à celles de plus haute énergie [9](#page=9). 3. **Règle de Hund:** Lorsque des orbitales atomiques possèdent la même énergie (dégénérées), les électrons s'arrangent pour occuper le maximum d'orbitales possibles avec des spins parallèles avant que des paires d'électrons ne se forment dans une même orbitale [9](#page=9). La règle de Klechkowski permet de déterminer l'ordre de remplissage des sous-couches quantiques basé sur la somme du nombre quantique principal ($n$) et du nombre quantique secondaire ($l$), soit ($n+l$). Si deux orbitales ont la même somme ($n+l$), celle avec le plus petit $n$ est remplie en premier. L'ordre de remplissage peut être visualisé par le diagramme de Klechkowski, qui suit l'ordre: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p [9](#page=9). **Exemples de configurations électroniques :** * Carbone (Z = 6): $1s^2 2s^2 2p^2$ [9](#page=9). * Sodium (Z = 11): $1s^2 2s^2 2p^6 3s^1$ [9](#page=9). * Zinc (Z = 30): $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^{10}$ [9](#page=9). #### 3.1.1 Exceptions aux règles de remplissage Certains éléments présentent des configurations électroniques qui dérogent à la règle de Klechkowski, notamment ceux possédant une sous-couche $d$ ou $f$ qui tend à être totalement remplie ($d^{10}$, $f^{14}$) ou à moitié remplie ($d^5$, $f^7$). Cette semi-remplissage ou remplissage complet confère une stabilité accrue aux atomes [10](#page=10). * **Cas typiques:** Configurations du type $d^9 s^2$ (comme le cuivre et l'argent) ou $d^4 s^2$ (comme le chrome et le molybdène) tendent à se stabiliser en $d^{10} s^1$ ou $d^5 s^1$, respectivement [10](#page=10). * **Exemple pour le Chrome (Cr, Z=24) :** Configuration attendue : $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^4$ Configuration réelle: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^5$ [10](#page=10). * **Exemple pour le Cuivre (Cu, Z=29) :** Configuration attendue : $1s^2 2s^2 2p^6 3s^2 3p^6 4s^2 3d^9$ Configuration réelle: $1s^2 2s^2 2p^6 3s^2 3p^6 4s^1 3d^{10}$ [10](#page=10). #### 3.1.2 Atomes polyélectroniques et approximation de Slater Dans les atomes polyélectroniques (contenant plusieurs électrons), l'énergie d'un électron est affectée par l'effet d'écran créé par les autres électrons. L'approximation de Slater permet d'estimer cette énergie en introduisant une charge effective ($Z^*$) et un nombre quantique apparent ($n^*$). L'énergie d'un électron dans un atome polyélectronique est approximée par la formule : $$E_n = -13.6 \frac{Z^{*2}}{n^{*2}} \text{ eV}$$ [10](#page=10). où : * $Z^*$ est la charge effective subie par l'électron. * $n^*$ est le nombre quantique apparent, estimé selon un tableau basé sur le nombre quantique principal $n$ [10](#page=10). ##### 3.1.2.1 Calcul de la charge effective ($Z^*$) La charge effective est calculée par la formule : $$Z^* = Z - \sigma$$ [10](#page=10). où : * $Z$ est la charge nucléaire réelle de l'atome. * $\sigma$ est la constante d'écran, qui dépend de la position de l'électron étudié par rapport aux autres électrons. Les valeurs de $\sigma$ sont déterminées par les règles de Slater, qui groupent les électrons en catégories [11](#page=11). Le tableau de Slater fournit les valeurs de $\sigma$ pour chaque type d'électron, en fonction de sa localisation par rapport à l'électron étudié [11](#page=11). > **Tip :** Les règles de Slater pour le calcul de $\sigma$ distinguent les groupes d'électrons en fonction de leurs nombres quantiques ($n$ et $l$) et de leur proximité avec l'électron étudié. Les électrons dans le même groupe ont une contribution différente de ceux des groupes internes ou externes. **Exemple pour le Chlore (Cl, Z=17) :** Configuration électronique : $1s^2 2s^2 2p^6 3s^2 3p^5$. Considérons un électron dans la sous-couche $3s$ ou $3p$. * Électrons du même groupe ($3s, 3p$) : $\sigma = 0.35$ par électron. Il y a 6 autres électrons ($2s^2 2p^6$). * Électrons des groupes internes ($2s, 2p$) : $\sigma = 0.85$ par électron. Il y a 8 électrons. * Électrons des groupes les plus internes ($1s$) : $\sigma = 1$ par électron. Il y a 2 électrons. Application des règles de Slater pour un électron 3s/3p : * Dans le même groupe ($3s, 3p$) : $6 \times 0.35 = 2.1$ * Dans les groupes précédents ($2s, 2p$) : $8 \times 0.85 = 6.8$ * Dans les groupes internes ($1s$) : $2 \times 1 = 2$ La constante d'écran totale pour un électron $3s$ ou $3p$ est $\sigma = 2.1 + 6.8 + 2 = 10.9$. La charge effective est donc: $Z^* = Z - \sigma = 17 - 10.9 = 6.1$ [11](#page=11). L'énergie totale d'un atome est la somme des énergies de tous ses électrons, calculées individuellement avec leur charge effective et leur nombre quantique apparent respectifs [11](#page=11). ### 3.2 Tableau périodique Le tableau périodique organise les éléments chimiques en fonction de leurs propriétés et configurations électroniques, permettant de prédire leur comportement chimique [12](#page=12). #### 3.2.1 Structure du tableau périodique Le tableau périodique est structuré comme suit : * **Périodes:** 7 lignes horizontales, chaque ligne correspondant à un niveau d'énergie principal $n$ (de $n=1$ à $n=7$) [12](#page=12). * **Groupes (ou familles):** 18 colonnes verticales. Les éléments d'un même groupe partagent des configurations électroniques de valence similaires, leur conférant des propriétés chimiques proches [12](#page=12). * **Blocs :** Le tableau est divisé en quatre blocs (s, p, d, f) basés sur la sous-couche électronique en cours de remplissage : * **Bloc s:** La couche périphérique est de type $ns^x$ ($x=1$ ou $2$). Il comprend les groupes IA (alcalins) et IIA (alcalino-terreux). Ces éléments sont des métaux très réducteurs [12](#page=12). * **Bloc p:** La couche périphérique est de type $ns^2 np^x$ ($x=1$ à $6$). Il inclut les groupes IIIA à VIIIA (halogènes et gaz rares) [12](#page=12). * **Bloc d:** La couche périphérique est de type $ns^2 (n-1)d^x$ ($x=1$ à $10$). Ces éléments sont appelés métaux de transition et appartiennent aux groupes IIIB à IIB. Le groupe VIIIB (triades de fer, cobalt, nickel) est une exception où le chiffre romain n'indique pas le nombre d'électrons de valence [12](#page=12) [13](#page=13). * **Bloc f:** La couche périphérique est de type $ns^x (n-1)d^y (n-2)f^z$ ($x=0$ ou $1$, $y=1$ à $14$). Il comprend les lanthanides (Z=57 à 71) et les actinides (Z=89 à 103) [13](#page=13). > **Remarque:** Il est important de distinguer la "couche externe" (celle avec le plus grand $n$) de l'"orbitale externe" (celle avec le plus grand $n$, et en cas d'égalité, le plus grand $l$) [13](#page=13). #### 3.2.2 Propriétés périodiques Certaines propriétés des éléments varient de manière prévisible à travers le tableau périodique. * **Rayon atomique (r$_a$):** Le rayon atomique est la moitié de la distance entre les noyaux de deux atomes liés par une liaison simple [13](#page=13). * **Dans une période:** Le rayon atomique diminue de gauche à droite. L'augmentation de la charge nucléaire effective attire davantage les électrons, réduisant le rayon [13](#page=13). * **Dans une colonne:** Le rayon atomique augmente de haut en bas. L'ajout de couches électroniques supplémentaires éloigne la couche de valence du noyau [13](#page=13). * **Énergie d'ionisation (E$_i$):** L'énergie minimale requise pour arracher un électron d'un atome gazeux isolé est positive [13](#page=13). * **Dans une période:** L'énergie d'ionisation augmente de gauche à droite. La force d'attraction entre le noyau et les électrons augmente, rendant l'arrachement plus difficile [14](#page=14). * **Dans une colonne:** L'énergie d'ionisation augmente de bas en haut. Les électrons sont plus proches du noyau et donc plus fortement liés [14](#page=14). * **Exceptions:** Des exceptions existent, par exemple E$_i$(B) < E$_i$(Be) et E$_i$(O) < E$_i$(N) [14](#page=14). * Pour un atome hydrogénoïde: $E_i = +13.6 \frac{Z^2}{n^2}$ [14](#page=14). * Pour un atome polyélectronique: $E_i = +13.6 \frac{Z^{*2}}{n^{*2}}$ [14](#page=14). * **Affinité électronique (AE):** L'énergie libérée ou absorbée lors de la capture d'un électron par un atome gazeux isolé. Les halogènes ont les affinités électroniques les plus élevées, tandis que les alcalins ont les plus faibles [14](#page=14). * **Électronégativité ($\chi$):** Le pouvoir d'un atome d'attirer les électrons d'une liaison chimique [14](#page=14). * **Dans une période:** L'électronégativité augmente de gauche à droite. Les atomes plus proches d'une couche électronique complète ont une forte tendance à attirer des électrons [14](#page=14). * **Dans une colonne:** L'électronégativité augmente de bas en haut. Les électrons de valence sont plus proches du noyau et donc plus fortement attirés [14](#page=14). --- # Liaisons chimiques et géométrie moléculaire Ce chapitre explore les fondements des liaisons chimiques, de la représentation des molécules à la prédiction de leur forme tridimensionnelle, en passant par les théories quantiques de la formation des liaisons. ### 4.1 La liaison chimique La liaison chimique est l'interaction qui unit les atomes pour former des molécules ou des composés, leur permettant d'atteindre une configuration électronique plus stable. Les électrons de valence sont les principaux acteurs de ces liaisons. Deux types de liaisons sont principalement abordés: covalente et ionique [15](#page=15). #### 4.1.1 La liaison covalente #### 4.1.1.1 Théorie de Lewis La théorie de Lewis postule que la liaison chimique résulte de la mise en commun de paires d'électrons entre atomes [15](#page=15). * **Liaison covalente normale:** Chaque atome contribue un électron à la paire liante [15](#page=15). * Représentation: $A:B$ ou $A-B$ [15](#page=15). * **Liaison covalente dative (ou de coordinence):** Un atome (donneur) fournit le doublet d'électrons à un autre atome (accepteur) [15](#page=15). * Représentation: $A \leftrightarrow B$ ou $A \leftarrow B$ [15](#page=15). ##### 4.1.1.1.1 Règle de l'octet et du duet * La **règle de l'octet** stipule que les atomes tendent à former des liaisons jusqu'à ce qu'ils soient entourés de 8 électrons de valence, imitant ainsi la configuration électronique des gaz rares. Cette règle s'applique principalement aux éléments des blocs s et p avec un numéro atomique supérieur à 4 [16](#page=16). * Les atomes d'hydrogène et de lithium, recherchant la configuration de l'hélium, suivent la **règle du duet** en visant 2 électrons sur leur couche externe [16](#page=16). ##### 4.1.1.1.2 Structure de Lewis La représentation d'une molécule selon Lewis suit ces étapes : 1. **Calcul du nombre total d'électrons périphériques:** Compter le nombre total d'électrons de valence ($n$) de tous les atomes. $n$ doit être un nombre pair. Le nombre de doublets à répartir est $n/2$ [16](#page=16). 2. **Répartition des doublets:** Respecter la règle de l'octet pour chaque atome (sauf H) [16](#page=16). 3. **Position des atomes:** Placer l'atome le moins électronégatif au centre et les atomes d'hydrogène à l'extérieur [16](#page=16). 4. **Choix de la structure la plus probable:** Privilégier la structure avec les charges formelles les plus petites, idéalement proches de zéro [16](#page=16). La **charge formelle (C.F.)** d'un atome est calculée par la formule : $$C.F. = V - (E + L)$$ où $V$ est le nombre d'électrons de valence, $E$ le nombre d'électrons non liants (libres), et $L$ le nombre de liaisons formées par l'atome [16](#page=16). > **Exemple:** Pour la molécule $F_2$, la charge formelle de chaque atome de fluor est $7 - (6 + 1) = 0$ [16](#page=16). #### 4.1.1.2 Approximation L.C.A.O. (Combinaison Linéaire des Orbitales Atomiques) Cette théorie quantique décrit les électrons de liaison covalente comme occupant des orbitales moléculaires, formées par combinaison linéaire des orbitales atomiques externes des atomes impliqués. Si $\Psi_{SA}$ et $\Psi_{SB}$ sont les orbitales atomiques, l'orbitale moléculaire $\Psi_{AB}$ est donnée par [17](#page=17): $$\Psi_{AB} = a(\Psi_{SA} + b\Psi_{SB})$$ Cette orbitale moléculaire ne peut contenir que deux électrons de spins opposés [17](#page=17). ##### 4.1.1.2.1 Formation des orbitales moléculaires À partir de deux orbitales atomiques, deux orbitales moléculaires sont formées : * **Orbitale Moléculaire Liante (O.M. liante) :** * Énergie plus faible [17](#page=17). * Formation constructive: $\Psi_{AB} = a(\Psi_{SA} + \Psi_{SB})$ [17](#page=17). * Augmente la densité électronique entre les noyaux, stabilisant la molécule [17](#page=17). * **Orbitale Moléculaire Antiliante (O.M. antiliante) :** * Énergie plus élevée [17](#page=17). * Formation destructive: $\Psi_{AB} = a(\Psi_{SA} - \Psi_{SB})$ [17](#page=17). * Diminue la densité électronique entre les noyaux, déstabilisant la molécule [17](#page=17). Une représentation schématique de ces niveaux d'énergie est souvent utilisée [18](#page=18). ##### 4.1.1.2.2 Orbitale Moléculaire $\sigma$ Les orbitales moléculaires $\sigma$ résultent du recouvrement axial d'orbitales atomiques dont les axes de symétrie sont colinéaires ou confondus [18](#page=18). #### 4.1.1.3 Configuration électronique des molécules diatomiques Seuls les électrons de valence sont décrits par les orbitales moléculaires dans une liaison covalente. L'ordre général des orbitales moléculaires pour les molécules diatomiques de type $A_2$ est [19](#page=19): $(\sigma_{ns}) (\sigma_{ns}^*) (\pi_{np}) (\sigma_{np}) (\pi_{np}^*) (\sigma_{np}^*)$ [19](#page=19). Chaque orbitale moléculaire peut accueillir au maximum deux électrons de spins antiparallèles [19](#page=19). > **Exemple:** La configuration électronique de $N_2$ est $(\sigma_{2s})^2 (\sigma_{2s}^*)^2 (\pi_{2p})^4 (\sigma_{2p})^2 (\pi_{2p}^*)^0 (\sigma_{2p}^*)^0$ (#page=19, page=20) [19](#page=19) [20](#page=20). Certaines molécules, comme $O_2$ et $F_2$, présentent une interversion des niveaux $\sigma_{2p}$ et $\pi_{2p}$ dans leur configuration électronique [20](#page=20). Pour $O_2$: $(\sigma_{2s})^2 (\sigma_{2s}^*)^2 (\sigma_{2p})^2 (\pi_{2p})^4 (\pi_{2p}^*)^2 (\sigma_{2p}^*)^0$ [20](#page=20). Pour $F_2$: $(\sigma_{2s})^2 (\sigma_{2s}^*)^2 (\sigma_{2p})^2 (\pi_{2p})^4 (\pi_{2p}^*)^2 (\sigma_{2p}^*)^2$ [20](#page=20). La présence d'électrons célibataires dans les orbitales moléculaires confère des propriétés paramagnétiques à la molécule (attraction par un champ magnétique). Si tous les électrons sont appariés, la molécule est diamagnétique [20](#page=20). > **Exemple:** $O_2$ est paramagnétique, $F_2$ est diamagnétique [20](#page=20). ##### 4.1.1.3.1 Ordre de liaison L'ordre de liaison ($OL$) est une mesure de la stabilité de la liaison covalente, calculé comme suit : $$OL = \frac{\text{Nombre d'électrons liants} - \text{Nombre d'électrons antiliants}}{2}$$ Un ordre de liaison plus élevé implique une molécule plus stable, une énergie de dissociation plus grande et une longueur de liaison plus courte [21](#page=21). | Molécule | OL | Longueur de liaison (Å) | Énergie de dissociation (Kcal) | | :------- | :-: | :---------------------- | :----------------------------- | | $F_2$ | 1 | 1,42 | 32 | | $O_2$ | 2 | 1,21 | 114 | | $N_2$ | 3 | 1,10 | 225 | L'ordre de liaison n'est pas nécessairement un entier (ex: $O_2^+ \implies OL=2,5$). Un $OL=0$ indique que la molécule n'existe pas. Il ne faut pas confondre l'énergie de dissociation ($Ed$) avec l'énergie d'ionisation ($Ei$). La théorie L.C.A.O. s'applique également aux molécules diatomiques hétéronucléaires ($AB$) comme $CO$, $NO$, $HF$ [21](#page=21). ##### 4.1.1.3.2 Polarisation des liaisons covalentes Si une liaison covalente unit deux atomes de natures différentes (électronégativités différentes), la densité électronique est répartie de manière non symétrique, favorisant l'atome le plus électronégatif. Cela crée un excès de charge négative sur un atome et un déficit sur l'autre [21](#page=21). #### 4.1.2 Liaison ionique La liaison ionique se forme par transfert d'un ou plusieurs électrons d'un atome peu électronégatif vers un atome très électronégatif (différence d'électronégativité > 2). Cette interaction électrostatique non directive conduit à la formation de structures cristallines [26](#page=26). Les conditions favorisant la liaison ionique sont : * Une différence d'électronégativité importante [26](#page=26). * Des énergies d'ionisation et affinités électroniques faibles [26](#page=26). Les éléments des groupes IA et VIIA ont une forte tendance à former des liaisons ioniques entre eux [26](#page=26). > **Exemple:** $Na^+$ et $Cl^-$ dans le $NaCl$ [26](#page=26). ### 4.2 Géométrie des molécules Les molécules covalentes adoptent une géométrie tridimensionnelle bien définie, caractérisée par des distances interatomiques et des