CBI.Sum.notes -1-7.pdf

# Atomic structure and quantum mechanics This section outlines the fundamental principles governing the structure of atoms and the behavior of electrons within them, as described by quantum mechanics. ### 1.1 The atom: fundamental components An atom is composed of a central nucleus surrounded by electrons [1](#page=1). * **Nucleus:** Contains protons (positively charged, +1) and neutrons (neutral, 0). The number of protons determines the nucleus's overall positive charge. Protons and neutrons have approximately the same mass [1](#page=1). * **Atomic Number ($Z$)**: The number of protons in an atom's nucleus. This number defines the element and dictates 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). * **Electrons:** Negatively charged particles orbiting the nucleus. Their mass is approximately 1/2000th of a proton's mass, contributing negligibly to the atomic mass [1](#page=1). The **unified atomic mass unit (u or Da)** is defined as 1/12th the mass of a Carbon-12 atom, and is roughly equivalent to the mass of a proton or neutron (approximately $1.66 \times 10^{-27}$ kg) [1](#page=1). ### 1.2 The periodic table and atomic properties The periodic table organizes elements, currently ordered by atomic number ($Z$) [1](#page=1). * **Groups (Columns 1-18):** Elements within the same group share similar numbers of valence electrons and exhibit comparable chemical properties [1](#page=1). * **Periods (Rows 1-7):** Indicate the number of electron shells an atom possesses [1](#page=1). Key atomic properties and trends include: * **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 nucleus's center to the atom's outer shell [1](#page=1). * **Electronegativity:** An atom's ability to attract electrons to itself, often estimated by averaging ionization energy and electron affinity. This property helps determine bond type [1](#page=1). ### 1.3 Isotopes Isotopes are atoms of the same element (same atomic number) with differing numbers of neutrons, resulting in different atomic mass numbers. Despite differing neutron counts, isotopes of an element have the same number of electrons, leading to similar chemical behavior. The atomic mass listed on the periodic table is the average of the masses of an element's naturally occurring isotopes [1](#page=1). ### 1.4 Atomic models and quantum mechanics * **Bohr Model:** Successfully described the hydrogen atom, depicting electrons in specific orbits around a positive nucleus, but failed for atoms with more than one electron [1](#page=1). * **Quantum Mechanical Model:** Treats electrons as having wave-like properties (matter waves). The **Schrödinger wave equation** describes the behavior of these matter waves [1](#page=1). * **Wave Functions ($\psi$)**: Solutions to the Schrödinger equation, used to define **probability density** – the likelihood of finding an electron in a specific region around an atom [1](#page=1). * **Atomic Orbital:** A region in an atom where there is a 90% probability of finding an electron [1](#page=1). ### 1.5 Quantum numbers and atomic orbitals Atomic orbitals are characterized by four quantum numbers: * **Principal Quantum Number ($n$)**: Defines the electron shell and determines the size and energy of an atomic orbital. Values are positive integers ($1, 2, 3, \dots$). Higher $n$ values correspond to larger orbitals and higher energy levels [1](#page=1). * **Orbital (Azimuthal) Quantum Number ($\ell$)**: Defines the subshell and the shape of the orbital. For a given $n$, $\ell$ can range from $0$ to $(n-1)$ [1](#page=1). * $\ell = 0$ corresponds to the 's' subshell (spherical shape) [1](#page=1). * $\ell = 1$ corresponds to the 'p' subshell (dumbbell shape with lobes along axes) [1](#page=1). * $\ell = 2$ corresponds to the 'd' subshell (more complex shapes, often lobed) [1](#page=1). * $\ell = 3$ corresponds to the 'f' subshell (even more complex shapes) [1](#page=1). * **Magnetic Quantum Number ($m_{\ell}$)**: Specifies the number and orientation of orbitals within a subshell [1](#page=1). * For a given $\ell$, $m_{\ell}$ ranges from $-\ell$ to $+\ell$ [1](#page=1). * s subshell ($\ell=0$): 1 orbital ($m_{\ell}=0$) [1](#page=1). * p subshell ($\ell=1$): 3 orbitals ($m_{\ell}=-1, 0, +1$) [1](#page=1). * d subshell ($\ell=2$): 5 orbitals ($m_{\ell}=-2, -1, 0, +1, +2$) [1](#page=1). * f subshell ($\ell=3$): 7 orbitals ($m_{\ell}=-3, -2, -1, 