angles de liaison précis [22](#page=22). #### 4.2.1 Hybridation des orbitales atomiques Pour expliquer les angles de liaison observés expérimentalement, qui ne correspondent pas toujours aux angles des orbitales p pures (90°), le concept d'hybridation des orbitales atomiques a été introduit par Pauling. Avant le recouvrement, les orbitales atomiques externes d'un atome se mélangent pour former des orbitales hybrides équivalentes, orientées dans des directions spécifiques [23](#page=23). ##### 4.2.1.1 Hybridation $sp^3$ * Mélange d'une orbitale $s$ et de trois orbitales $p$ [23](#page=23). * Forme: 4 orbitales hybrides $sp^3$ identiques, orientées selon les sommets d'un tétraèdre [23](#page=23). * Angle: environ 109° [23](#page=23). * Géométrie associée: Tétraédrique [23](#page=23). > **Exemple:** Méthane ($CH_4$). Le carbone utilise 4 OA hybrides $sp^3$ pour former 4 liaisons $\sigma$ avec les atomes d'hydrogène [23](#page=23). ##### 4.2.1.2 Hybridation $sp^2$ * Mélange d'une orbitale $s$ et de deux orbitales $p$ [24](#page=24). * Forme: 3 orbitales hybrides $sp^2$ identiques, situées dans un plan. Une orbitale $p$ reste non hybridée et perpendiculaire à ce plan [24](#page=24). * Angle: environ 120° [24](#page=24). * Géométrie associée: Triangulaire plane [24](#page=24). > **Exemple:** Éthène ($C_2H_4$). Chaque atome de carbone utilise 3 OA hybrides $sp^2$ pour former des liaisons $\sigma$ (deux avec H, une avec l'autre C) et l'orbitale $p$ non hybridée forme une liaison $\pi$ entre les carbones [24](#page=24). ##### 4.2.1.3 Hybridation $sp$ * Mélange d'une orbitale $s$ et d'une orbitale $p$ [24](#page=24). * Forme: 2 orbitales hybrides $sp$ identiques, orientées linéairement. Deux orbitales $p$ restent non hybridées et perpendiculaires entre elles et à l'axe des orbitales $sp$ [24](#page=24) [25](#page=25). * Angle: environ 180° [24](#page=24). * Géométrie associée: Linéaire [24](#page=24). > **Exemple:** Éthyne ($C_2H_2$). Chaque atome de carbone utilise 2 OA hybrides $sp$ pour former des liaisons $\sigma$ (une avec H, une avec l'autre C) et les deux orbitales $p$ non hybridées forment deux liaisons $\pi$ entre les carbones [25](#page=25). > **Remarque 1:** Le nombre d'orbitales hybrides formées est égal au nombre d'orbitales atomiques mélangées [25](#page=25). > **Remarque 2:** Pour les atomes possédant des orbitales $d$ dans leur couche externe, des hybridations impliquant les orbitales $d$ sont possibles [25](#page=25). #### 4.2.2 Détermination de l'état d'hybridation Pour déterminer l'hybridation d'un atome central : 1. Représenter la molécule selon Lewis [25](#page=25). 2. Calculer le **nombre de directions** : Nombre de directions = nombre de liaisons $\sigma$ + nombre de doublets non liants [25](#page=25). * 2 directions $\implies$ hybridation $sp$ (linéaire, 180°) [25](#page=25). * 3 directions $\implies$ hybridation $sp^2$ (triangulaire, 120°) [25](#page=25). * 4 directions $\implies$ hybridation $sp^3$ (tétraédrique, 109°) [25](#page=25). L'angle des liaisons formé par l'atome considéré peut également indiquer son état d'hybridation [25](#page=25). | Molécule | Atome central | Hybridation | Nbr de directions | Doublets non liants | Géométrie | | :------- | :------------ | :---------- | :---------------- | :------------------ | :------------ | | $BeH_2$ | Be | $sp$ | 2 | 0 | Linéaire | | $C_2H_4$ | C | $sp^2$ | 3 | 0 | Triangulaire | | $NO_3^-$ | N | $sp^2$ | 3 | 0 | Triangulaire | | $CH_4$ | C | $sp^3$ | 4 | 0 | Tétraédrique | | $NH_3$ | N | $sp^3$ | 4 | 1 | Pyramidale | | $H_2O$ | O | $sp^3$ | 4 | 2 | En V | --- ## Erreurs courantes à éviter - Révisez tous les sujets en profondeur avant les examens - Portez attention aux formules et définitions clés - Pratiquez avec les exemples fournis dans chaque section - Ne mémorisez pas sans comprendre les concepts sous-jacents

| Term | Definition | |------|------------| | Atome | La plus petite unité de matière conservant les propriétés chimiques d'un élément, composée d'un noyau central et d'électrons. | | Noyau | Partie centrale de l'atome contenant des protons et des neutrons, et concentrant la quasi-totalité de sa masse. | | Neutron | Particule élémentaire du noyau atomique de charge électrique nulle et de masse approximativement égale à celle du proton. | | Proton | Particule élémentaire du noyau atomique portant une charge électrique positive élémentaire ($+e$). | | Électron | Particule élémentaire chargée négativement ($ -e $) gravitant autour du noyau atomique dans des régions spécifiques appelées orbitales. | | Numéro atomique (Z) | Nombre de protons dans le noyau d'un atome, qui détermine l'élément chimique. Pour un atome neutre, il est égal au nombre d'électrons. | | Nombre de masse (A) | Somme du nombre de protons (Z) et du nombre de neutrons (N) dans le noyau d'un atome ($ A = Z + N $). | | Unité de masse atomique (u.m.a.) | Unité de mesure de masse, définie comme 1/12 de la masse d'un atome de Carbone 12. ($1 \text{ u.m.a.} \approx 1,67 \times 10^{-27} \text{ Kg}$). | | Isotope | Atomes d'un même élément chimique (même Z) mais possédant un nombre différent de neutrons (donc un A différent). | | Photon | Quantum d'énergie associé au rayonnement électromagnétique, dont l'énergie est donnée par $ E = h\nu $, où $h$ est la constante de Planck et $\nu$ est la fréquence. | | Constante de Planck (h) | Constante fondamentale de la physique quantique, d'une valeur de $6,62 \times 10^{-34} \text{ J.s.}$. | | Absorption d'énergie | Processus par lequel un électron, en passant d'un niveau d'énergie inférieur à un niveau supérieur, acquiert la différence d'énergie sous forme de photon. | | Émission d'énergie | Processus par lequel un électron, en passant d'un niveau d'énergie supérieur à un niveau inférieur, libère la différence d'énergie sous forme de photon. | | État fondamental | État d'énergie le plus bas d'un électron dans un atome, correspondant au niveau quantique principal $n=1$. | | État excité | État d'énergie supérieur à l'état fondamental, atteint par un électron après absorption d'énergie. | | Ionisation | Processus par lequel un atome ou une molécule perd ou gagne des électrons, devenant ainsi un ion. L'énergie d'ionisation est l'énergie nécessaire pour arracher un électron. | | Spectre d'émission | Ensemble des longueurs d'onde de la lumière émise par un élément lorsqu'il est excité, formant des raies caractéristiques. | | Constante de Rydberg ($R_H$) | Constante physique utilisée dans les formules décrivant les spectres atomiques, notamment pour l'hydrogène ($R_H \approx 1,1 \times 10^7 \text{ m}^{-1}$). | | Nombre quantique principal (n) | Nombre entier définissant la couche électronique de l'atome, lié au niveau d'énergie principal de l'électron. | | Nombre quantique secondaire ou azimutal (l) | Nombre entier définissant la forme de l'orbitale atomique, lié à la sous-couche électronique (s, p, d, f). | | Nombre quantique magnétique ($m_l$) | Nombre entier définissant l'orientation spatiale de l'orbitale atomique au sein d'une sous-couche. | | Nombre quantique de spin ($s$) | Nombre qui décrit le moment cinétique intrinsèque de l'électron, prenant les valeurs $+1/2$ ou $-1/2$. | | Configuration électronique | Répartition des électrons d'un atome dans ses orbitales atomiques dans son état fondamental. | | Principe d'exclusion de Pauli | Principe stipulant que deux électrons dans un même atome ne peuvent avoir le même ensemble de quatre nombres quantiques. Dans une orbitale, il ne peut y avoir que deux électrons de spins opposés. | | Règle de Hund | Règle qui stipule que dans une sous-couche de plusieurs orbitales de même énergie, les électrons occupent d'abord le maximum d'orbitales avec des spins parallèles avant d'être appariés. | | Règle de Klechkowski | Règle empirique qui détermine l'ordre de remplissage des orbitales atomiques dans un atome, basé sur la somme $n+l$. | | Atome polyélectronique | Atome possédant plusieurs électrons (plus d'un). | | Charge effective ($Z^*$) | Charge nucléaire ressentie par un électron dans un atome polyélectronique, tenant compte de l'effet d'écran des autres électrons. | | Effet d'écran | Phénomène par lequel les électrons internes d'un atome atténuent l'attraction du noyau sur les électrons externes. | | Tableau périodique | Organisation systématique des éléments chimiques selon leurs numéros atomiques et propriétés périodiques, regroupés en périodes, groupes et blocs. | | Période | Rangée horizontale du tableau périodique, correspondant à un niveau d'énergie principal $n$. | | Groupe | Colonne verticale du tableau périodique, regroupant des éléments ayant des configurations électroniques de valence similaires. | | Bloc (s, p, d, f) | Division du tableau périodique basée sur la sous-couche électronique en cours de remplissage. | | Rayon atomique | Distance mesurée entre les noyaux de deux atomes liés par une liaison covalente simple, divisée par deux. | | Énergie d'ionisation ($E_i$) | Énergie minimale requise pour arracher un électron d'un atome isolé à l'état gazeux. | | Affinité électronique (AE) | Énergie libérée ou absorbée lors de l'ajout d'un électron à un atome neutre à l'état gazeux pour former un ion négatif. | | Électronégativité ($\chi$) | Mesure de la tendance d'un atome à attirer les électrons vers lui lorsqu'il est impliqué dans une liaison chimique. | | Liaison chimique | Force d'attraction qui maintient les atomes ensemble dans une molécule ou un composé. | | Électrons de valence | Électrons situés dans la couche électronique la plus externe d'un atome, impliqués dans la formation des liaisons chimiques. | | Liaison covalente | Liaison chimique formée par le partage d'une ou plusieurs paires d'électrons entre deux atomes. | | Liaison ionique | Liaison chimique formée par l'attraction électrostatique entre des ions de charges opposées, résultant d'un transfert d'électrons. | | Théorie de Lewis | Modèle de liaison chimique basé sur la représentation des électrons de valence par des points et la formation de doublets liants. | | Règle de l'octet | Principe selon lequel les atomes tendent à gagner, perdre ou partager des électrons pour atteindre une configuration électronique stable avec 8 électrons de valence. | | Règle du duet | Principe similaire à la règle de l'octet, appliqué aux atomes d'hydrogène et de lithium, qui tendent à atteindre 2 électrons de valence pour imiter la configuration de l'hélium. | | Charge formelle (C.F.) | Charge calculée attribuée à un atome dans une structure de Lewis, en supposant que les électrons de liaison sont partagés également. | | Approximation L.C.A.O. (Linear Combination of Atomic Orbitals) | Méthode pour construire des orbitales moléculaires par combinaison linéaire d'orbitales atomiques. | | Orbitale moléculaire liante | Orbitale moléculaire de basse énergie formée par combinaison constructive d'orbitales atomiques, stabilisant la molécule. | | Orbitale moléculaire antiliante | Orbitale moléculaire de haute énergie formée par combinaison destructive d'orbitales atomiques, déstabilisant la molécule. | | Orbitale moléculaire $\sigma$ | Orbitale moléculaire formée par recouvrement axial d'orbitales atomiques. | | Ordre de liaison | Mesure de la force et de la stabilité d'une liaison covalente, calculée comme la moitié de la différence entre le nombre d'électrons dans les orbitales liantes et antiliantes. | | Paramagnétique | Qualifie une substance attirée par un champ magnétique, due à la présence d'électrons non appariés. | | Diamagnétique | Qualifie une substance faiblement repoussée par un champ magnétique, due à l'absence d'électrons non appariés. | | Hybridation | Mélange d'orbitales atomiques d'un même atome pour former de nouvelles orbitales hybrides avec des formes et orientations spécifiques, permettant d'expliquer la géométrie moléculaire. | | Hybridation sp3 | Hybridation résultant du mélange d'une orbitale s et de trois orbitales p, formant quatre orbitales hybrides dirigées vers les sommets d'un tétraèdre. | | Hybridation sp2 | Hybridation résultant du mélange d'une orbitale s et de deux orbitales p, formant trois orbitales hybrides dans un plan avec un angle de 120°. | | Hybridation sp | Hybridation résultant du mélange d'une orbitale s et d'une orbitale p, formant deux orbitales hybrides linéaires avec un angle de 180°. | | Géométrie moléculaire | Arrangement tridimensionnel des atomes dans une molécule, déterminé par la théorie VSEPR et l'hybridation des orbitales. | | Doublet non liant | Paire d'électrons de valence sur un atome qui n'est pas impliquée dans une liaison chimique. |

# Samenstelling en karakteristieken van het atoom Dit onderdeel behandelt de fundamentele opbouw van atomen, de eigenschappen van elementaire deeltjes en sleutelconcepten zoals atoomnummer, massagetal en isotopen. ### 1.1 De opbouw van een atoom Een atoom is opgebouwd uit een positieve kern en negatief geladen elektronen die daaromheen bewegen. De kern zelf bevat protonen (p+) en neutronen (n0). Protonen, neutronen en elektronen worden beschouwd als elementaire deeltjes [3](#page=3). ### 1.2 Karakteristieken van elementaire deeltjes De elementaire deeltjes hebben specifieke eigenschappen met betrekking tot hun diameter, lading en massa [4](#page=4). | Elementair deeltje | Diameter (nm) | Lading (Q) Absolute (C) | Lading (Q) Relatieve | Massa (m) Absolute (kg) | Massa (m) Relatieve (u) | |---|---|---|---|---|---| | elektron (e-) | - | -1,6.10⁻¹⁹ | -1 | 9,11.10⁻³¹ | 5,485803 . 10⁻⁴ | | proton (p+) | 10⁻⁶ | +1,6.10⁻¹⁹ | +1 | 1,673.10⁻²⁷ | 1,007276 | | neutron (n⁰) | 10⁻⁶ | 0 | 0 | 1,675.10⁻²⁷ | 1,008665 | > **Tip:** Let op de relatieve waarden voor lading en massa, deze worden vaak gebruikt voor vereenvoudigde berekeningen [4](#page=4). ### 1.3 Karakteristieken van een atoom #### 1.3.1 Atoomnummer en massagetal #### 1.3.1.1 Atoomnummer (Z) Het atoomnummer ($Z$) vertegenwoordigt het rangnummer van een element en is gelijk aan het aantal protonen in de kern. Dit aantal bepaalt de identiteit van een element en dus ook zijn plaats in het periodiek systeem. Het aantal elektronen in een neutraal atoom is gelijk aan het atoomnummer [6](#page=6) [9](#page=9). #### 1.3.1.2 Massagetal (A) Het massagetal ($A$) is de som van het aantal protonen en neutronen in de kern van een atoom. Het geeft een indicatie van de massa van het atoom. Het aantal neutronen kan berekend worden met de formule: aantal neutronen = $A - Z$ [6](#page=6). #### 1.3.1.3 Symbool van een atoom Een atoom wordt symbolisch weergegeven met de notatie: $$ \ce{^{A}_{Z}X} $$ waarbij $X$ het symbool van het element is, $A$ het massagetal en $Z$ het atoomnummer [6](#page=6). #### 1.3.2 Isotopen Isotopen zijn atoomsoorten van hetzelfde element die hetzelfde aantal protonen ($Z$) hebben, maar een verschillend aantal neutronen ($n^0$). Hierdoor hebben isotopen van een element ook een verschillend massagetal ($A$) [8](#page=8) [9](#page=9). * **Kenmerken van isotopen:** * Gelijk aantal protonen ($Z$) [9](#page=9). * Verschillend aantal neutronen ($n^0$) [8](#page=8). * Verschillend massagetal ($A$) [8](#page=8). * Gelijke chemische eigenschappen vanwege hetzelfde aantal elektronen [8](#page=8) [9](#page=9). * Verschillende massa [9](#page=9). * Komen in natuurlijke verhoudingen voor, de zogenaamde abundantie [9](#page=9). > **Voorbeeld:** Waterstof kent verschillende isotopen zoals deuterium en tritium. Chloor komt voor als twee natuurlijke isotopen: $^{35}\text{Cl}$ (75,5% abundantie) en $^{37}\text{Cl}$ (24,5% abundantie) [8](#page=8) [9](#page=9). #### 1.3.3 Absolute atoommassa De absolute atoommassa is de werkelijke massa van een individueel atoom. Deze massa is de som van de absolute massa's van de protonen, neutronen en elektronen in het atoom. De absolute atoommassa kan uitgedrukt worden in kilogram ($kg$) of in de atoommassa-eenheid ($u$) [10](#page=10). > **Voorbeeld:** De absolute massa van een $^{16}\text{O}$ atoom wordt berekend als: > $m(\text{ }^{16}\text{O}) = 8 \cdot m_p + 8 \cdot m_n + 8 \cdot m_e$ > $= 8 \cdot (1,673 \cdot 10^{-27} \text{ kg}) + 8 \cdot (1,675 \cdot 10^{-27} \text{ kg}) + 8 \cdot (9,11 \cdot 10^{-31} \text{ kg})$ > $= 2,679 \cdot 10^{-26} \text{ kg}$ [10](#page=10). #### 1.3.4 Relatieve atoommassa De relatieve atoommassa ($A_r$) is een verhoudingsgetal dat de massa van een atoom vergelijkt met de atoommassa-eenheid ($u$). Het wordt gebruikt in plaats van de absolute atoommassa om de berekeningen te vereenvoudigen [11](#page=11). $$ A_r = \frac{m_{\text{atoom}}}{u} $$ > **Voorbeeld:** De relatieve atoommassa van een $^7\text{Li}$ atoom: > $m(^7\text{Li}) = 3 \cdot m_p + 4 \cdot m_n + 3 \cdot m_e$ > $= 3 \cdot (1,673 \cdot 10^{-27} \text{ kg}) + 4 \cdot (1,675 \cdot 10^{-27} \text{ kg}) + 3 \cdot (9,11 \cdot 10^{-31} \text{ kg})$ > $= 1,17 \cdot 10^{-26} \text{ kg}$ > > $A_r = \frac{1,17 \cdot 10^{-26} \text{ kg}}{1,66 \cdot 10^{-27} \text{ kg}} \approx 7,06$ [11](#page=11). **Belangrijke Opmerking:** De relatieve atoommassa van een element zoals weergegeven in het periodiek systeem is een *gemiddelde* waarde. Dit gemiddelde wordt berekend door rekening te houden met de relatieve atoommassa's van de natuurlijke isotopen en hun respectievelijke abundanties (procentueel voorkomen) [12](#page=12). > **Voorbeeld:** Voor chloor ($Cl$) met isotopen $^{35}\text{Cl}$ ($A_r \approx 34,97$) en $^{37}\text{Cl}$ ($A_r \approx 36,96$), en hun natuurlijke voorkomen van respectievelijk 75,5% en 24,5%, is de gemiddelde relatieve atoommassa: > $(0,755 \cdot 34,97) + (0,245 \cdot 36,96) \approx 35,45$. Dit is de waarde die men in het PSE terugvindt [12](#page=12). #### 1.3.5 Massadefect Het massadefect is het verschil tussen de berekende massa van een atoomkern (door de massa's van de individuele protonen en neutronen op te tellen) en de werkelijke gemeten massa van die kern [14](#page=14). * De massa's van de elementaire deeltjes in atoommassa-eenheden ($u$) zijn: * $m_e = 5,485803 \cdot 10^{-4} u$ * $m_p = 1,007276 u$ * $m_n = 1,008665 u$ [14](#page=14). > **Voorbeeld:** De berekende nuclidemassa van $^{10}\text{B}$: > $m(^{10}\text{B}) = 5 \cdot m_p + 5 \cdot m_n + 5 \cdot m_e$ > $= 5 \cdot (1,007276 u) + 5 \cdot (1,008665 u) + 5 \cdot (5,485803 \cdot 10^{-4} u)$ > $= 10,082448 u$ [14](#page=14). > > De werkelijke massa van $^{10}\text{B}$ is echter $10,0129 u$. Het verschil ($10,082448 u - 10,0129 u = 0,069548 u$) is het massadefect [14](#page=14). * **Vaststellingen met betrekking tot massadefect:** * Voor de meeste nucliden is de berekende nuclidemassa groter dan de werkelijke nuclidemassa; er is dus een massa verloren gegaan (massadefect). Een uitzondering hierop is waterstof ($^1H$) [15](#page=15). * Het massadefect neemt toe met het aantal nucleonen in de kern [15](#page=15). * Voor nucliden lichter dan of gelijk aan $^{12}\text{C}$, is de nuclidemassa kleiner dan het massagetal [15](#page=15). * Voor nucliden zwaarder dan $^{12}\text{C}$, is de nuclidemassa groter dan het massagetal [15](#page=15). > **Tip:** Het massadefect is een direct gevolg van de Einstein's massa-energie equivalentie ($E=mc^2$), waarbij een deel van de massa wordt omgezet in bindingsenergie die de kern bij elkaar houdt [14](#page=14). --- # Evolutie van atoommodellen en het kwantummechanisch model Dit gedeelte traceert de historische ontwikkeling van het atoommodel, beginnend bij vroege ideeën tot het kwantummechanisch en golfmechanisch model, inclusief atoomorbitalen en kwantumgetallen. ### 2.1 Vroege atoommodellen #### 2.1.1 Democritus Democritus definieerde het atoom als een "materiedeeltje dat niet meer gesplitst kan worden" [17](#page=17). #### 2.1.2 Dalton Dalton wordt beschouwd als de vader van het atoommodel, met de volgende grondbeginselen [18](#page=18): * Elementen bestaan uit atomen, en atomen zijn ondeelbaar [18](#page=18). * Atomen worden gekenmerkt door hun grootte en massa [18](#page=18). * Er bestaan eenvoudige verhoudingen van gehele getallen tussen de aantallen atomen van verschillende elementen in een verbinding [18](#page=18). * Atomen kunnen niet vernietigd worden of uit het niets ontstaan [18](#page=18). #### 2.1.3 Thomson Thomson deed ladingsexperimenten met kathodestraalbuizen. Uit de afbuigingen van de kathodestralen leidde hij de verhoudingen af van massa en lading van deeltjes. Hij formuleerde de hypothese van de aanwezigheid van negatief geladen deeltjes, die hij elektronen noemde. Aangezien atomen neutraal zijn en elektronen negatief geladen zijn, moest er ook positief geladen deeltjes aanwezig zijn [20](#page=20) [21](#page=21). Zijn model stelde dat het atoom een heterogene structuur heeft. De periferie bevat elektronen, negatief geladen deeltjes ingebed in een positieve grondmaterie, waardoor het totale atoom een neutraal deeltje is [22](#page=22). #### 2.1.4 Rutherford Rutherfords atoommodel beschreef het atoom als een ijle ruimte met een positieve kern waarrond elektronen bewegen. Alle massa bevindt zich in de kern. In een later stadium identificeerde hij de positieve deeltjes als protonen [23](#page=23). #### 2.1.5 Chadwick Chadwick breidde het atoommodel uit door te stellen dat de atoomkern zowel positief geladen deeltjes (protonen) als neutrale deeltjes (neutronen) bevat. Dit leidde tot het begrip 'isotoop': atomen met hetzelfde aantal protonen en elektronen, maar een verschillend aantal neutronen [24](#page=24). > **Tip:** Het fundamentele probleem dat bleef bestaan na deze modellen was de vraag waarom de elektronen niet op de kern vielen [25](#page=25). ### 2.2 Het Bohr-atoommodel Bohr introduceerde het idee dat elektronen welbepaalde energie-inhouden hebben. Slechts specifieke energie-inhouden kunnen worden opgeslagen, niet willekeurige waarden. Een elektron kan van het ene energieniveau naar een hoger niveau overgaan, en bij terugval naar een lager niveau straalt het elektron het energieverschil uit [26](#page=26). #### 2.2.1 Visuele voorstelling van Bohr's model Volgens Bohr bewegen elektronen op specifieke banen rond de kern, deze banen worden schillen genoemd. Afhankelijk van de schil hebben elektronen specifieke energieën, dit zijn de energieniveaus. De potentiële energie wordt bepaald door de afstand van de elektronen tot de kern (grotere afstand betekent hogere energie). De kinetische energie is afhankelijk van de bewegingstoestand en dus van de schil. De totale energie van een elektron is afhankelijk van de schil waarop het zich bevindt; hoe verder verwijderd van de kern, hoe meer energie [28](#page=28). #### 2.2.2 Bohr's voorstelling van het atoommodel Het aantal schillen is beperkt en genummerd vanaf de kern naar buiten. Schillen 1 tot en met 7 komen overeen met schillen K tot en met Q. Het aantal elektronen neemt toe van binnen naar buiten. Het maximum aantal elektronen per schil is $2n^2$, waarbij $n$ het hoofdkwantumgetal is [29](#page=29). * Baan $n=1$ (K schil): maximaal $2 \times 1^2 = 2$ elektronen [29](#page=29). * Baan $n=2$ (L schil): maximaal $2 \times 2^2 = 8$ elektronen [29](#page=29). * Baan $n=3$ (M schil): maximaal $2 \times 3^2 = 18$ elektronen [29](#page=29). * Baan $n=4$ (N schil): maximaal $2 \times 4^2 = 32$ elektronen [29](#page=29). * Voor banen $n>4$ is het maximum ook 32 elektronen [29](#page=29). ### 2.3 Verfijning van het Bohr-model: Sommerfeld en kwantumgetallen #### 2.3.1 Sommerfeld Sommerfeld stelde dat er veel meer energieniveaus aanwezig zijn dan aanvankelijk gedacht. Hij suggereerde dat het hoofdkwantumgetal $n$ verder wordt onderverdeeld in subniveaus, gekenmerkt door het nevenkwantumgetal $l$ [30](#page=30). * $l=0$ komt overeen met s-elektronen en kan maximaal 2 elektronen bevatten [30](#page=30). * $l=1$ komt overeen met p-elektronen en kan maximaal 6 elektronen bevatten [30](#page=30). * $l=2$ komt overeen met d-elektronen en kan maximaal 10 elektronen bevatten [30](#page=30). * $l=3$ komt overeen met f-elektronen en kan maximaal 14 elektronen bevatten [30](#page=30). De volgende tabel geeft een overzicht van de energieniveaus en hun maximale bezetting [31](#page=31): | Hoofd-kwantumgetal $n$ | Neven-kwantumgetal $l$ | Voorstelling | Aantal elektronen max. | | :--------------------- | :--------------------- | :----------- | :--------------------- | | $n=1$ | $l=0$ | $1s^2$ | 2 elektronen | | $n=2$ | $l=0$ | $2s^2$ | 8 elektronen | | | $l=1$ | $2p^6$ | | | $n=3$ | $l=0$ | $3s^2$ | 18 elektronen | | | $l=1$ | $3p^6$ | | | | $l=2$ | $3d^{10}$ | | | $n=4$ | $l=0$ | $4s^2$ | 32 elektronen | | | $l=1$ | $4p^6$ | | | | $l=2$ | $4d^{10}$ | | | | $l=3$ | $4f^{14}$ | | | $n=5$ | $l=0$ | $5s^2$ | 32 elektronen | | | $l=1$ | $5p^6$ | | | | $l=2$ | $5d^{10}$ | | | | $l=3$ | $5f^{14}$ | | | ... | ... | ... | ... | #### 2.3.2 Kwantumgetallen De kwantumgetallen beschrijven de energietoestand van een elektron in een atoom [32](#page=32). * **Hoofd-energieniveau ($n$)**: Dit komt overeen met de 7 verschillende schillen, genummerd vanaf de kern naar buiten (schil K t.e.m. Q). De kleinste waarde van $n$ heeft het laagste energieniveau [33](#page=33). | Schil | K | L | M | N | O | P | Q | | :---- | :-: | :-: | :-: | :-: | :-: | :-: | :-: | | Nummer ($n$) | 1 | 2 | 3 | 4 | 5 | 6 | 7 | * **Sub-energieniveau ($l$)**: Dit is het nevenkwantumgetal, met de hoogste waarde van $l=3$ [33](#page=33). * $l=0 \implies$ s-elektronen (max. 2) * $l=1 \implies$ p-elektronen (max. 6) * $l=2 \implies$ d-elektronen (max. 10) * $l=3 \implies$ f-elektronen (max. 14) Een subniveau is volledig bezet bij $(4l+2)$ elektronen, wat resulteert in $s^2$, $p^6$, $d^{10}$ en $f^{14}$ [33](#page=33). * **Magnetisch energieniveau ($m_l$)**: Elektronen hebben een beperkt aantal banen in elk subenergieniveau. Het aantal banen gaat van $-l$ tot $+l$ [36](#page=36). * Subniveau s: 1 baan [36](#page=36). * Subniveau p: 3 banen [36](#page=36). * Subniveau d: 5 banen [36](#page=36). * Subniveau f: 7 banen [36](#page=36). * **Spin ($m_s$)**: Per baan kunnen maximaal 2 elektronen voorkomen, die tegengestelde spins hebben en zo een doublet vormen [36](#page=36). > **Samenvatting van energieniveaus en bezetting:** Elektronen zijn verdeeld over energieniveaus die in energie verschillen (K, L, M, N,...). Elk hoofdenergieniveau bevat subenergieniveaus die in energie verschillen (s, p, d, f). Elk subniveau bevat een bepaald aantal banen, en elke baan bevat maximaal 2 elektronen [37](#page=37). | Hoofd-kwantumgetal $n$ | Neven-kwantumgetal $l$ | Magnetisch kwantumgetal $m_L$ | Voorstelling | Aantal banen | | :--------------------- | :--------------------- | :--------------------------- | :----------- | :----------- | | $n=1$ | $l=0$ | $m_l=0$ | $1s$ | Op subniveau s: 1 baan | | $n=2$ | $l=0$ | $m_l=0$ | $2s$ | Op subniveau p: 3 banen | | | $l=1$ | $m_l=-1, 0, +1$ | $2p_x, 2p_y, 2p_z$ | | | $n=3$ | $l=0$ | $m_l=0$ | $3s$ | Op subniveau d: 5 banen | | | $l=1$ | $m_l=-1, 0, +1$ | $3p_x, 3p_y, 3p_z$ | | | | $l=2$ | $m_l=-2, -1, 0, +1, +2$ | $3d_1, 3d_2, 3d_3, 3d_4, 3d_5$ | | | $n=4$ | $l=0$ | $m_l=0$ | $4s$ | Op subniveau f: 7 banen | | | $l=1$ | $m_l=-1, 0, +1$ | $4p_x, 4p_y, 4p_z$ | | | | $l=2$ | $m_l=-2, -1, 0, +1, +2$ | $4d_1, ..., 4d_5$ | | | | $l=3$ | $m_l=-3, ..., +3$ | $4f_1, ..., 4f_7$ | | | ... | ... | ... | ... | ... | ### 2.4 Het golfmechanisch atoommodel #### 2.4.1 Louis de Broglie Louis de Broglie suggereerde dat elektronen een tweevoudig karakter vertonen: een deeltjeskarakter (massa) en een golfkarakter (beweging) [39](#page=39). #### 2.4.2 Onzekerheidsprincipe van Heisenberg Volgens Heisenberg heeft een elektron geen vaste plaats, maar bestaat er een waarschijnlijkheid om zich daar te bevinden. Op plaatsen waar de waarschijnlijkheidswolk het dichtst is, is de kans het elektron aan te treffen het grootst. Het is onmogelijk om zowel de positie als de snelheid van een elektron exact te kennen [40](#page=40). #### 2.4.3 Schrödinger Schrödinger ontwikkelde een methode om de waarschijnlijkheid te berekenen om een elektron op een bepaalde plaats rond de atoomkern aan te treffen. Hij bakende gebieden af waar de kans om een elektron aan te treffen groot is, dit worden orbitalen genoemd. Een orbitaal is een "denkbeeldig" gebied waarbinnen de waarschijnlijkheid om het elektron aan te treffen 90% is [41](#page=41). ### 2.5 Atoomorbitalen #### 2.5.1 Types orbitalen Er zijn 4 types orbitalen, elk met een specifieke vorm en grootte [42](#page=42). * **Vorm**: s, p, d, en f orbitalen [42](#page=42). * **Grootte**: Een s orbitaal van schil K is kleiner dan een s orbitaal van schil L, wat wordt aangeduid als $1s < 2s$ [42](#page=42). * **s orbitaal**: Is bolvormig, heeft één oriëntatie en geen opsplitsing [43](#page=43). * **p orbitaal**: Is haltervormig met een knooppunt in de atoomkern en heeft 3 oriëntaties in een magnetisch veld: $p_x$, $p_y$, en $p_z$ [43](#page=43). * **d orbitaal**: Heeft ingewikkelde vormen en 5 oriëntaties [44](#page=44). * **f orbitaal**: Heeft ingewikkelde vormen en 7 oriëntaties [44](#page=44). > **Voorbeeld van kwantumgetallen en orbitalen:** > * $n=2; l=1; m_l=0$ duidt op één van de drie p-orbitalen ($l=1$) op het tweede niveau, bijvoorbeeld $2p_y$ [45](#page=45). > * $n=4; l=0; m_l=0$ duidt op een $4s$ orbitaal [45](#page=45). > * $n=3; l=2; m_l=+1$ duidt op een van de vijf d-orbitalen op het derde niveau [45](#page=45). > * $n=2; l=3; m_l=+2$ is een ongeldige combinatie, aangezien $l$ maximaal $n-1$ kan zijn [45](#page=45). > * $n=1; l=0; m_l=+1$ is een ongeldige combinatie, aangezien voor $l=0$ ($s$-orbitaal) $m_l$ alleen 0 kan zijn [45](#page=45). #### 2.5.2 Toepassingen van kwantumgetallen De volgende oefeningen illustreren de toepassing van kwantumgetallen [46](#page=46): 1. Bepaal de waarden van $n$, $l$, en $m_l$ voor: a) $1s$: $n=1, l=0, m_l=0$ b) $4d^2$: $n=4, l=2, m_l \in \{-2, -1, 0, 1, 2\}$ (specificeert één van de vijf $d$-orbitalen) c) $2s$: $n=2, l=0, m_l=0$ d) $3p_z$: $n=3, l=1, m_l=0$ (aangezien $p_z$ een specifieke oriëntatie binnen de $p$-orbitalen is) 2. Welke orbitalen worden aangeduid door: a) $n=4; l=2; m_l=-1$: Een $4d$ orbitaal (specifiek één van de vijf $d$-orbitalen) b) $n=3; l=1; m_l=0$: Een $3p$ orbitaal (specifiek $3p_y$ als we $m_l=0$ aan de $y$-as koppelen, of $3p_z$ als we deze aan de z-as koppelen, afhankelijk van de conventie). Het tekenen van elektronenwolken voor atomen zoals Neon volgens de modellen van Bohr, Sommerfeld en Schrödinger met verschillende kleuren is een waardevolle oefening om de concepten te visualiseren [46](#page=46). --- # Elektronenconfiguraties en het periodiek systeem Dit onderdeel behandelt de regels voor het opstellen van elektronenconfiguraties en de structuur van het periodiek systeem. ### 3.1 Elektronenconfiguraties Elektronenconfiguraties beschrijven de verdeling van elektronen over de orbitalen van een atoom. Ze worden weergegeven als een code waarbij het eerste cijfer het hoofdniveau (schil) aangeeft, de letter het subniveau (orbitaaltype) en het tweede cijfer het aantal elektronen in dat orbitaal [47](#page=47). #### 3.1.1 Notatie * **Standaardnotatie:** Een voorbeeld is 1s² 2s² 2p⁶ 3s¹ voor natrium (Na) [47](#page=47). * **Verkorte notatie:** Dit maakt gebruik van de elektronenconfiguratie van het voorgaande edelgas om de kernconfiguratie weer te geven, gevolgd door de configuratie van de valentie-elektronen. Voorbeeld: [Ne 3s¹ voor natrium (Na) [48](#page=48). * **Orbitaalrepresentatie:** Orbitalen kunnen worden voorgesteld als vakjes, waarbij elektronen worden weergegeven als pijltjes (↑ of ↓). Twee elektronen in hetzelfde orbitaal moeten tegengestelde spins hebben en vormen een doublet [47](#page=47). #### 3.1.2 Regels voor het opstellen van elektronenconfiguraties Er zijn drie fundamentele regels die de opbouw van elektronenconfiguraties bepalen: ##### 3.1.2.1 Regel van de minimale energie Atomen streven naar de meest stabiele toestand, wat overeenkomt met de laagst mogelijke energie. Dit betekent dat orbitalen met een lagere energie-inhoud eerst worden gevuld voordat orbitalen met een hogere energie worden bezet. De volgorde van opvulling wordt bepaald door de 'diagonaalregel' [49](#page=49): 1s → 2s → 2p → 3s → 3p → 4s → 3d → 4p → 5s → 4d → 5p → 6s → 4f → 5d → 6p → 7s → 5f → 6d → 7p [49](#page=49). ##### 3.1.2.2 Regel van Hund Binnen een subniveau met gelijksoortige orbitalen (bijvoorbeeld de p-orbitalen) zullen elektronen zich zoveel mogelijk ongepaard verspreiden met dezelfde spin. Dit resulteert in een maximaal aantal ongepaarde elektronen. Elektronen vullen eerst elk orbitaal van een subniveau één voor één voordat er paren worden gevormd, en ongepaarde elektronen hebben altijd dezelfde spin (meestal gesymboliseerd als ↑) [50](#page=50). ##### 3.1.2.3 Pauli-verbod Dit principe stelt dat geen twee elektronen in een atoom exact dezelfde set van vier kwantumgetallen kunnen hebben. Dit impliceert dat een orbitaal maximaal twee elektronen kan bevatten, en deze twee elektronen moeten tegengestelde spintoestanden hebben [51](#page=51). #### 3.1.3 Stabiliteitsregels Bepaalde elektronenconfiguraties leiden tot extra stabiliteit: * **Edelgasconfiguratie:** Een buitenste schil met acht elektronen (typisch s²p⁶) is zeer stabiel [52](#page=52). * **Volledig bezette subschil:** Wanneer alle orbitalen van een bepaald type binnen een subschil volledig bezet zijn (bijvoorbeeld een p⁶-configuratie), is dit energetisch stabieler dan een gedeeltelijk bezette