0, +1, +2, +3$) [1](#page=1). * **Degenerate Orbitals:** Orbitals within the same subshell have the same energy level, regardless of their orientation (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 have two possible values: $+\frac{1}{2}$ or $-\frac{1}{2}$ [1](#page=1). > **Tip:** Remember that the number of subshells in a shell is equal to the principal quantum number ($n$). For $n=3$, there are three subshells: 3s, 3p, and 3d, corresponding to $\ell$ values of 0, 1, and 2, respectively [7](#page=7). ### 1.6 Electron configuration principles Electron configuration describes the arrangement of electrons in atomic orbitals. * **Pauli Exclusion Principle:** No two electrons in an atom can have the same set of four quantum numbers. This means each orbital can hold a maximum of two electrons, and they must have opposite spins [1](#page=1). * **Aufbau Principle:** Electrons fill atomic orbitals in order of increasing energy [1](#page=1) [2](#page=2). * **Hund's Rule:** Within a subshell, electrons will occupy individual orbitals singly before pairing up. Unpaired electrons in the same subshell will have the same spin [1](#page=1) [2](#page=2). **Example Electron Configurations:** * **Hydrogen (H):** 1 electron. Configuration: $1s^1$ [2](#page=2). * **Nitrogen (N):** 7 electrons. Configuration: $1s^2, 2s^2, 2p^3$. The three electrons in the 2p subshell are in separate orbitals with parallel spins, according to Hund's rule [2](#page=2). **Shorthand Notation:** The electron configuration of noble gases represents the inner shell electrons, simplifying the notation for outer shell electrons. For example, Potassium (K) (19 electrons) is $[Ar 4s^1$ [2](#page=2). ### 1.7 Chemical Bonding * **Ionic Bonds:** Formed between atoms with a large electronegativity difference (typically > 2), involving the transfer of electrons to form ions. The electrostatic attraction between oppositely charged ions holds them together [2](#page=2). * **Cation:** A positively charged ion [1](#page=1). * **Anion:** A negatively charged ion [1](#page=1). * Cations are generally smaller than their parent atoms, while anions are larger due to increased electron-electron repulsion [1](#page=1). * **Covalent Bonds:** Formed when atoms share pairs of electrons, typically between non-metals. The degree of sharing can be unequal if electronegativity differs, leading to polar covalent bonds [1](#page=1) [2](#page=2). ### 1.8 Molecular Geometry (VSEPR Theory) **Valence Shell Electron Pair Repulsion (VSEPR) Theory** predicts the geometry of molecules based on minimizing electron repulsion around a central atom. Electron pairs (bonding and lone pairs) arrange themselves as far apart as possible [2](#page=2). * **Methane (CH$_4$):** Tetrahedral geometry due to four bonding pairs and no lone pairs on the central carbon atom [2](#page=2). * **Ammonia (NH$_3$):** Tetrahedral electron geometry, but trigonal pyramidal molecular geometry due to one lone pair on nitrogen [2](#page=2). * **Water (H$_2$O):** Tetrahedral electron geometry, but bent molecular geometry due to two lone pairs on oxygen. The repulsion order is: Lone Pair-Lone Pair (LP-LP) > Lone Pair-Bonding Pair (LP-BP) > Bonding Pair-Bonding Pair (BP-BP) [2](#page=2). --- # Chemical bonding and molecular geometry This section explores the fundamental principles governing how atoms bond together to form molecules and how the arrangement of these atoms dictates the molecule's three-dimensional structure. ### 2.1 Types of chemical bonds Chemical bonds are the forces that hold atoms together in molecules and compounds. The primary types of chemical bonds are ionic and covalent. #### 2.1.1 Ionic bonds Ionic bonds are formed when there is a large difference in electronegativity between two atoms, typically between a metal and a non-metal. This large electronegativity difference (greater than 2) causes one atom to effectively transfer an electron to another, creating ions. The positively charged ion (cation) and negatively charged ion (anion) are then held together by strong electrostatic attraction in a crystal lattice structure. Cations are generally smaller than their neutral atomic counterparts because they have fewer electrons to repel each other. Conversely, anions are larger due to increased electron-electron repulsion [2](#page=2). #### 2.1.2 Covalent bonds Covalent bonds involve the sharing of electron pairs between atoms, typically between non-metal atoms. These shared electrons belong to the valence shells of both participating atoms. The electrons that are shared are called bonding pairs, while electrons not involved in sharing are called lone pairs [2](#page=2). * **Polar Covalent Bonds:** When atoms with different electronegativities form a covalent bond, they share the bonding pair of electrons unequally. The atom with higher electronegativity attracts the electron pair more strongly, resulting in a partial negative charge ($ \delta^- $) on that atom and a partial positive charge ($ \delta^+ $) on the less electronegative atom. This uneven distribution of charge creates a permanent dipole within the molecule. For instance, in water ($ \text{H}_2\text{O} $), the oxygen atom is more electronegative than hydrogen, leading to polar O-H bonds and an overall polar molecule due to its bent geometry. Carbon dioxide ($ \text{CO}_2 $), despite having polar C=O bonds, is nonpolar overall because its linear geometry causes the bond dipoles to cancel out [2](#page=2). * **Nonpolar Covalent Bonds:** When two atoms with equal electronegativity form a covalent bond, the electron pair is shared equally, resulting in a nonpolar covalent bond [2](#page=2). ### 2.2 Theories of Chemical Bonding Two key theories explain the nature of chemical bonding: Valence Bond Theory and Molecular Orbital Theory. #### 2.2.1 Valence Bond Theory Valence Bond Theory focuses on the combination and overlap of atomic orbitals to form chemical bonds [2](#page=2). * **Overlap:** In covalent bonding, the atomic orbitals of the participating atoms overlap, allowing them to share electrons. Greater overlap generally leads to a stronger bond [2](#page=2). * **Sigma ($ \sigma $) Bonds:** These are formed by the end-to-end overlap of atomic orbitals along the internuclear axis. Sigma bonds can form from the overlap of s-s, s-p, p-p, or hybridized orbitals [2](#page=2). * **Pi ($ \pi $) Bonds:** These are formed by the sideways (parallel) overlap of atomic orbitals, usually p orbitals, above and below or in front and behind the sigma bond axis. Pi bonds are generally weaker than sigma bonds due to less effective overlap and restrict rotation around the bond axis. Multiple bonds (double and triple bonds) consist of one sigma bond and one or two pi bonds, respectively [2](#page=2). #### 2.2.2 Hybridization Hybridization is a theoretical concept where atomic orbitals within an atom mix to form new hybrid orbitals that are better suited for bonding, explaining observed molecular geometries and bond properties. For example, in methane ($ \text{CH}_4 $), the carbon atom's 2s and three 2p orbitals hybridize to form four degenerate $ \text{sp}^3 $ hybrid orbitals. These $ \text{sp}^3 $ orbitals are arranged tetrahedrally, allowing for the formation of four identical C-H bonds with equal energy and length [2](#page=2). The type of hybridization determines the geometry: * Two regions of electron density: $ \text{sp} $ hybridization, leading to a linear geometry [2](#page=2). * Three regions of electron density: $ \text{sp}^2 $ hybridization, leading to a trigonal planar geometry [2](#page=2). * Four regions of electron density: $ \text{sp}^3 $ hybridization, leading to a tetrahedral geometry [2](#page=2). #### 2.2.3 Molecular Orbital Theory Molecular Orbital Theory (MO Theory) offers an alternative explanation for bonding by considering the combination of atomic orbitals to form molecular orbitals that encompass entire molecules. MO theory can accurately predict bond lengths and energies. While effective for simple molecules, it becomes complex for larger ones [2](#page=2). ### 2.3 Molecular Geometry and VSEPR Theory Valence Shell Electron Pair Repulsion (VSEPR) theory is used to predict the shapes of molecules based on the repulsion between electron pairs (both bonding and lone pairs) around a central atom. Electron pairs arrange themselves to minimize repulsion, leading to specific geometries [2](#page=2). * **Electron Regions:** The number of regions of electron density (bonding pairs and lone pairs) around the central atom determines the electron geometry. * 2 regions: Linear * 3 regions: Trigonal planar * 4 regions: Tetrahedral * 5 regions: Trigonal bipyramidal * 6 regions: Octahedral * **Molecular Geometry:** This describes the arrangement of atoms only, ignoring lone pairs for the final shape prediction. * **Methane ($ \text{CH}_4 $):** Central carbon with 4 bonding pairs (no lone pairs) leads to a tetrahedral electron geometry and molecular geometry [2](#page=2). * **Ammonia ($ \text{NH}_3 $):** Central nitrogen with 3 bonding pairs and 1 lone pair (4 regions total) results in a tetrahedral electron geometry, but a trigonal pyramidal molecular geometry [2](#page=2). * **Water ($ \text{H}_2\text{O} $):** Central oxygen with 2 bonding pairs and 2 lone pairs (4 regions total) results in a tetrahedral electron geometry, but a bent molecular geometry [2](#page=2). **Repulsion Hierarchy:** The repulsion between electron pairs follows this order: Lone Pair-Lone Pair (LP-LP) > Lone Pair-Bonding Pair (LP-BP) > Bonding Pair-Bonding Pair (BP-BP). This hierarchy influences bond angles, with lone pairs compressing bond angles slightly compared to ideal geometries [2](#page=2). ### 2.4 Electronegativity and Bond Polarity Electronegativity is the measure of an atom's ability to attract electrons to itself within a chemical bond. The difference in electronegativity between two bonded atoms is crucial for determining the type and polarity of the bond [2](#page=2). * **Ionic Bonds:** Formed when the electronegativity difference is large (typically > 2) [2](#page=2). * **Polar Covalent Bonds:** Formed when there is a moderate electronegativity difference, leading to unequal sharing of electrons and a partial charge separation across the bond [2](#page=2). * **Nonpolar Covalent Bonds:** Formed when the electronegativity difference is zero or very small, resulting in equal sharing of electrons [2](#page=2). Fajans' Rules describe when ionic compounds gain covalent character: 1. **Small and highly charged cation:** This strongly distorts the electron cloud of the anion, increasing covalent character [2](#page=2). 2. **Large and highly charged anion:** Its electron cloud is more easily distorted [2](#page=2). 3. **High charge density of the cation:** Leads to greater polarizing power [2](#page=2). ### 2.5 Bond Strength and Length Generally, as bond length increases, bond strength decreases. Longer bonds are weaker. Greater 's' character in hybridization leads to shorter, stronger bonds, while greater 'p' character leads to longer, weaker bonds [2](#page=2). --- # Thermodynamics and chemical kinetics Thermodynamics and chemical kinetics are fundamental concepts that govern the energy changes and rates of chemical reactions. ### 3.1 The laws of thermodynamics Thermodynamics provides a framework for understanding energy transfer and transformation within chemical systems. #### 3.1.1 First law of thermodynamics The first law of thermodynamics, also known as the law of conservation of energy, states that energy cannot be created or destroyed, only transferred or changed from one form to another. In any isolated system (or the system and its surroundings combined), the total energy remains constant. This means energy can be converted from potential energy (stored energy, such as in chemical bonds) to kinetic energy (energy of motion) and vice versa. Thermodynamic temperature, measured in Kelvin (K), is a measure of the average kinetic energy of particles within a system [5](#page=5) [6](#page=6) [7](#page=7). * **Heat:** Energy transferred due to a temperature difference. It flows spontaneously from hotter to colder objects. Heat absorbed is positive ($+ve$), while heat released is negative ($-ve$) [5](#page=5) [6](#page=6) [7](#page=7). * **Enthalpy (H):** Represents the total energy of a chemical system, encompassing kinetic (heat) energy and chemical potential energy stored in bonds. The change in enthalpy ($\Delta H$) at constant pressure is equal to the heat absorbed or released by the system [5](#page=5) [6](#page=6) [7](#page=7). * **Exothermic reactions:** Release heat, resulting in a negative enthalpy change ($\Delta H < 0$). Products have lower chemical potential energy than reactants [5](#page=5) [6](#page=6) [7](#page=7). * **Endothermic reactions:** Absorb heat, resulting in a positive enthalpy change ($\Delta H > 0$). Products have greater chemical potential energy than reactants [5](#page=5) [6](#page=6) [7](#page=7). * **Bond energy:** The average energy required to break one mole of a specific type of bond. The sum of bond energies can be used to estimate enthalpy changes in reactions [5](#page=5) [6](#page=6) [7](#page=7). #### 3.1.2 Second law