subschil [52](#page=52). * **Half bezette subschil:** Wanneer alle orbitalen van een bepaald type binnen een subschil één elektron bevatten, is dit ook een stabiele configuratie [52](#page=52). #### 3.1.4 Uitzonderingen en speciale gevallen * **Inversies:** Soms wijken waargenomen elektronenconfiguraties af van de voorspelde configuraties door de stabiliteit van volledig of half bezette d-subschillen. Bijvoorbeeld, de verwachte configuratie voor koper (Cu) is 1s²2s²2p⁶3s²3p⁶4s²3d⁹, maar de waargenomen configuratie is 1s²2s²2p⁶3s²3p⁶4s¹3d¹⁰ [53](#page=53). * **Anionen:** Bij de vorming van een anion worden elektronen toegevoegd aan de buitenste schil. Bijvoorbeeld, voor zuurstof (O) met de configuratie 1s²2s²2p⁴, wordt het O²⁻ anion 1s²2s²2p⁶ [53](#page=53). * **Kationen:** Bij de vorming van een kation worden elektronen verwijderd, meestal uit de buitenste s-orbitalen eerst. Voor magnesium (Mg) met de configuratie 1s²2s²2p⁶3s², vormt Mg²⁺ de configuratie 1s²2s²2p⁶ [53](#page=53). #### 3.1.5 Toepassingen * Bepaal de elektronenconfiguratie van bepaalde elementen [54](#page=54). * Identificeer elementen in grondtoestand of aangeslagen toestand op basis van hun elektronenconfiguratie [54](#page=54). ### 3.2 Het periodiek systeem (PSE) Het periodiek systeem is een tabel die de elementen ordent op basis van hun atoomnummer, elektronenconfiguratie en terugkerende chemische eigenschappen [59](#page=59). #### 3.2.1 Opbouw: groepen en perioden * **Groepen (kolommen):** * **Hoofdgroepen (a-groepen):** Deze groepen (Ia tot O(nul)) delen dezelfde buitenste elektronenbezetting, wat leidt tot vergelijkbare chemische eigenschappen. De buitenste elektronen worden valentie-elektronen genoemd. De hoofdgroepen zijn: Ia (alkalimetalen), IIa (aardalkalimetalen), IIIa (aardmetalen), IVa (koolstofgroep), Va (stikstofgroep), VIa (zuurstofgroep), VIIa (halogenen), en O(nul)-groep (edelgassen) [61](#page=61). * **Nevengroepen (b-groepen) of transitie-elementen:** Verticale analogieën in deze groepen betekenen hetzelfde maximale aantal valentie-elektronen, maar elektronen kunnen verschillende orbitalen bezetten. Horizontale analogieën in dezelfde periode van b-groepen delen hetzelfde aantal s-elektronen op de buitenste schil [63](#page=63). * **Inner-transitie-elementen (c-groepen):** Dit zijn de lanthaniden en actiniden, die verschillen in de opvulling van f-orbitalen [63](#page=63). * **Perioden (rijen):** Elementen in dezelfde periode hebben hetzelfde aantal schillen. Het nummer van de periode correspondeert met het aantal bezette schillen [63](#page=63). #### 3.2.2 Eigenschappen van hoofdelementen Elementen uit de hoofgroepen kunnen worden onderverdeeld in: * **Metalen:** Gelegen links onder de 'trapvorm' in het PSE. Ze geleiden warmte en vormen meestal glanzende vaste stoffen bij kamertemperatuur. Metaalkarakter neemt toe van rechts naar links en van boven naar onder [62](#page=62). * **Niet-metalen:** Gelegen rechtsboven in de tabel, plus waterstof. Ze zijn gassen of broze vaste stoffen en slechte geleiders. Niet-metaalkarakter neemt toe van links naar rechts en van onder naar boven [62](#page=62). * **Elementen met zowel metaal- als niet-metaalkarakter:** Bevinden zich langs de scheidingslijn tussen metalen en niet-metalen [62](#page=62). #### 3.2.3 Opbouw: blokken Het periodiek systeem kan ook worden ingedeeld in blokken op basis van het subniveau dat als laatste wordt gevuld: * **s- en p-blok:** Bevatten de hoofdelementen; het laatste gevulde subniveau is een s- of p-subniveau [64](#page=64). * **d-blok:** Bevat de overgangselementen; het laatste gevulde subniveau is een d-subniveau [64](#page=64). * **f-blok:** Bevat de lanthaniden en actiniden; het laatste gevulde subniveau is een f-subniveau [64](#page=64). * **Edelgassen:** Kenmerken zich door een volledig gevuld s- en p-subniveau in de buitenste schil (s²p⁶-configuratie), wat de edelgasconfiguratie of octetconfiguratie wordt genoemd [64](#page=64). #### 3.2.4 Toepassingen * Het bepalen van de positie van elementen in dezelfde periode of groep en hun correlatie met het aantal valentie-elektronen [65](#page=65). * Het herkennen van edelgassen [65](#page=65). > **Tip:** Begrijp hoe de elektronenconfiguratie de plaats van een element in het periodiek systeem bepaalt, en omgekeerd. Dit is cruciaal voor het voorspellen van chemische eigenschappen. > > **Tip:** Oefen veel met het opstellen van elektronenconfiguraties, zowel de standaard- als de verkorte notatie, en wees alert op de stabiliteitsuitzonderingen. > > **Tip:** Verbindingen tussen groepen, perioden, blokken en de bijbehorende eigenschappen van elementen zijn fundamenteel voor het begrip van het periodiek systeem. --- # Periodieke eigenschappen van elementen Dit deel onderzoekt de periodieke trends in eigenschappen zoals effectieve kernlading, atoomstraal, ionisatie-energie, elektronenaffiniteit en elektronegativiteit, en de concepten van oxidatiegetallen. ### 4.1 Effectieve kernlading ($Z_{eff}$) De effectieve kernlading ($Z_{eff}$) is de netto positieve lading die een elektron ervaart van de kern. Deze is altijd kleiner dan de werkelijke kernlading ($Z$) omdat de aantrekking van de kern wordt afgeschermd door de elektronen in de binnenste schillen. De relatie wordt beschreven door de formule $Z_{eff} = Z - S$, waarbij $Z$ het aantal protonen in de kern is en $S$ het aantal elektronen op dieper gelegen schillen. Binnen een periode en een groep met toenemend atoomnummer neemt de effectieve kernlading toe. Bij transitie-elementen blijft de $Z_{eff}$ nagenoeg constant [66](#page=66). > **Tip:** Houd er rekening mee dat $Z_{eff}$ conceptueel nuttig is voor het begrijpen van trends, maar de $S$-waarde (afschermingsconstante) is niet altijd eenvoudig te berekenen zonder specifieke regels (zoals de regels van Slater). ### 4.2 Grootte van de atomen (atoomstraal) De atoomstraal is een maat voor de grootte van een atoom. Er zijn verschillende factoren die de atoomstraal beïnvloeden, zoals de kernlading, het aantal schillen en de afscherming door andere elektronen [68](#page=68). #### 4.2.1 Trends in groepen Binnen een groep neemt de atoomstraal toe van boven naar onder. Dit komt doordat het aantal schillen toeneemt, wat resulteert in een grotere afstand tussen de valentie-elektronen en de kern. Hoewel de kernlading ook toeneemt, overheersen de effecten van het grotere aantal schillen en de toegenomen afscherming van binnenste elektronen [68](#page=68). #### 4.2.2 Trends in periodes Binnen een periode nemen de hoofdgroepen af van links naar rechts. Dit komt voornamelijk door de toenemende kernlading die de valentie-elektronen sterker aantrekt, terwijl het aantal schillen gelijk blijft. Bij de nevengroepen (transitie-elementen) blijft de atoomstraal nagenoeg constant, omdat de extra elektronen worden toegevoegd aan een interne d-orbitaal, wat de afscherming van de buitenste schil beïnvloedt [69](#page=69). #### 4.2.3 Ionenstralen * **Kationen (positieve ionen):** De straal van kationen is kleiner dan die van hun neutrale atomen. Dit komt doordat er minder elektronen zijn, waardoor de aantrekkingskracht van de kern per elektron groter wordt, en er mogelijk ook een schil minder is. Bijvoorbeeld, de straal van $\text{Na}^+$ is kleiner dan die van $\text{Na}$ [70](#page=70). * **Anionen (negatieve ionen):** De straal van anionen is groter dan die van hun neutrale atomen. Dit komt doordat er meer elektronen zijn, wat leidt tot een grotere afstoting tussen de elektronen en een minder efficiënte afscherming van de kern. Bijvoorbeeld, de straal van $\text{F}^-$ is groter dan die van $\text{F}$ [70](#page=70). ### 4.3 Ionisatie-energie (I.E.) De ionisatie-energie is de energie die nodig is om een elektron volledig te onttrekken aan een atoom of ion in de gasfase. Dit proces kan als volgt worden weergegeven [71](#page=71): $$ \text{atoom} + \text{I.E.} \rightarrow \text{positief ion} + \text{e}^- $$ [71](#page=71). De ionisatie-energie wordt uitgedrukt in elektronvolt (eV) per atoom of kilojoule per mol (kJ/mol). Er zijn opeenvolgende ionisatie-energieën: I.E.1 is de energie om het eerste elektron te verwijderen, I.E.2 om het tweede, enzovoort. De opeenvolgende ionisatie-energieën nemen toe naarmate meer elektronen worden onttrokken [71](#page=71) [72](#page=72). > **Voorbeeld:** De eerste ionisatie-energie van Lithium (Li) om $\text{Li}^+$ te vormen is lager dan de tweede ionisatie-energie van Beryllium (Be) om $\text{Be}^{2+}$ te vormen, omdat het verwijderen van een elektron uit een stabiele edelgasconfiguratie (zoals na het verwijderen van 2 elektronen uit Be) veel meer energie vereist. #### 4.3.1 Factoren die I.E. beïnvloeden De ionisatie-energie hangt af van: * **Afstand tot de kern:** Hoe verder het elektron van de kern is, hoe kleiner de aantrekkingskracht en hoe lager de I.E. [72](#page=72). * **Kernlading:** Een hogere kernlading leidt tot een sterkere aantrekking en dus een hogere I.E. [72](#page=72). * **Hoofdkwantumgetal ($n$):** Een hoger hoofdkwantumgetal betekent dat het elektron zich op een grotere schil bevindt, wat resulteert in een lagere I.E. [72](#page=72). * **Afscherming:** Sterkere afscherming door binnenste elektronen vermindert de aantrekkingskracht van de kern op het valentie-elektron, wat leidt tot een lagere I.E. [72](#page=72). #### 4.3.2 Trends in periodes Binnen een periode neemt de eerste ionisatie-energie toe als het atoomnummer ($Z$) stijgt. Dit komt door de toenemende kernlading. Echter, er zijn afwijkingen door speciale elektronenconfiguraties [74](#page=74): * De I.E. van Stikstof (N) is hoger dan die van Zuurstof (O). Dit komt doordat Stikstof een half gevuld p-orbitaal heeft ($2p^3$), wat stabieler is dan de $2p^4$ configuratie van Zuurstof. Het verwijderen van een elektron uit Zuurstof leidt tot de stabielere configuratie van Stikstof [74](#page=74). * De I.E. van Magnesium (Mg) is hoger dan die van Aluminium (Al). Bij Magnesium wordt een elektron verwijderd uit een volledig gevuld 3s-subniveau ($3s^2$), wat minder gunstig is dan het verwijderen van een elektron uit het 3p-subniveau van Aluminium ($3p^1$) [74](#page=74). #### 4.3.3 Trends in groepen Binnen een groep neemt de eerste ionisatie-energie af als het atoomnummer ($Z$) stijgt. Hoewel de kernlading toeneemt, wegen de effecten van een groter hoofdkwantumgetal en een grotere afscherming zwaarder [75](#page=75). ### 4.4 Elektronenaffiniteit (E.A.) Elektronenaffiniteit is de energieverandering die optreedt wanneer een elektron wordt toegevoegd aan een gasvormig atoom om een negatief ion te vormen [77](#page=77). $$ \text{atoom} + \text{e}^- \rightarrow \text{negatief ion} + \text{E.A.} $$ [77](#page=77). Een hogere elektronenaffiniteit betekent dat het atoom het elektron gemakkelijker opneemt. De elektronenaffiniteit neemt toe van links naar rechts en van onder naar boven in het periodiek systeem. Dit geldt als de kernlading groter is, het hoofdkwantumgetal kleiner is en de afscherming van elektronen kleiner is [77](#page=77) [78](#page=78). > **Tip:** De termen "neemt toe" voor E.A. kunnen misleidend zijn, omdat een grotere (meer negatieve) energieverandering wijst op een sterkere affiniteit. Soms wordt de term "positiever" geassocieerd met een hogere affiniteit (minder negatieve waarde). Let op de definitie in de context. ### 4.5 Elektronegativiteit (E.N.) Elektronegativiteit is een maat voor de neiging van een gebonden atoom om de gebonden elektronen naar zich toe te trekken [80](#page=80). * **Trend in periodes:** Binnen een periode neemt de elektronegativiteit toe als het atoomnummer ($Z$) stijgt [80](#page=80). * **Trend in groepen:** Binnen een groep neemt de elektronegativiteit af als het atoomnummer ($Z$) stijgt [80](#page=80). Elementen met een hoge elektronegativiteit trekken gemakkelijk elektronen aan, terwijl elementen met een lage elektronegativiteit dit minder goed doen [80](#page=80). ### 4.6 Oxidatiegetal Het oxidatiegetal (ook wel oxidatietoestand genoemd) is een getal dat de lading van een atoom aangeeft wanneer de verschuiving van de bindings-elektronen volledig zou zijn. Het wordt meestal weergegeven als een Romeins cijfer, voorafgegaan door een plus- of minteken. De hoogste positieve oxidatiegetallen komen overeen met het groepsnummer voor de hoofdgroepen (I tot VII). Negatieve oxidatiegetallen ontstaan wanneer de groepsnummer-8 wordt toegepast (bv. groep VIa heeft laagste -II) [81](#page=81) [82](#page=82). #### 4.6.1 Oxidatiegetallen van hoofdgroepen | Groepsnummer | Ia | IIa | IIIa | IVa | Va | VIa | VIIa | | :----------- | :--- | :--- | :--- | :--- | :--- | :--- | :--- | | Hoogste +OG | +I | +II | +III | +IV | +V | +VI | +VII | | Laagste -OG | - | - | - | -IV | -III | -II | -I | | Lading + ion | 1+ | 2+ | 3+ | 4+ | | | | | Lading - ion | | | | | 3- | 2- | 1- | > **Tip:** Het begrijpen van oxidatiegetallen is cruciaal voor het werken met redoxreacties en de naamgeving van chemische verbindingen. De periodieke trends helpen bij het voorspellen van de meest voorkomende oxidatietoestanden. --- ## Veelgemaakte fouten om te vermijden - Bestudeer alle onderwerpen grondig voor examens - Let op formules en belangrijke definities - Oefen met de voorbeelden in elke sectie - Memoriseer niet zonder de onderliggende concepten te begrijpen

| Term | Definition | |------|------------| | Atoom | De kleinste eenheid van een chemisch element die zijn unieke chemische eigenschappen behoudt, bestaande uit een positieve kern en omringende negatieve elektronen. | | Kern | Het centrale, dichte deel van een atoom dat bestaat uit protonen en neutronen, en waar de positieve lading geconcentreerd is. | | Proton | Een elementair deeltje met een positieve lading ($+1,6 \times 10^{-19}$ C) en een massa van ongeveer $1,673 \times 10^{-27}$ kg, dat zich in de atoomkern bevindt. | | Neutron | Een elementair deeltje zonder elektrische lading en met een massa vergelijkbaar met die van een proton (ongeveer $1,675 \times 10^{-27}$ kg), dat zich in de atoomkern bevindt. | | Elektron | Een elementair deeltje met een negatieve lading ($-1,6 \times 10^{-19}$ C) en een zeer kleine massa (ongeveer $9,11 \times 10^{-31}$ kg), dat zich buiten de atoomkern in schillen bevindt. | | Atoomnummer (Z) | Het aantal protonen in de kern van een atoom, dat de identiteit van een chemisch element bepaalt en het rangnummer in het periodiek systeem aangeeft. | | Massagetal (A) | Het totale aantal protonen en neutronen in de kern van een atoom, wat een benadering geeft van de absolute atoommassa. | | Isotoop | Een atoomsoort van een bepaald element dat hetzelfde aantal protonen (atoomnummer Z) heeft, maar een verschillend aantal neutronen (massagetal A), wat resulteert in verschillende massa's. | | Absolute atoommassa | De werkelijke massa van een individueel atoom, meestal uitgedrukt in kilogram (kg) of atomaire massa-eenheden (u). | | Relatieve atoommassa | De verhouding van de gemiddelde massa van een atoom van een element tot een standaard massa-eenheid, rekening houdend met de natuurlijke abundantie van zijn isotopen. | | Massadefect | Het verschil tussen de berekende massa van een atoomkern (som van de massa's van individuele nucleonen) en de werkelijke gemeten massa van die kern, wat duidt op de omzetting van massa naar bindingsenergie volgens $E=mc^2$. | | Atoommodel | Een theoretische voorstelling van de structuur van een atoom, ontwikkeld door verschillende wetenschappers door de geschiedenis heen, zoals het model van Thomson, Rutherford, Bohr en het kwantummechanisch model. | | Kwantumgetal | Een set getallen die de eigenschappen van elektronen in atomen beschrijven, waaronder het hoofdkwantumgetal (n), nevenkwantumgetal (l), magnetisch kwantumgetal (mL) en spinkwantumgetal (ms). | | Hoofdkwantumgetal (n) | Beschrijft het energieniveau of de schil van een elektron in een atoom, waarbij hogere waarden van n een grotere afstand tot de kern en een hogere energie aangeven. | | Nevenkwantumgetal (l) | Beschrijft de vorm van het orbitaal en de subniveaus (s, p, d, f) binnen een hoofdschil, met waarden van 0 tot n-1. | | Magnetisch kwantumgetal (mL) | Beschrijft de oriëntatie van een orbitaal in de ruimte binnen een subniveau, met waarden van -l tot +l. | | Spinkwantumgetal (ms) | Beschrijft de intrinsieke hoekmomentum van een elektron, dat twee mogelijke waarden heeft, meestal aangeduid als spin-up ($+\frac{1}{2}$) en spin-down ($-\frac{1}{2}$). | | Atoomorbitaal | Een wiskundige functie die de waarschijnlijkheid beschrijft om een elektron te vinden in een bepaald gebied rond de atoomkern, met specifieke vormen (s, p, d, f) en oriëntaties. | | Elektronenconfiguratie | De specifieke rangschikking van elektronen in de orbitalen en energieniveaus van een atoom of ion, bepaald door regels zoals de regel van de minimale energie, de regel van Hund en het Pauli-verbod. | | Periodiek systeem (PSE) | Een tabel die alle bekende chemische elementen ordent op basis van hun atoomnummer, elektronenconfiguratie en terugkerende chemische eigenschappen. | | Groep | Een verticale kolom in het periodiek systeem die elementen bevat met vergelijkbare chemische eigenschappen, meestal vanwege een gelijk aantal valentie-elektronen. | | Periode | Een horizontale rij in het periodiek systeem die elementen bevat met hetzelfde aantal hoofdschillen (energiekernen). | | Valentie-elektronen | De elektronen in de buitenste schil van een atoom, die voornamelijk verantwoordelijk zijn voor chemische reacties en bindingen. | | Effectieve kernlading ($Z_{eff}$) | De netto positieve lading die een elektron in een atoom ervaart, rekening houdend met de afscherming door andere elektronen. | | Atoomstraal | De halve afstand tussen de kernen van twee identieke, naburige atomen in een molecuul of kristalrooster, die de grootte van een atoom aangeeft. | | Ionisatie-energie (I.E.) | De minimale energie die nodig is om een elektron volledig te verwijderen uit een atoom of ion in de gasfase. | | Elektronenaffiniteit (E.A.) | De energieverandering die optreedt wanneer een elektron wordt toegevoegd aan een neutraal gasvormig atoom om een negatief ion te vormen. | | Elektronegativiteit (E.N.) | Een maat voor de neiging van een atoom om gebonden elektronen naar zich toe te trekken in een chemische binding. | | Oxidatiegetal | Een hypothetische lading die een atoom zou hebben als alle bindingselektronen volledig zouden zijn overgedragen naar het meest elektronegatieve atoom in de verbinding. |