of thermodynamics The second law of thermodynamics states that the total entropy of the universe always increases over time [5](#page=5) [6](#page=6) [7](#page=7). * **Entropy (S):** A thermodynamic measure of energy dispersal or randomness within a system. An increase in the number of ways energy can be distributed leads to an increase in entropy. A perfect solid at absolute zero has zero entropy [5](#page=5) [6](#page=6) [7](#page=7). #### 3.1.3 Gibbs free energy (G) Gibbs free energy ($G$) is a thermodynamic potential that combines enthalpy and entropy changes to predict the spontaneity of a process [5](#page=5) [6](#page=6) [7](#page=7). The change in Gibbs free energy ($\Delta G$) is calculated as: $$ \Delta G = \Delta H - T\Delta S $$ where $\Delta H$ is the change in enthalpy, $T$ is the temperature (in Kelvin), and $\Delta S$ is the change in entropy [5](#page=5) [6](#page=6) [7](#page=7). * If $\Delta G < 0$, the reaction is spontaneous and exergonic (releases free energy) [5](#page=5) [6](#page=6) [7](#page=7). * If $\Delta G > 0$, the reaction is non-spontaneous and endergonic (absorbs free energy) [5](#page=5) [6](#page=6) [7](#page=7). * If $\Delta G = 0$, the reaction is at equilibrium [5](#page=5) [6](#page=6) [7](#page=7). A spontaneous and exergonic reaction is favoured when $\Delta H$ is negative (exothermic) and $\Delta S$ is positive (increasing disorder) [5](#page=5) [6](#page=6) [7](#page=7). ### 3.2 Chemical kinetics Chemical kinetics studies the rates of chemical reactions. #### 3.2.1 Reaction rates and activation energy The rate of a reaction is influenced by several factors, including temperature, concentration, and the presence of catalysts. * **Transition state:** The highest energy point along the reaction pathway, representing an unstable arrangement of atoms where bonds are breaking and forming [5](#page=5) [6](#page=6) [7](#page=7). * **Activation energy ($E_a$):** The minimum energy required for a reaction to occur, overcoming the energy barrier of the transition state. Reactions with higher activation energies proceed more slowly [5](#page=5) [6](#page=6) [7](#page=7). * **Activated complex:** A general term for the arrangement of atoms at the transition state [5](#page=5) [6](#page=6) [7](#page=7). #### 3.2.2 Catalysts Catalysts are substances that increase the rate of a chemical reaction without being consumed in the process. They achieve this by providing an alternative reaction pathway with a lower activation energy ($E_a$). Catalysts do not alter the position of equilibrium but allow it to be reached faster. Enzymes are biological catalysts [5](#page=5) [6](#page=6) [7](#page=7). ### 3.3 Relationship between thermodynamics and kinetics Thermodynamics predicts whether a reaction is possible (spontaneity) based on free energy changes ($\Delta G$), while kinetics determines how fast a reaction will occur based on activation energy ($E_a$) and other kinetic factors. A reaction can be thermodynamically favourable but kinetically slow if it has a high activation energy [5](#page=5) [6](#page=6) [7](#page=7). --- # Biomolecules: structure and function Biomolecules are the fundamental organic molecules that comprise living organisms, playing crucial roles in cellular structure, energy storage, and biological processes [4](#page=4). ### 4.1 Carbohydrates Carbohydrates, often referred to as sugars, are a major class of biomolecules with the general formula $(CH_2O)_n$. Their primary functions include serving as energy stores, fuels, metabolic substrates, and structural components within cells [6](#page=6). #### 4.1.1 Monosaccharides Monosaccharides are the simplest form of carbohydrates and are classified based on their functional group: * **Aldoses:** Contain an aldehyde group (R-CH=O) at the end of the carbon chain. Examples include glucose and glyceraldehyde [6](#page=6). * **Ketoses:** Contain a ketone group (R-C(=O)-R') in the middle of the carbon chain. Examples include fructose and dihydroxyacetone [6](#page=6). Fischer projections are used to represent the three-dimensional structure of monosaccharides in a two-dimensional format, indicating the arrangement of groups around chiral carbon atoms [6](#page=6). In aqueous solutions, monosaccharides can exist in both open-chain and cyclic (ring) forms. Cyclization occurs when a hydroxyl (-OH) group within the molecule attacks the carbonyl (C=O) group, forming