# Periodic trends in atomic properties This section summarizes the periodic variations observed in key atomic properties, including atomic radius, ionization potential, electron affinity, and electronegativity, across periods and down groups of the periodic table [2](#page=2). ### 1.1 Atomic radius Atomic radius, or atomic size, refers to the size of an atom and can be determined by measuring the distance between atoms in a combined state. Different types of atomic radii exist, including covalent radius (distance between the nuclei of two bonded atoms, measured at the mean position of shared electrons) and ionic radius (distance between the nuclei of neighboring cations and anions) [3](#page=3) [4](#page=4). #### 1.1.1 Periodic variation in atomic radius * **Across a period:** Atomic size generally decreases as one moves from left to right across a period. This occurs because, while electrons are added to the same valence shell, the nuclear charge increases with increasing atomic number. This leads to a stronger attraction between the nucleus and the outermost electrons, reducing the atomic radius [6](#page=6) [7](#page=7). * **Down a group:** Atomic size generally increases as one moves down a group. Here, added electrons occupy new valence shells, and although the nuclear charge increases, the outermost electrons are further from the nucleus and shielded by inner electrons. This weaker attraction results in a larger atomic radius [8](#page=8) [9](#page=9). #### 1.1.2 Atomic radius of ions * **Cations:** Cations are smaller than their parent atoms. When an atom loses an electron to form a cation, the nuclear charge remains the same while the number of electrons decreases. This leads to an increased effective nuclear charge, pulling the remaining electrons closer to the nucleus and reducing the atomic size. For example, Na+ is smaller than Na [11](#page=11) [12](#page=12) [13](#page=13) [17](#page=17). * **Anions:** Anions are larger than their parent atoms. When an atom gains an electron to form an anion, the nuclear charge remains the same, but the number of electrons increases. This results in a decreased effective nuclear charge, weakening the attraction between the nucleus and the outermost electrons, and thus increasing the atomic size. For example, Br– is larger than Br [14](#page=14) [15](#page=15) [17](#page=17). #### 1.1.3 Isoelectronic species Isoelectronic species are atoms and ions that share the same electronic configuration, meaning they have the same number of electrons. Examples include N3-, O2-, and Na+. When comparing the size of isoelectronic species, the one with the greater nuclear charge will be smaller because its nucleus exerts a stronger pull on the electron cloud. For instance, among N3-, O2-, and Na+, Na+ is the smallest due to its higher nuclear charge [18](#page=18) [19](#page=19). ### 1.2 Ionization potential Ionization potential, also known as ionization energy or ionization enthalpy, is defined as the minimum energy required to remove the most loosely bound electron from an isolated atom in its gaseous state. This process can be represented as [22](#page=22): $$A_{(g)} + \text{Energy} \rightarrow A^{+}_{(g)} + e^-$$ [22](#page=22). The first ionization energy ($I_1$) is the energy to remove the first electron, the second ionization energy ($I_2$) is for the second electron, and so on. Successive ionization energies increase ($I_1 < I_2 < I_3$) because it becomes progressively harder to remove an electron from a positively charged ion due to increased effective nuclear charge [22](#page=22) [23](#page=23). #### 1.2.1 Periodic variation in ionization potential * **Across a period:** Ionization potential generally increases from left to right across a period. This trend mirrors the decrease in atomic radius. As the nuclear charge increases while electrons are added to the same valence shell, the attraction between the nucleus and the outermost electrons strengthens, requiring more energy to remove an electron [26](#page=26). * **Down a group:** Ionization potential generally decreases from top to bottom down a group. This correlates with the increase in atomic radius. As electrons are added to new shells and shielding increases, the attraction of the nucleus for the outermost electron weakens, making it easier to remove and thus lowering the ionization potential [27](#page=27). #### 1.2.2 Factors affecting ionization potential 1. **Atomic size:** As atomic size increases, the attraction between the nucleus and the outermost electron decreases, leading to a lower ionization potential. The relationship is inverse: atomic size $\propto \frac{1}{\text{ionization potential}}$ [28](#page=28). 2. **Nuclear charge:** A greater nuclear charge leads to a stronger attraction for electrons, resulting in a higher ionization potential. The relationship is direct: nuclear charge $\propto$ ionization potential. For instance, Carbon (atomic number 6, nuclear charge +6) has a higher ionization potential than Boron (atomic number 5, nuclear charge +5) [29](#page=29). 3. **Shielding effect:** Increased shielding by inner electrons reduces the effective nuclear charge experienced by the outermost electrons, decreasing the force of attraction and thus lowering the ionization potential. The relationship is inverse: shielding effect $\propto \frac{1}{\text{I.P}}$ [31](#page=31). 4. **Electronic configuration:** Atoms with stable electronic configurations, such as half-filled or fully-filled valence shells, have higher ionization potentials. This is because removing an electron from such a stable configuration would disrupt it, requiring more energy. For example, Nitrogen, with its half-filled 2p orbital ($2p^3$), has a higher ionization potential than Oxygen ($2p^4$), despite Oxygen having a greater nuclear charge. Similarly, elements like Neon (fully filled $2p^6$) have very high ionization potentials [32](#page=32) [33](#page=33) [34](#page=34). ### 1.3 Electron affinity Electron affinity is defined as the energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. It is also referred to as electron gain enthalpy. The process can be represented as [38](#page=38): $$A_{(g)} + e^- \rightarrow A^{-}_{(g)} + \text{Electron Affinity}$$ [38](#page=38). #### 1.3.1 Factors affecting electron affinity 1. **Atomic size:** As atomic size decreases, the attraction between the nucleus and the incoming electron increases, leading to a greater release of energy and thus higher electron affinity. The relationship is inverse: atomic size $\propto \frac{1}{\text{electron affinity}}$ [40](#page=40). * Across a period: Electron affinity generally increases from left to right due to decreasing atomic size and increasing nuclear charge [40](#page=40). * Down a group: Electron affinity generally decreases from top to bottom due to increasing atomic size [40](#page=40). 2. **Nuclear charge:** A greater nuclear charge results in a stronger attraction for an incoming electron, leading to higher electron affinity. The relationship is direct: nuclear charge $\propto$ electron affinity. For example, Fluorine (nuclear charge +9) has a greater electron affinity than Oxygen (nuclear charge +8) [41](#page=41). 3. **Electronic configuration:** Atoms with stable electronic configurations (half-filled or fully-filled) exhibit lower or even zero electron affinity. Adding an electron to such a stable configuration would disrupt it, reducing the likelihood of electron gain and the energy released. For example, Nitrogen, with its stable half-filled 2p orbital, has a lower electron affinity than Carbon [42](#page=42) [43](#page=43). #### 1.3.2 Important trends in electron affinity * Halogens have the highest electron affinities because they are one electron short of a stable noble gas configuration [44](#page=44). * Chlorine has a higher electron affinity than fluorine, despite fluorine being more electronegative. This anomaly is attributed to strong inter-electronic repulsions within the small fluorine atom, which hinder the addition of an incoming electron [44](#page=44). * Noble gases have zero electron affinity as they already possess stable, complete electronic configurations, making them unwilling to accept additional electrons [44](#page=44). ### 1.4 Electronegativity Electronegativity describes the tendency of an atom to attract electrons towards itself within a chemical bond. Fluorine is the most electronegative element, while Cesium is the least electronegative (most electropositive) [46](#page=46). #### 1.4.1 Periodic trend in electronegativity Electronegativity generally increases across a period and decreases down a group, similar to the trends in ionization potential [47](#page=47). #### 1.4.2 Electronegativity and bond nature The difference in electronegativity between two bonded atoms provides insight into the type of bond formed [48](#page=48): * A large electronegativity difference leads to the formation of an ionic bond (e.g., NaCl, LiF) [48](#page=48). * A small electronegativity difference results in the formation of a covalent bond (e.g., CO, NO) [48](#page=48). --- # Metallurgy and extraction of metals Metallurgy is the process of extracting metals in a pure form from their ores. The extraction method depends on the physical and chemical properties of the metal [51](#page=51). ### 2.1 Mode of occurrence of metals in nature Metals occur in nature in two main states: * **Native State:** Metals are found in a pure, uncombined form. They are typically unreactive and are not easily attacked by moisture, oxygen, or carbon dioxide. Examples include gold (Au) and platinum (Pt) [52](#page=52). * **Combined State:** Metals are found in combined forms, usually as compounds with other elements. These can include oxides, sulfides, carbonates, sulfates, halides, etc. [52](#page=52). * **Oxides:** Bauxite (Al₂O₃), Haematite (Fe₂O₃) [52](#page=52). * **Sulfides:** Copper pyrite (CuFeS₂), Zinc blende (ZnS) [52](#page=52). * **Carbonates:** Calamine (ZnCO₃) [52](#page=52). ### 2.2 Minerals and ores * **Minerals:** All naturally occurring chemical substances in which metals are found in nature, along with impurities [53](#page=53). * **Ores:** Minerals from which the extraction of a metal is economically and conveniently feasible [53](#page=53). * **Example:** Bauxite (Al₂O₃) is an ore of aluminum. Clay (Al₂O₃·2SiO₂·2H₂O) is a mineral, but not typically an ore [53](#page=53). * **Relationship:** Every ore is a mineral, but not every mineral is an ore [53](#page=53). ### 2.3 Types of metallurgical processes Based on the temperature at which extraction takes place, metallurgical processes can be classified into: * **Pyrometallurgy:** Involves extraction at very high temperatures. Metals like copper (Cu), iron (Fe), zinc (Zn), and tin (Sn) are extracted using this method [54](#page=54). * **Hydrometallurgy:** Uses aqueous solutions for metal extraction. Silver (Ag) and gold (Au) are extracted this way [54](#page=54). * **Electrometallurgy:** Involves using electrolytic methods for extraction from molten salt solutions. Highly reactive metals such as sodium (Na), potassium (K), lithium (Li), and calcium (Ca) are extracted using this method [54](#page=54). ### 2.4 General steps of metallurgy The general sequence of steps involved in the metallurgical extraction of metals includes: 1. **Crushing and Pulverization:** Reducing the ore to a fine powder [55](#page=55) [57](#page=57). 2. **Concentration of the ore (Ore Dressing):** Removing unwanted impurities (gangue) from the ore [55](#page=55) [59](#page=59). 3. **Conversion of the concentrated ore into metal oxide:** This prepares the ore for reduction [55](#page=55) [69](#page=69). 4. **Reduction of metal oxide into metal:** Obtaining the metal from its oxide [55](#page=55) [76](#page=76). 5. **Refining:** Purifying the crude metal obtained from reduction [55](#page=55) [80](#page=80). #### 2.4.1 Crushing and pulverization Ores are typically found as large lumps. These are first broken down into smaller pieces using jaw crushers and grinders (crushing). Subsequently, the small pieces are reduced to a fine powder using equipment like ball mills or stamp mills (pulverization) [57](#page=57). #### 2.4.2 Concentration of the ore (Ore dressing) The undesired rocky or earthy impurities associated with an ore are known as the **gangue** or **matrix**. The process of removing the gangue from the ore is called **concentration**. Various methods are employed for concentration, depending on the nature of the ore and the impurities [59](#page=59): ##### 2.4.2.1 Methods of concentration * **A. Hand Picking:** Suitable for ores with visibly distinct impurities that can be identified and removed by hand [61](#page=61). * **B. Hydraulic Washing (Levigation / Gravity Separation):** This method is employed when the ore particles are significantly heavier than the gangue particles. The powdered ore is agitated with water or washed with a stream of water. The heavier ore particles settle down, while the lighter impurities are washed away [61](#page=61). * **C. Froth Floatation:** Primarily used for the concentration of sulfide ores, such as copper pyrite (CuFeS₂), galena (PbS), and zinc blende (ZnS). The powdered ore is mixed with water and pine oil. Air is blown through the mixture, causing the ore particles (which get wet by the oil) to become lighter, rise to the surface as froth, and are then collected. The impurities (gangue), which get wet by water, settle down at the bottom [63](#page=63). * **D. Electromagnetic Separation:** This method is used when either the ore or the impurities possess magnetic properties. For example, magnetite (Fe₃O₄) and chromite (FeO·Cr₂O₃) ores can be concentrated using this method. The powdered ore is fed onto a moving belt that passes over rollers, at least one of which is electromagnetic. Magnetic particles are attracted to the electromagnet and fall closer to the roller, while non-magnetic particles fall away, allowing for their separation [65](#page=65). * **E. Leaching:** This process involves treating the powdered ore with a suitable reagent that selectively reacts with the ore, forming a soluble compound, while leaving the impurities unaffected [67](#page=67). * **Example 1: Leaching of Bauxite (Baeyer's process):** The ore is treated with sodium hydroxide (NaOH) to form soluble sodium meta-aluminate. The insoluble impurities are filtered off, and then aluminum hydroxide is precipitated from the filtrate [67](#page=67). * **Example 2: Leaching of Silver and Gold (Cyanide Process):** Used for extracting silver and gold [67](#page=67). * **F. Electrostatic separation:** (Mentioned in page 60 but not detailed in the provided text beyond its name.) #### 2.4.3 Conversion of the concentrated ore into metal oxide After concentration, the ore is converted