a hemiacetal (from aldoses) or hemiketal (from ketoses). This process creates a new stereocenter at the anomeric carbon, leading to two possible configurations [6](#page=6): * **Alpha (α) anomer:** The hydroxyl group at C1 is on the opposite side of the ring relative to the C5 or C4 group [6](#page=6). * **Beta (β) anomer:** The hydroxyl group at C1 is on the same side of the ring relative to the C5 or C4 group [6](#page=6). Haworth projections are commonly used to depict these cyclic forms, with both alpha and beta anomers typically existing in equilibrium [6](#page=6). #### 4.1.2 Disaccharides Disaccharides are formed when two monosaccharides are linked together via a covalent **glycosidic bond** through a condensation reaction, releasing a molecule of water. The name of the disaccharide reflects the constituent monosaccharides and the specific glycosidic linkage. For instance, sucrose has an $\alpha, \beta$-1,2 glycosidic bond where both anomeric carbons are involved. Lactose, however, has a $\beta$-1,4 glycosidic bond, involving only one anomeric carbon, allowing for the presence of both anomers simultaneously in solution [6](#page=6). #### 4.1.3 Polysaccharides Polysaccharides are long chains of monosaccharides linked by glycosidic bonds. The type of monosaccharide and the linkage determine the polysaccharide's structure and function. Examples include [6](#page=6): * **Cellulose:** Formed by $\beta$-D-glucose units linked via $\beta$-1,4 glycosidic bonds, providing structural support [6](#page=6). * **Starch and Glycogen:** Composed of $\alpha$-D-glucose units linked via $\alpha$-1,4 glycosidic bonds, serving as energy storage. Starch contains amylose (unbranched $\alpha$-1,4 linkages) and amylopectin (branched $\alpha$-1,6 linkages), while glycogen is highly branched amylopectin [6](#page=6). Carbohydrates also function in post-translational modifications, such as in the ABO blood groups [6](#page=6). ### 4.2 Lipids Lipids are a diverse group of molecules characterized by their insolubility in water and solubility in nonpolar solvents. They play vital roles in energy storage, membrane structure, and signaling [6](#page=6). #### 4.2.1 Fatty Acids Fatty acids are long hydrocarbon chains with a terminal carboxylic acid group [-COOH. They are classified based on saturation [6](#page=6): * **Saturated Fatty Acids (SFAs):** Contain only single bonds between carbon atoms in the hydrocarbon chain [6](#page=6). * **Unsaturated Fatty Acids (UFAs):** Contain one or more carbon-carbon double bonds ($\text{C=C}$). The stereochemistry of these double bonds (cis/Z or trans/E) significantly impacts their properties. Cis double bonds introduce a kink or bend in the hydrocarbon chain, affecting how fatty acids pack in membranes [6](#page=6). Fatty acids are 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 denoted as "number of carbons:number of double bonds," e.g., 16:0 for palmitic acid, and the position of double bonds can be indicated from either end of the chain [6](#page=6). #### 4.2.2 Triacylglycerols Triacylglycerols (triglycerides) are formed by the esterification of glycerol with three fatty acids. They serve as a major form of chemical energy storage in adipose tissue [6](#page=6). #### 4.2.3 Glycerophospholipids Glycerophospholipids are key components of cellular membranes. They consist of a glycerol backbone esterified to two fatty acids and a phosphate group, which is further linked to a polar head group (e.g., choline, serine, inositol). These molecules are amphiphilic, possessing both hydrophobic tails and hydrophilic heads, enabling them to form lipid bilayers in membranes, vesicles, and micelles [6](#page=6). #### 4.2.4 Sphingolipids Sphingolipids are another class of membrane lipids characterized by a sphingoid backbone, synthesized from serine and acetyl-CoA. They contain a fatty acid and a head group attached to the sphingoid base. Sphingomyelins (with a phosphate-linked choline or ethanolamine head group) are crucial in myelin sheaths, while cerebrosides and gangliosides (with carbohydrate head groups) are important in neural membranes [6](#page=6). #### 4.2.5 Sterol Lipids Sterol lipids, such as cholesterol, are predominantly found in eukaryotes and feature a rigid ring system. The ring structure is hydrophobic, while the hydroxyl (-OH) group is hydrophilic. Cholesterol is a precursor for steroid hormones like