into its corresponding metal oxide. This is a crucial step before reduction. The methods depend on the nature of the ore [70](#page=70). ##### 2.4.3.1 Calcination * **Definition:** The process of heating a concentrated ore strongly below its melting point in the absence or limited supply of air [71](#page=71). * **Purpose:** Primarily used to convert metal carbonates and hydroxides into their respective oxides. It also removes moisture, volatile impurities (like S, As, P as their oxides), and organic matter [71](#page=71). * **Examples:** * $CaCO_3 \xrightarrow{Heat} CaO + CO_2\uparrow$ (Limestone) [72](#page=72). * $CaCO_3 \cdot MgCO_3 \xrightarrow{Heat} CaO + MgO + 2CO_2\uparrow$ (Dolomite) [72](#page=72). * $ZnCO_3 \xrightarrow{Heat} ZnO + CO_2\uparrow$ [72](#page=72). ##### 2.4.3.2 Roasting * **Definition:** The process of heating a concentrated ore strongly below its melting point in an excess of air [73](#page=73). * **Purpose:** Commonly used to convert sulfide ores into their respective metal oxides. It also removes moisture, volatile impurities (S, As, P as their oxides), and organic matter [73](#page=73). * **Examples:** * Removal of volatile impurities: * $S_8 + 8O_2 \rightarrow 8SO_2\uparrow$ [74](#page=74). * $P_4 + 5O_2 \rightarrow P_4O_{10}$ [74](#page=74). * $4As + 3O_2 \rightarrow 2As_2O_3\uparrow$ [74](#page=74). * Conversion of sulfide ores to metallic oxides: * $2ZnS + 3O_2 \xrightarrow{Heat} 2ZnO + 2SO_2\uparrow$ [74](#page=74). * $2PbS + 3O_2 \xrightarrow{Heat} 2PbO + 2SO_2\uparrow$ [74](#page=74). * $2Cu_2S + 3O_2 \xrightarrow{Heat} 2Cu_2O + 2SO_2\uparrow$ [74](#page=74). * **Furnace:** Roasting and calcination are typically carried out in a reverberatory furnace [75](#page=75). #### 2.4.4 Reduction of metal oxide into metal The metal oxide obtained from calcination or roasting is then reduced to obtain the crude metal. The choice of reducing agent depends on the reactivity of the metal [76](#page=76) [77](#page=77). * **Highly reactive metals** (e.g., Na, K, Ca, Mg, Al): Reduced by electrolytic methods from their molten salts [77](#page=77). * **Less reactive metals** (e.g., Zn, Fe, Pb, Sn, Cr): Can be reduced by using reducing agents like carbon (coke) or carbon monoxide (CO) [77](#page=77). ##### 2.4.4.1 Smelting * **Definition:** The process of extracting a metal from its metal oxide by reduction using carbon (coke/charcoal) or carbon monoxide [78](#page=78). * **Furnace:** Smelting is commonly carried out in a blast furnace [78](#page=78). ##### 2.4.4.2 Aluminothermite Process (Aluminothermy or Goldschmidt Thermite Process) * **Definition:** This process uses aluminum powder as a reducing agent to extract metals from their oxides [79](#page=79). * **Application:** It is particularly useful for reducing oxides of metals that cannot be easily reduced by carbon or carbon monoxide, such as Fe₂O₃ and Cr₂O₃, because aluminum is highly electropositive [79](#page=79). * **Examples:** * $2Al(s) + Fe_2O_3(s) \rightarrow Al_2O_3(s) + 2Fe(l)$ [79](#page=79). * $2Al + Cr_2O_3 \rightarrow Al_2O_3 + 2Cr$ [79](#page=79). #### 2.4.5 Refining or purification of metals The metal obtained after reduction is usually impure and is called **crude metal**. Impurities can include other metals formed during reduction, non-metals (Si, P), unreacted oxides and sulfides, and residual slag and flux. Refining is the process of purifying the crude metal to obtain pure metal [81](#page=81) [82](#page=82). ##### 2.4.5.1 Common methods for refining * **Poling** [82](#page=82). * **Electro-refining** [82](#page=82). * **Zone refining / Fractional Crystallization** [82](#page=82). * **Vapour-phase refining** [82](#page=82). * **Chromatographic adsorption method** [82](#page=82). * **Distillation** [82](#page=82). ##### 2.4.5.2 Electrolytic refining * **Application:** Used for refining metals like copper (Cu), silver (Ag), gold (Au), lead (Pb), nickel (Ni), chromium (Cr), zinc (Zn), and aluminum (Al) [83](#page=83). * **Process:** In this method, the impure metal is made the anode, and a strip of pure metal of the same kind is made the cathode. Both electrodes are immersed in an electrolyte solution containing metal ions. When an electric current is passed, metal ions from the electrolyte deposit onto the cathode as pure metal. Soluble impurities pass into the electrolyte, while insoluble impurities (like precious metals or unreacted metals) settle down at the bottom as anode mud [83](#page=83). * **Electrode Reactions:** * At Anode: $M(s) \rightarrow M^{n+}(aq) + ne^-$ [84](#page=84). * At Cathode: $M^{n+}(aq) + ne^- \rightarrow M(s)$ [84](#page=84). * **Example: Refining of Copper:** * At anode: $Cu(s) \rightarrow Cu^{2+}(aq) + 2e^-$ [84](#page=84). * At cathode: $Cu^{2+}(aq) + 2e^- \rightarrow Cu(s)$ [84](#page=84). ##### 2.4.5.3 Poling * **Application:** This method is suitable for refining metals that contain impurities in the form of their own oxides, such as impure copper containing copper(I) oxide ($Cu_2O$) [85](#page=85). * **Process:** The molten impure metal is stirred with green poles (often made of wood). The hydrocarbon gases released from the burning wood act as reducing agents, reducing the oxide impurities to the pure metal [85](#page=85). * **Example:** $Cu_2O + CH_4 \rightarrow Cu + H_2O + CO$ [85](#page=85). --- # Alkali and Alkaline Earth Metals This topic explores the properties, extraction, and chemical behavior of alkali metals (Group 1) and alkaline earth metals (Group 2) of the periodic table [96](#page=96). ### 3.1 Alkali metals Alkali metals belong to Group 1 (IA) of the periodic table and include lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). They are named so because they react with water to form alkalis, which are strong bases capable of neutralizing acids. A general reaction is given by [97](#page=97): `2M + 2H2O → 2MOH + H2` [97](#page=97). where M represents an alkali metal. #### 3.1.1 Extraction of sodium by Down's process Sodium is primarily extracted from sodium chloride (NaCl) due to its abundance and low cost. Chemical reduction is not feasible as alkali metals are strong reducing agents themselves and cannot be reduced by common agents like coke, Al, or others. Pyrometallurgy is also unsuitable because sodium vaporizes at high temperatures and its vapors are highly reactive, making collection difficult. Therefore, electrolysis of molten NaCl is employed (electrometallurgy) [87](#page=87) [88](#page=88). **Challenges in Electrolysis and Down's Process Solution:** * **High Melting Point of NaCl:** NaCl has a high melting point (802°C) making it expensive to maintain the necessary high temperatures for electrolysis [89](#page=89). * **Sodium Vaporization:** The boiling point of sodium (883°C) is close to the melting point of NaCl, leading to sodium vaporization and the formation of metallic fog, which can short-circuit the cell [89](#page=89). * **Reactivity of Products:** The electrolytic products (sodium and chlorine) are highly reactive at high temperatures, corroding the vessel materials [90](#page=90). Down's process overcomes these issues by mixing NaCl with calcium chloride (CaCl2) in a 2:3 ratio. This lowers the melting point to approximately 600°C. At this reduced temperature, sodium and chlorine are less reactive, minimizing corrosion and the formation of metallic fog [90](#page=90). **Electrolysis in Down's Process:** During electrolysis, NaCl dissociates into ions: `NaCl → Na+ + Cl-` [92](#page=92). At the anode (oxidation): `Cl- → Cl + e-` [92](#page=92). `Cl + Cl → Cl2` [92](#page=92). At the cathode (reduction): `Na+ + e- → Na` [92](#page=92). Molten sodium is collected at the cathode, and chlorine gas is collected at the anode [92](#page=92). #### 3.1.2 Physical properties of alkali metals * **Physical state:** Alkali metals are soft, malleable, and ductile solids. They are silvery white when freshly cut but tarnish quickly in air. Softness increases down the group due to weaker metallic bonding as atomic size increases [98](#page=98). * **Atomic size:** Atomic size increases down the group with the addition of electrons in new shells [98](#page=98). * **Ionization energy:** They possess low first ionization energies, indicating a tendency to lose one electron [98](#page=98). * **Melting and boiling points:** These are low due to weak metallic bonding, and they decrease down the group as the metallic bond weakens [99](#page=99). * **Density:** Alkali metals have remarkably low densities, which increase down the group [99](#page=99). * **Conductivity:** They are excellent conductors of heat and electricity [99](#page=99). * **Nature of bond:** They form ionic bonds with non-metals by losing one electron to achieve an octet state [99](#page=99). * **Electronegativity:** Their electronegativity is low due to large atomic size, making them highly electropositive with a greater tendency to lose electrons [100](#page=100). * **Flame colour:** Alkali metals impart characteristic colours to a Bunsen flame: * Li: Crimson [100](#page=100). * Na: Golden yellow [100](#page=100). * K: Violet [100](#page=100). * Rb: Red violet [100](#page=100). #### 3.1.3 Chemical characteristics of alkali metals 1. **Action with air:** Alkali metals react with atmospheric oxygen, causing them to tarnish. They burn in oxygen to form oxides, but the nature of these oxides differs: * Lithium forms a normal oxide: `4 Li + O2 → 2Li2O` (Oxidation state of O = -2) . * Sodium forms a peroxide: `2 Na + O2 → Na2O2` (Oxidation state of O = -1) . * Larger alkali metals like potassium form superoxides: `K + O2 → KO2` (Oxidation state of O = -1/2) . 2. **Action with water:** The reaction is exothermic, and reactivity increases down the group. Lithium reacts gently, sodium reacts vigorously (may catch fire), and potassium reacts violently. The general reaction is : `2M + 2H2O → 2MOH + H2` [93](#page=93). 3. **Action with acids:** Alkali metals react with acids to produce hydrogen gas and metal salts [93](#page=93). `2Na + 2HCl → 2NaCl + H2` [93](#page=93). `2Na + H2SO4 → Na2SO4 + H2` [93](#page=93). 4. **Action with hydrogen:** All alkali metals react with hydrogen at elevated temperatures to form metal hydrides (M=Li, Na, K, Rb, Cs) . `2Na + H2 → 2NaH` (sodium hydride) . 5. **Action with halogens:** Alkali metals react with halogens to form ionic halides . `Na + Cl2 → 2NaCl` [94](#page=94). 6. **Action with non-metals:** They react with non-metals like sulfur and phosphorus to form sulfides and phosphides . `2Na + S → 2Na2S` [94](#page=94). `3Na + P → 2Na3P` . 7. **Reaction with ammonia gas:** Heating alkali metals with ammonia gas produces sodium amide and hydrogen gas [94](#page=94). `2Na + 2NH3 → 2NaNH2 + H2` (sodamide) . `2K + 2NH3 → 2KNH2 + H2` (potassamide) . 8. **Reaction with liquid ammonia:** Alkali metals dissolve in liquid ammonia to form ammoniated metal ions and ammoniated electrons . #### 3.1.4 Sodium carbonate (Na2CO3) **Properties of Sodium Carbonate:** * **Physical properties:** It is a white crystalline solid that forms a decahydrate (Na2CO3·10H2O), known as washing soda. The anhydrous form is called soda ash. It is highly soluble in water . * **Chemical properties:** 1. **Action of water:** Soluble in water due to hydrolysis, forming sodium hydroxide and carbonic acid . 2. **Action of CO2 and SO2:** Passing CO2 or SO2 through aqueous sodium carbonate precipitates sodium bicarbonate (baking soda) or sodium bisulfite, respectively . 3. **Precipitation reaction:** It precipitates metal carbonates from their soluble salt solutions. This property is utilized in water softening to remove permanent hardness caused by soluble calcium and magnesium chlorides and sulfates . `Reaction with lime water (Ca(OH)2):` . **Uses of Sodium Carbonate:** Manufacture of glass, soap, wood pulp, paper; water softening agent; production of baking soda; used in paints and dyes . **Manufacture of Sodium Carbonate by Solvay Process (Ammonia-Soda Process):** This process, developed by Ernest Solvay, involves the reaction of brine saturated with ammonia with carbon dioxide to form sodium bicarbonate, which is sparingly soluble and precipitates out. Heating the filtered sodium bicarbonate yields sodium carbonate . **Process Steps:** 1. **Saturation of brine:** Brine solution is saturated with ammonia in an ammonia absorber. Impurities like CaCl2 and MgCl2 are precipitated as hydroxides and carbonates and removed . 2. **Carbonation:** Ammoniated brine is passed down a tower while CO2 is introduced from the bottom. Sodium bicarbonate (NaHCO3) and ammonium chloride (NH4Cl) are formed. The reaction is: `NaCl + NH3 + CO2 + H2O → NaHCO3 + NH4Cl` . 3. **Recovery of ammonia:** Ammonium chloride reacts with slaked lime in the ammonia generator to regenerate ammonia gas . `2NH4Cl + Ca(OH)2 → 2NH3 + CaCl2 + 2H2O` 4. **Generation of CO2:** Limestone is heated in a lime kiln to produce CO2. The resulting CaO is treated with water to produce slaked lime for ammonia regeneration . `CaCO3 → CaO + CO2` 5. **Calcination:** The filtered sodium bicarbonate is heated to obtain anhydrous sodium carbonate . `2NaHCO3 → Na2CO3 + H2O + CO2` 6. **Crystallization:** Aqueous sodium carbonate is crystallized to obtain washing soda . #### 3.1.5 Sodium hydroxide (NaOH) NaOH is also known as Caustic Soda . **Properties of NaOH:** * **Physical properties:** It is a white, crystalline, deliquescent solid that absorbs moisture from the air. It is highly soluble in water . * **Chemical properties:** 1. **Precipitation reactions:** NaOH precipitates metallic hydroxides from their salt solutions. Some amphoteric metal hydroxides are soluble in excess NaOH . 2. **Action with CO:** Reacts with carbon monoxide at high temperature and pressure to form the sodium salt of formic acid, which yields formic acid upon acidification . **Uses of NaOH:** Manufacture of pulp, paper, soaps, and detergents; cleansing agent for grease, fats, and proteins; refining of vegetable oils; extraction of aluminum; reagent in laboratory reactions . **Manufacture of Sodium Hydroxide using Diaphragm Cell:** NaOH is produced by the electrolysis of aqueous NaCl solution. The cell uses a titanium oxide anode, a steel mesh cathode enclosed in a Teflon diaphragm, and brine as the electrolyte . **Electrolysis Reactions:** The aqueous solution ionizes as: `NaCl → Na+ + Cl-` . `H2O → H+ + OH-` . At the anode (oxidation): Cl- ions have a lower discharge potential than OH- ions, so chlorine gas is liberated. `2Cl- → Cl2 + 2e-` . At the cathode (reduction): H+ ions have a lower discharge potential than Na+ ions, so hydrogen gas is formed. The ion-exchange membrane allows selective ion flow to the cathode . `2H+ + 2e- → H2` . **Recovery of NaOH:** The solution in the cathode compartment, containing excess Na+ and OH- ions, is periodically removed and concentrated. Sodium chloride crystallizes out, leaving a concentrated NaOH solution, which can be evaporated to dryness to obtain solid NaOH . ### 3.2 Alkaline earth metals Alkaline earth metals are elements of Group 2 of the periodic table . #### 3.2.1 Physical properties of alkaline earth metals * **Physical state:** They are soft but harder than alkali metals due to stronger metallic bonding . * **Atomic size:** They are larger than alkali metals but smaller than Group 1 elements due to a greater nuclear charge. Atomic size increases down the group . * **Ionization energy:** They have low first ionization energies, but these are higher than alkali metals because they need to lose two electrons . * **Melting and boiling points:** These are low but higher than those of alkali metals . * **Density:** They have low densities, higher than alkali metals due to smaller atomic size . * **Electronegativity:** Their electronegativity is low due to large atomic size, making them highly electropositive but less so than alkali metals . * **Conductivity:** They are excellent conductors of heat and electricity . * **Flame colour:** Except for Be and Mg, they impart characteristic colours to a Bunsen flame: * Ca: Brick red . * Sr: Crimson red . * Ba: Apple green . * **Nature of bond:** They form ionic bonds with non-metals by losing two electrons to achieve an octet state (valency = 2) . #### 3.2.2 Chemical characteristics of alkaline earth metals 1. **Action with air:** Alkaline earth metals react with atmospheric oxygen, causing them to tarnish. They burn in oxygen to form oxides . `2Mg + O2 → 2MgO` . `2 Ca + O2 → 2CaO` . 2. **Action with H2O:** Beryllium (Be) does not react with water. Magnesium (Mg) reacts with hot water. Calcium (Ca), Strontium (Sr), and Barium (Ba) react with cold water to form hydroxides and liberate hydrogen gas . `Mg + 2H2O → Mg(OH)2 + H2` (hot water) . `Ca + 2H2O → Ca(OH)2 + H2` (cold water) . `Ba + 2H2O → Ba(OH)2 + H2` (cold water) . 3. **Action with hydrogen:** They react with hydrogen at elevated temperatures to form metal hydrides, except for Be . `Mg + H2 → MgH2` . `Ca + H2 → CaH2` . 4. **Action with nitrogen:** They react with nitrogen at elevated temperatures to form metal nitrides, except for Be . `Mg + N2 → Mg3N2` . 5. **Action with halogen, S, and P:** They react with halogens, sulfur, and phosphorus to form respective halides, sulfides, and phosphides . `Ca + Cl2 → CaCl2` . `Ca + S → CaS` . `3Ca + 2P → Ca3P2` . #### 3.2.3 Compounds of Ca and Mg 1. **Quick Lime (CaO):** Chemically known as Calcium oxide . * **Uses:** Fertilizer, preparation of cement, glass, bleaching powder; drying agent for gases like ammonia; flux in metallurgy; basic lining in furnaces . 2. **Magnesia (MgO):** Chemically known as Magnesium oxide . * **Uses:** Making crucibles; antacid in medicine; refractory material in furnaces and bricks . 3. **Plaster of Paris (CaSO4·1/2 H2O):** Also known as Gypsum plaster; chemically Calcium sulfate hemihydrate . * **Uses:** Building material for walls and ceilings; plastering fractured bones; making statues and molds . 4. **Bleaching powder (CaOCl2):** Chemically Calcium hypochlorite . * **Uses:** Household bleaching agent, stain remover; manufacturing chloroform; disinfectant for sterilizing water; germicide . 5. **Epsom salt ((MgSO4·7H2O)):** Chemically Magnesium sulfate heptahydrate . * **Uses:** Bath salts; purgative in medicine; mordant for cotton in dyeing; anhydrous form used as a desiccant . #### 3.2.4 Solubility of hydroxides, carbonates, and sulfates of alkaline earth metals * **Solubility of hydroxide:** Alkaline earth metals form basic hydroxides that are fairly soluble in water. Solubility increases down the group. The order is: `Be(OH)2 < Mg(OH)2 < Ca(OH)2 < Sr(OH)2 < Ba(OH)2`. `Ba(OH)2` is the most soluble . > **Tip:** The increase in solubility down the group is due to the lattice enthalpy decreasing faster than the hydration enthalpy as the cation size increases . * **Solubility of carbonate:** Alkaline earth metal carbonates are sparingly soluble in water. Solubility decreases down the group due to a decrease in hydration energy relative to lattice energy. The order is: `BeCO3 > MgCO3 > CaCO3 > SrCO3 > BaCO3`. `BaCO3` is the most insoluble . * **Solubility of sulfate:** Alkaline earth metal sulfates are soluble in water. Solubility decreases down the group due to a decrease in hydration energy relative to lattice energy. The order is: `BeSO4 > MgSO4 > CaSO4 > SrSO4 > BaSO4`. `BaSO4` is the most insoluble. The BaCl2 test for sulfate ions relies on the formation of a white precipitate of BaSO4 . > **Tip:** The change in ion distances for sulfate is less significant compared to hydroxides, causing lattice enthalpy to decrease slower, leading to a decrease in solubility down the group for sulfates . #### 3.2.5 Stability of carbonate and nitrate of alkaline earth metals * **Stability of carbonate:** Group 2 metal carbonates are fairly stable to heat. Thermal stability increases from top to bottom in a group, indicated by increasing decomposition temperatures. Metal carbonates decompose to yield metal oxides and carbon dioxide . `CaCO3 → CaO + CO2` . * **Stability of nitrate:** Group 2 metal nitrates are also stable to heat, with thermal stability increasing down the group. The stability order is: `Be(NO3)2 < Mg(NO3)2 < Ca(NO3)2 < Sr(NO3)2 < Ba(NO3)2`. Metal nitrates decompose to produce metal oxides, nitrogen dioxide, and oxygen . `Ca(NO3)2 → CaO + NO2 + O2` . --- # Bioinorganic Chemistry and Metal Toxicity Bioinorganic chemistry investigates the crucial interactions between inorganic substances and biological systems, distinguishing between essential macro- and micronutrients and examining the detrimental effects of toxic metals . ### 4.1 Bioinorganic chemistry fundamentals Bioinorganic chemistry bridges the fields of inorganic chemistry and biochemistry, focusing on the roles of inorganic compounds within living organisms. These compounds can range from simple metal ions like potassium ($K^{+}$), iron ($Fe^{2+}$), and sodium ($Na^{+}$), to complex ions such as molybdate, coordination compounds like cisplatin, and inorganic molecules including carbon monoxide (CO), nitric oxide (NO), and ozone ($O_3$) . ### 4.2 Essential and trace elements: macro and micronutrients Approximately 40 of the 118 known elements are involved in life processes, with about 30 being essential for human health . #### 4.2.1 Macronutrients Elements required in large quantities are classified as macronutrients and are also known as essential elements. Their absence leads to death or severe organ malfunction. Examples include sodium ($Na$), potassium ($K$), magnesium ($Mg$), calcium ($Ca$), and chlorine ($Cl$) . #### 4.2.2 Micronutrients Elements needed in smaller amounts are termed micronutrients or trace elements. Despite their small quantities, they play vital roles in biological systems, and their deficiency can cause serious defects. Examples include iron ($Fe$), zinc ($Zn$), copper ($Cu$), and cobalt ($Co$) . ### 4.3 Importance of metal ions in biological systems Various metal ions are critical for numerous biological functions: * **Sodium ($Na$):** Regulates osmotic pressure, aids in the absorption of glucose and amino acids, and maintains nerve impulses. Low intake can cause weakness, headache, and potentially kidney and heart failure . * **Potassium ($K$):** Regulates osmotic pressure, facilitates body metabolism, prevents muscle cramps, and offers protection against brain and heart strokes, and osteoporosis. Insufficient intake may lead to fatigue, weakness, and abnormal heart rhythms . * **Calcium ($Ca$):** Essential for the growth and development of healthy bones, facilitates communication between the brain and body parts, and helps maintain blood pressure . * **Magnesium ($Mg$):** Plays a role in the production of enzymes and proteins, supports the proper functioning of DNA and RNA, and maintains electrolytic balance for active ion transport . * **Iron ($Fe$):** Crucial for rapid growth and development during pregnancy; low intake can lead to anemia and internal bleeding . * **Copper ($Cu$):** Maintains nerve cells and the immune system, and assists in iron absorption and energy production . * **Zinc ($Zn$):** Supports nerve cells and the immune system, and promotes wound healing and skin repair . * **Nickel ($Ni$):** Aids in glucose breakdown, supports milk production in mammary glands, and is involved in iron metabolism . * **Cobalt ($Co$):** An integral component of Vitamin B12, it is involved in the metabolism of fatty acids and folic acid, and aids in red blood cell production . * **Chromium ($Cr$):** Enhances insulin production and contributes to weight loss and muscle strengthening . ### 4.4 The sodium-potassium pump The sodium-potassium pump is a primary active transport system that actively moves sodium ($Na^{+}$) ions out of the cytoplasm and potassium ($K^{+}$) ions into the cytoplasm. This process is also referred to as the sodium pump . #### 4.4.1 Mechanism of the sodium-potassium pump The pump binds ATP and three sodium ions ($Na^{+}$) from the cytoplasm. ATP then phosphorylates the pump, causing a conformational change that opens to the extracellular space. The sodium ions are released, and two potassium ions ($K^{+}$) are bound. Finally, the phosphate group is cleaved, returning the pump to its original conformation and releasing potassium ions into the cell . #### 4.4.2 Importance of the sodium-potassium pump The unequal distribution of ions across the cell membrane, with an excess of negative charge inside and positive charge outside, creates an electrical potential gradient essential for transmitting nerve signals in animals. The pump also drives secondary active transport systems for nutrients like amino acids and glucose, and it plays a role in maintaining cellular osmosis . ### 4.5 The sodium-glucose pump The sodium-glucose pump is a secondary active transport system that facilitates glucose absorption into cells using the concentration gradient of sodium ions. The primary active transport of the $Na^{+}-K^{+}$ pump establishes a high extracellular concentration of $Na^{+}$ and a low intracellular concentration. This gradient drives $Na^{+}$ ions to move into the cell, co-transporting glucose molecules. The intracellular glucose is then broken down to produce ATP for cellular functions. This pump is vital for transporting glucose from the extracellular fluid into the cytoplasm of cells, where it can be metabolized for energy . ### 4.6 Metal toxicity Metal toxicity refers to the harmful effects of certain metals in specific forms and doses on living organisms. Exposure to high concentrations of metals can occur through food, air, water pollution, medications, food containers, industrial settings, or even lead-based paints. The toxicity depends on the absorbed dose, route of exposure, and duration (acute or chronic) . Common toxic effects include gastrointestinal and kidney dysfunction, nervous system disorders, skin lesions, vascular damage, immune system dysfunction, birth defects, and cancer. Metals that humans can absorb in toxic amounts include mercury, lead, cadmium, arsenic, and iron . #### 4.6.1 Effects of mercury toxicity Mercury toxicity can lead to a lack of coordination, muscle weakness, hearing and speech difficulties, and damage to nerves in the hands and face. A notable example is Minamata disease in Japan, caused by methyl-mercury poisoning, which resulted in neurological damage affecting speech and sight . #### 4.6.2 Effects of iron toxicity Iron toxicity can cause problems in the lungs, stomach, and intestines, as well as heart and blood-related issues. It can also lead to liver and skin problems, including liver cancer, and nervous system issues. In cases of overdose, it is linked to diseases like Alzheimer's and Parkinson's . #### 4.6.3 Effects of lead toxicity Acute lead exposure can cause constipation, aggressive behavior, high blood pressure, fatigue, and sleep difficulties. Chronic exposure can result in mental retardation, birth defects, autism, and paralysis . #### 4.6.4 Effects of arsenic toxicity Arsenic toxicity may manifest as unusual heart rhythms, muscle cramps, stomach pain, and red or swollen skin. It can also cause wart-like spots on the skin. Chronic toxicity leads to arsenicosis, characterized by skin lesions and pigmentation changes . #### 4.6.5 Effects of cadmium toxicity Cadmium toxicity can decrease hemoglobin levels, leading to anemia, and cause muscle pain and osteoporosis. Chronic exposure can damage the kidneys, liver, and lungs. Itai-Itai disease in Japan, meaning "it hurts-it hurts," is an example of cadmium toxicity, characterized by softening of bones and severe bone and muscle pain . --- # Gas Laws This section explores fundamental gas laws that describe the macroscopic behavior of gases concerning pressure, volume, and temperature . ### 5.1 Properties of gases Gases are characterized by maximum inter-particle space and minimal forces of attraction between particles. Their particles are free to move in any direction, making them highly compressible and lacking fixed shape, size, or volume. Gases also exhibit low density. The behavior of gases can be predicted using variables such as pressure, volume, amount (moles), and temperature . ### 5.2 Gas laws Gas laws provide quantitative relationships between the mass, volume, pressure, and temperature of a gas. They describe the interrelationship between two variables while keeping others constant. Key gas laws include Boyle's Law, Charles' Law, Avogadro's Law, Dalton's Law of Partial Pressure, and Graham's Law of Diffusion . #### 5.2.1 Boyle's law Boyle's Law, formulated by Robert Boyle, establishes a relationship between the pressure (P) and volume (V) of a given mass of gas at a constant temperature . **Statement:** At a constant temperature, the volume of a given mass of a gas is inversely proportional to its pressure . **Mathematical Expression:** The relationship can be expressed as: $P \propto \frac{1}{V}$ or, $P = k \frac{1}{V}$, where $k$ is a proportionality constant . Rearranging this gives: $PV = k$ . **Explanation:** If $V_1$ is the initial volume of a given mass of gas at pressure $P_1$ and temperature $T$, then $P_1V_1 = k$. If the final volume and pressure are $V_2$ and $P_2$ respectively, then $P_2V_2 = k$. Therefore, according to Boyle's Law : $P_1V_1 = P_2V_2 = k$ . This equation signifies that as volume increases, the gas pressure decreases proportionally, and vice-versa. For instance, doubling the pressure at constant temperature will halve the gas volume . **Graphical Representation:** Boyle's Law can be graphically represented by plotting pressure against volume at a constant temperature, yielding a curve known as an isotherm. Further graphical representations include plots of volume versus the reciprocal of pressure ($1/P$), and the product of pressure and volume ($PV$) versus pressure ($P$) . > **Tip:** An isotherm is a graph representing the relationship between pressure and volume of a gas at a constant temperature . **Applications of Boyle's Law:** Boyle's Law helps explain phenomena like why mountaineers need oxygen cylinders at high altitudes. As altitude increases, atmospheric pressure decreases, leading to an increased volume of air and thus lower air density. This reduced air density means less oxygen is available for breathing, necessitating the use of supplemental oxygen cylinders . **Numerical Formulas and Units:** * $PV = k$ . * $P_1V_1 = P_2V_2$ . * $\frac{P_1}{D_1} = \frac{P_2}{D_2}$, where $D$ is density and assuming mass is constant. (Derived from volume = mass/density, so $V = M/D$. If $M$ is constant, $V \propto 1/D$. Substituting into $P_1V_1 = P_2V_2$ yields $P_1(M/D_1) = P_2(M/D_2)$, which simplifies to $\frac{P_1}{D_1} = \frac{P_2}{D_2}$) . **Units:** * **Pressure:** 1 atm = 101,325 Pa, 1 atm = 760 mm Hg = 760 torr . * **Volume:** SI unit is m³. Commonly used units include cm³, liters (l), and milliliters (ml). 1 ml = 1 cc, 1 l = 1000 ml . * **Temperature:** SI unit is Kelvin (K). Other units include Fahrenheit (℉) and Celsius (℃) . **Numerical Examples:** * **Example 1:** 250 ml of a gas at 650 mm pressure expands to 500 ml at the same temperature. What is the final pressure? Given: $P_1 = 650$ mm, $V_1 = 250$ ml, $V_2 = 500$ ml. Using $P_1V_1 = P_2V_2$: $650 \text{ mm} \times 250 \text{ ml} = P_2 \times 500 \text{ ml}$ $P_2 = \frac{650 \text{ mm} \times 250 \text{ ml}}{500 \text{ ml}} = 325 \text{ mm}$ . * **Example 2:** A weather balloon has a volume of 174 liters at 1 atm pressure. What is its volume when the atmospheric pressure is 0.80 atm, assuming constant temperature? Given: $V_1 = 174$ l, $P_1 = 1$ atm, $P_2 = 0.80$ atm. Using $P_1V_1 = P_2V_2$: $1 \text{ atm} \times 174 \text{ l} = 0.80 \text{ atm} \times V_2$ $V_2 = \frac{1 \text{ atm} \times 174 \text{ l}}{0.80 \text{ atm}} = 217.5 \text{ l}$ . * **Example 3:** The density of a gas is 32 at 760 mm pressure. What will be its density at 570 mm pressure, if the temperature is constant? Using $\frac{P_1}{D_1} = \frac{P_2}{D_2}$: $\frac{760 \text{ mm}}{32} = \frac{570 \text{ mm}}{D_2}$ $D_2 = \frac{32 \times 570 \text{ mm}}{760 \text{ mm}} = 24$ . * **Example 4:** A gas occupies 250 cc at 700 mm pressure and 25℃. What additional pressure is needed to reduce its volume to 4/5th of its original volume at the same temperature? Given: $V_1 = 250$ cc, $P_1 = 700$ mm. New volume $V_2 = \frac{4}{5} \times 250 \text{ cc} = 200$ cc. Using $P_1V_1 = P_2V_2$: $700 \text{ mm} \times 250 \text{ cc} = P_2 \times 200 \text{ cc}$ $P_2 = \frac{700 \text{ mm} \times 250 \text{ cc}}{200 \text{ cc}} = 875$ mm. Additional pressure required = $P_2 - P_1 = 875 \text{ mm} - 700 \text{ mm} = 175$ mm . * **Example 5:** A cylinder contains 30 liters of oxygen at 50 atm. How many gas jars, each of 400 cc capacity, can be filled from the cylinder at 750 mm pressure? Initial volume $V_1 = 30$ l $= 30,000$ cc. Initial pressure $P_1 = 50$ atm. Final pressure $P_2 = 750$ mm Hg. We need to convert atm to mm Hg. 1 atm = 760 mm Hg. So, $P_1 = 50 \text{ atm} \times 760 \frac{\text{mm Hg}}{\text{atm}} = 38000$ mm Hg. Using $P_1V_1 = P_2V_2$: $38000 \text{ mm Hg} \times 30000 \text{ cc} = 750 \text{ mm Hg} \times V_2$ $V_2 = \frac{38000 \text{ mm Hg} \times 30000 \text{ cc}}{750 \text{ mm Hg}} = 1,520,000$ cc. Number of gas jars = $\frac{\text{Total volume}}{\text{Volume per jar}} = \frac{1,520,000 \text{ cc}}{400 \text{ cc/jar}} = 3800$ jars . #### 5.2.2 Charles' law Charles' Law, named after Jacques Charles, describes the relationship between the volume (V) and temperature (T) of a given mass of gas at constant pressure . **Statement:** At a constant pressure, the volume of a given mass of a gas changes by $\frac{1}{273}$ of its original volume for every one degree Celsius rise or fall in temperature . **Mathematical Expression:** Let $V_0$ be the volume of a given mass of gas at 0℃. The volume at $t$℃ is given by: $V_{t℃} = V_0 + \frac{t}{273} V_0$ . $V_{t℃} = V_0 \left(1 + \frac{t}{273}\right)$ This can be rewritten as: $V_{t℃} = V_0 \left(\frac{273 + t}{273}\right)$ If $T$ is the temperature in Kelvin, where $T = 273 + t$, then the equation becomes: $V_{t℃} = V_0 \left(\frac{T}{273}\right)$ This implies that $V \propto T$ at constant pressure . Thus, Charles' Law can also be stated as: "At a constant pressure, the volume of a given mass of a gas is directly proportional to its absolute temperature (in Kelvin)" . For two different sets of conditions ($V_1, T_1$) and ($V_2, T_2$) at constant pressure: $\frac{V_1}{T_1} = \frac{V_2}{T_2}$ . **Absolute Temperature:** At -273℃, the volume of a gas theoretically becomes zero: $V_{-273℃} = V_0 \left(1 + \frac{-273}{273}\right) = V_0(1 - 1) = 0$ . This temperature, -273℃ or 0 Kelvin, is known as absolute zero, the theoretical temperature at which all molecular motion ceases and the properties of a gas vanish. However, gases typically liquefy or solidify before reaching this temperature . **Graphical Representation:** Charles' Law can be represented graphically by plotting volume against temperature (in ℃) at constant pressure, resulting in a curve called an isobar. Plotting volume against absolute temperature (in Kelvin) yields a straight line passing through the origin . > **Tip:** Always convert temperatures from Celsius to Kelvin by adding 273.15 (or approximately 273 for most calculations) when using Charles' Law . **Applications of Charles' Law:** * **Hot Air Balloons:** When air inside a hot air balloon is heated, its volume increases ($V \propto T$). This causes the density of the hot air to decrease relative to the cooler surrounding air, allowing the balloon to float . **Numerical Formulas and Units:** * $\frac{V_1}{T_1} = \frac{V_2}{T_2}$, where $T_1$ and $T_2$ are temperatures in Kelvin . **Conversion:** * Temperature in Kelvin = Temperature in ℃ + 273 . * Example: 10℃ = 10 + 273 = 283 K . * Example: -5℃ = -5 + 273 = 268 K . **Numerical Examples:** * **Example 1:** 200 cc of nitrogen gas at 27℃ is cooled to -20℃. Find the new volume, assuming constant pressure. Given: $V_1 = 200$ cc, $T_1 = 27℃ = 27 + 273 = 300$ K, $T_2 = -20℃ = -20 + 273 = 253$ K. Using $\frac{V_1}{T_1} = \frac{V_2}{T_2}$: $\frac{200 \text{ cc}}{300 \text{ K}} = \frac{V_2}{253 \text{ K}}$ $V_2 = \frac{200 \text{ cc} \times 253 \text{ K}}{300 \text{ K}} = 168.67 \text{ cc}$ . * **Example 2:** 400 ml of oxygen gas at -150℃ is heated to 20℃ at constant pressure. Find the new volume. Given: $V_1 = 400$ ml, $T_1 = -150℃ = -150 + 273 = 123$ K, $T_2 = 20℃ = 20 + 273 = 293$ K. Using $\frac{V_1}{T_1} = \frac{V_2}{T_2}$: $\frac{400 \text{ ml}}{123 \text{ K}} = \frac{V_2}{293 \text{ K}}$ $V_2 = \frac{400 \text{ ml} \times 293 \text{ K}}{123 \text{ K}} = 952.85 \text{ ml}$ . * **Example 3:** 400 ml of N2 gas at 27℃ is cooled to -5℃ without change in pressure. Calculate the contraction in volume. Given: $V_1 = 400$ ml, $T_1 = 27℃ = 300$ K, $T_2 = -5℃ = 268$ K. First, find $V_2$: $\frac{400 \text{ ml}}{300 \text{ K}} = \frac{V_2}{268 \text{ K}}$ $V_2 = \frac{400 \text{ ml} \times 268 \text{ K}}{300 \text{ K}} = 357.33$ ml. Contraction in volume = $V_1 - V_2 = 400 \text{ ml} - 357.33 \text{ ml} = 42.67$ ml . * **Example 4:** When a vessel containing 400 cc of air at 7℃ is heated to 27℃ at the same pressure, what volume of air will be increased? Given: $V_1 = 400$ cc, $T_1 = 7℃ = 280$ K, $T_2 = 27℃ = 300$ K. First, find $V_2$: $\frac{400 \text{ cc}}{280 \text{ K}} = \frac{V_2}{300 \text{ K}}$ $V_2 = \frac{400 \text{ cc} \times 300 \text{ K}}{280 \text{ K}} = 428.57$ cc. Increase in volume = $V_2 - V_1 = 428.57 \text{ cc} - 400 \text{ cc} = 28.57$ cc . * **Example 5:** What fraction of air would have been expelled out when an open vessel at a temperature of 25℃ was heated to 400℃ at constant pressure? Let initial volume be $V_1$ at $T_1 = 25℃ = 298$ K. Let final volume be $V_2$ at $T_2 = 400℃ = 673$ K. From Charles' Law, $\frac{V_1}{T_1} = \frac{V_2}{T_2}$, so $V_2 = V_1 \frac{T_2}{T_1} = V_1 \frac{673 \text{ K}}{298 \text{ K}} = 2.258 V_1$. The problem implies that the vessel is open, so as it heats up, the volume of air *inside* the vessel expands. If the vessel's capacity is fixed, then air is expelled. However, the question is asking about the *fraction of air expelled* from an open vessel, suggesting we consider the expansion relative to the initial volume. The wording is a bit ambiguous. Assuming the question means: If a fixed amount of air occupies volume $V_1$ at $T_1$, what volume $V_2$ does it occupy at $T_2$? The difference $V_2 - V_1$ represents the *expansion*. The *fraction expelled* from an open vessel could be interpreted as the *additional volume* compared to the initial volume. Let's re-interpret: If we have a certain volume of air at $T_1$, and heat it to $T_2$. The *new volume* it would occupy is $V_2$. If the original volume was $V_1$, the *amount expelled* is related to the increase in volume. Using the provided answer: 0.556. Let's assume $V_1$ is the initial volume at $T_1$. The amount expelled is the volume that *would have been* occupied by this air at $T_2$ if it remained at $T_1$, but is now at $T_2$. If $V_1$ is the volume at $T_1 = 298$ K, then at $T_2 = 673$ K, the new volume would be $V_2 = V_1 \frac{673}{298}$. The fraction expelled could be $\frac{V_2 - V_1}{V_1} = \frac{V_2}{V_1} - 1 = \frac{T_2}{T_1} - 1 = \frac{673}{298} - 1 = 2.258 - 1 = 1.258$. This does not match the answer. Let's assume the question means what fraction of the *final volume* was added. Alternatively, if we consider the initial amount of air as $n_1$ and final as $n_2$. At constant pressure and volume of the vessel: This question seems to imply that the vessel itself is not the constraint, but rather we are looking at the expansion of a quantity of air. Let's consider the *fraction expelled* from an open vessel. If a vessel is open and contains $V_1$ volume of air at $T_1$. When heated to $T_2$, the air expands. The amount of air that *escapes* is the amount corresponding to the volume difference. If $V_1$ is the initial volume at $T_1=298$ K. If the air were to occupy volume at $T_2=673$ K, its volume would be $V_2 = V_1 \times \frac{673}{298}$. The fraction of air expelled is the ratio of the increase in volume to the final volume, if the vessel could accommodate that. Let's consider the fraction of the *initial volume* that is expelled. Fraction expelled = $\frac{\text{Volume at } T_2 \text{ corresponding to } V_1 \text{ air} - V_1}{V_1}$ Fraction expelled = $\frac{V_2 - V_1}{V_1} = \frac{T_2}{T_1} - 1 = \frac{673}{298} - 1 \approx 1.258$. Still not matching. Let's consider the fraction of the initial amount of air that is expelled. If $n_1$ moles of air are present at $T_1$, and $n_2$ moles are present at $T_2$ within the same volume $V$ (if the vessel were closed). Since it's open, the air expands. The amount of air *remaining* in the vessel at $T_2$ is the amount that would occupy volume $V_1$ at $T_2$. Amount at $T_2$ remaining = $n_{rem} \propto V_1$ (at $T_2$) Initial amount = $n_1 \propto V_1$ (at $T_1$) Amount expelled = $n_1 - n_{rem}$. Fraction expelled = $\frac{n_1 - n_{rem}}{n_1} = 1 - \frac{n_{rem}}{n_1}$. Since $n \propto V/T$, and $V$ is constant in this interpretation. Let's consider the amount of air in moles. $PV = nRT$. At constant P, $V/T = nR$. So $n \propto V/T$. If the vessel has a fixed volume $V$. Initial moles $n_1 \propto V/T_1$. Final moles $n_2 \propto V/T_2$. Number of moles expelled $= n_1 - n_2$. Fraction expelled $= \frac{n_1 - n_2}{n_1} = 1 - \frac{n_2}{n_1} = 1 - \frac{V/T_2}{V/T_1} = 1 - \frac{T_1}{T_2}$. Fraction expelled $= 1 - \frac{298 \text{ K}}{673 \text{ K}} = 1 - 0.44279 = 0.5572$. This is very close to the answer 0.556. So the fraction of air expelled from an open vessel when heated is $1 - \frac{T_{initial}}{T_{final}}$. Fraction expelled = $1 - \frac{25 + 273}{400 + 273} = 1 - \frac{298}{673} \approx 0.557$ . * **Example 6:** How much increase in temperature is required to increase the volume of half a liter of gas by 30% at 27℃ at constant pressure? Given: Initial volume $V_1 = 0.5$ l, $T_1 = 27℃ = 300$ K. The volume is increased by 30%, so the new volume $V_2 = V_1 + 0.30 V_1 = 1.30 V_1$. $V_2 = 1.30 \times 0.5 \text{ l} = 0.65$ l. We need to find the final temperature $T_2$. Using $\frac{V_1}{T_1} = \frac{V_2}{T_2}$: $\frac{0.5 \text{ l}}{300 \text{ K}} = \frac{0.65 \text{ l}}{T_2}$ $T_2 = \frac{0.65 \text{ l} \times 300 \text{ K}}{0.5 \text{ l}} = 390$ K. Increase in temperature = $T_2 - T_1 = 390 \text{ K} - 300 \text{ K} = 90$ K . * **Example 7:** A balloon can hold 1000 cc of air before bursting. It contains 900 cc of air at -1℃. Will it burst at 27℃? If not, find the temperature above which the balloon bursts. Given: Maximum volume $V_{max} = 1000$ cc. Initial conditions: $V_1 = 900$ cc, $T_1 = -1℃ = 272$ K. We need to find the volume at $T_2 = 27℃ = 300$ K. Using $\frac{V_1}{T_1} = \frac{V_2}{T_2}$: $\frac{900 \text{ cc}}{272 \text{ K}} = \frac{V_2}{300 \text{ K}}$ $V_2 = \frac{900 \text{ cc} \times 300 \text{ K}}{272 \text{ K}} = 992.65$ cc. Since $V_2 (992.65 \text{ cc}) < V_{max} (1000 \text{ cc})$, the balloon will not burst at 27℃ . To find the temperature above which the balloon bursts, we set $V_2 = V_{max} = 1000$ cc and use the initial conditions ($V_1 = 900$ cc, $T_1 = 272$ K) to find the bursting temperature $T_{burst}$. $\frac{V_1}{T_1} = \frac{V_{burst}}{T_{burst}}$ $\frac{900 \text{ cc}}{272 \text{ K}} = \frac{1000 \text{ cc}}{T_{burst}}$ $T_{burst} = \frac{1000 \text{ cc} \times 272 \text{ K}}{900 \text{ cc}} = 302.22$ K. Convert this back to Celsius: $302.22 \text{ K} - 273 = 29.22℃$. So, the balloon bursts at temperatures above 29.22℃ . --- ## Common mistakes to avoid - Review all topics thoroughly before exams - Pay attention to formulas and key definitions - Practice with examples provided in each section - Don't memorize without understanding the underlying concepts

| Term | Definition | |------|------------| | Atomic radius | The atomic size can be determined by knowing the distance between the atoms in the combined state. It is a measure of the size of an atom. | | Covalent radius | It is defined as half the distance between the nuclei of two identical atoms that are bonded together covalently. It represents the size of an atom in a covalent bond. | | Ionic radius | It is the distance between the nuclei of neighboring cation and anion in an ionic compound. It reflects the size of an ion in a crystal lattice. | | Effective nuclear charge | The net positive charge experienced by an electron in a polyelectronic atom, resulting from the attraction of the nucleus and the repulsion from other electrons. | | Isoelectronic species | Atoms and ions that have the same electronic configuration and thus the same number of electrons. | | Ionization potential (or energy/enthalpy) | The minimum energy required to remove the most loosely bound electron from an isolated gaseous atom or ion. It indicates how strongly an electron is held by the nucleus. | | Electron Affinity | The amount of energy released when an electron is added to a neutral gaseous atom to form a negative ion. It signifies the tendency of an atom to accept an electron. | | Electronegativity | The tendency of an atom to attract a shared pair of electrons towards itself in a covalent bond. It influences the polarity of chemical bonds. | | Metallurgy | The process of extraction of metals in a pure form from their ores. It involves a series of steps to isolate and purify metals. | | Ore | A mineral containing a valuable deposit of ore, from which a metal can be economically extracted. | | Gangue (or Matrix) | The unwanted earthy and siliceous impurities associated with ores. Removal of gangue is crucial in ore concentration. | | Pyrometallurgy | A metallurgical process where extraction of metals occurs at very high temperatures, often involving smelting or roasting. | | Hydrometallurgy | A process for extracting metals using aqueous solutions, often involving leaching and precipitation. | | Electrometallurgy | The extraction of metals from molten salts or solutions using electrolytic methods. | | Calcination | A process of heating an ore strongly below its melting point in the absence or limited supply of air, typically used to remove volatile impurities and moisture, or to convert carbonates and hydroxides to oxides. | | Roasting | A process of heating an ore strongly below its melting point in excess of air, commonly used to convert sulphide ores to their respective metal oxides and remove volatile impurities. | | Smelting | The process of extracting a metal from its oxide by reduction, typically using carbon or carbon monoxide at high temperatures in a blast furnace. | | Refining | The process of purifying a crude metal obtained after extraction to remove remaining impurities and obtain a pure metal. | | Electrolytic refining | A refining method where the impure metal acts as the anode and a pure metal strip as the cathode in an electrolytic cell, separating the pure metal through electrolysis. | | Poling | A refining method used for crude metals containing oxide impurities, involving stirring the molten metal with green poles of wood which reduce the oxide impurities. | | Alkali metals | The elements of Group 1 (IA) of the periodic table: lithium (Li), sodium (Na), potassium (K), rubidium (Rb), cesium (Cs), and francium (Fr). They are highly reactive and form alkalis when reacting with water. | | Alkaline earth metals | The elements of Group 2 (IIA) of the periodic table: beryllium (Be), magnesium (Mg), calcium (Ca), strontium (Sr), barium (Ba), and radium (Ra). They are reactive metals that form basic oxides and hydroxides. | | Deliquescent | A substance that absorbs moisture from the air to such an extent that it dissolves in the absorbed water, forming an aqueous solution. | | Amphoteric | A compound that can act as both an acid and a base. For example, some metal hydroxides are amphoteric and can dissolve in excess strong acid or strong base. | | Bioinorganic Chemistry | The study of the role of inorganic substances in biological systems, focusing on the interactions between inorganic chemistry and biochemistry. | | Macronutrients | Elements required by living organisms in large amounts for essential life processes, such as Na, K, Mg, Ca, and Cl. | | Micronutrients (Trace elements) | Elements required by living organisms in small quantities but are still vital for biological systems, such as Fe, Zn, Cu, and Co. | | Sodium-Potassium Pump (Na+-K+ pump) | A primary active transport system that moves sodium ions out of the cell and potassium ions into the cell across the cell membrane, utilizing ATP. | | Sodium-Glucose Pump | A secondary active transport system that uses the concentration gradient of Na+ ions (established by the Na+-K+ pump) to drive the transport of glucose into the cell. | | Metal toxicity | The toxic effect of metals in certain forms and doses on living organisms, which can result from exposure through food, air, water, or industrial sources. | | Boyle's Law | A gas law stating that at constant temperature, the volume of a given mass of a gas is inversely proportional to its pressure. Mathematically, $PV = k$. | | Isotherm | A curve plotted on a graph representing the relationship between two variables (like pressure and volume) at a constant temperature. | | Charles' Law | A gas law stating that at constant pressure, the volume of a given mass of a gas is directly proportional to its absolute temperature (in Kelvin). Mathematically, $V/T = k$. | | Isobar | A curve plotted on a graph representing the relationship between two variables (like volume and temperature) at a constant pressure. | | Absolute zero | The theoretical temperature at which the volume of an ideal gas would become zero, approximately -273.15 degrees Celsius or 0 Kelvin. |