testosterone and progesterone and is vital for maintaining membrane stability and fluidity [6](#page=6). #### 4.2.6 Lipid Aggregation In aqueous environments, lipids spontaneously arrange into structures driven by the **hydrophobic effect**. Hydrophobic tails cluster inward to minimize contact with water, while hydrophilic heads face outward, interacting with water through hydrogen bonds and electrostatic forces. This self-assembly leads to the formation of micelles (from single fatty acids) or lipid bilayers (from phospholipids and sphingolipids), which form the basis of cell membranes [6](#page=6). The properties of lipid bilayers, such as permeability and fluidity, are influenced by the type of fatty acids present. Saturated fatty acids lead to tighter packing and higher melting temperatures ($T_m$), while unsaturated fatty acids, with their kinks, disrupt packing, lowering $T_m$ and increasing fluidity [6](#page=6). ### 4.3 Nucleotides Nucleotides are the building blocks of nucleic acids (DNA and RNA). Each nucleotide consists of a ribose sugar, a phosphate group, and a nitrogenous base. In nucleic acid polymers, the phosphate group forms phosphodiester bonds with the 3' hydroxyl group of one sugar and the 5' carbon of the next, creating a sugar-phosphate backbone. The sequence of nucleotides is read from the 5' to 3' direction [7](#page=7). DNA typically exists as a stable double helix, with base pairing (A with T, and G with C) stabilized by hydrogen bonds (2 between A-T, 3 between G-C). The structure has a thickness of 2 nm, with a helical repeat every 3.4 nm and bases spaced 0.34 nm apart. The number of G-C bonds influences DNA stability, as they require more thermal energy to separate [7](#page=7). Nucleotides also function as monomeric units in energy metabolism (ATP, GTP), as second messengers (cAMP), and as substrates for enzymes [7](#page=7). ### 4.4 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, the **primary structure**, dictates the protein's overall structure and function [7](#page=7). * **Secondary structure** involves the folding of the polypeptide chain into regular structures like $\alpha$-helices and $\beta$-sheets, stabilized by hydrogen bonds between backbone atoms [7](#page=7). * **Tertiary structure** refers to the overall three-dimensional conformation 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 [7](#page=7). * **Quaternary structure** describes the arrangement of multiple polypeptide subunits in a functional protein complex [7](#page=7). Protein folding is a complex process driven by thermodynamics, often assisted by chaperone proteins. Techniques like chromatography (size exclusion, ion exchange, affinity) and SDS-PAGE are used for protein purification and analysis [7](#page=7). ### 4.5 Enzymes Enzymes are biological catalysts, typically proteins, that accelerate biochemical reactions by lowering the activation energy ($E_a$) without being consumed in the process. They achieve this by providing an alternative reaction pathway and stabilizing the transition state [7](#page=7). * Enzyme activity is influenced by factors such as pH and temperature, with each enzyme having an optimal pH and temperature range. High temperatures can denature enzymes, disrupting their three-dimensional structure and thus their function [7](#page=7). * Enzymes bind specifically to their substrates at the **active site**, forming an enzyme-substrate complex (ES) [7](#page=7). * Enzyme kinetics, described by Michaelis-Menten kinetics, relates reaction rate ($v$) to substrate concentration ($[S]$), characterized by $V_{max}$ (maximum reaction velocity) and $K_m$ (substrate concentration at $v = V_{max}/2$), which indicates substrate affinity. The $k_{cat}$ (turnover number) measures catalytic efficiency [7](#page=7). * Enzyme activity can be regulated by **inhibitors**, which bind to enzymes and reduce their catalytic rate. Inhibitors can be competitive (binding to the active site), uncompetitive (binding to the ES complex), or non-competitive (binding to an allosteric site) [7](#page=7). * **Cofactors** (e.g., coenzymes) are non-protein molecules that often assist enzymes in their catalytic functions [7](#page=7). * **Allosteric regulation** involves the binding of molecules to sites other than the active site, altering enzyme conformation and activity. Many enzymes, especially those in complex metabolic pathways, exhibit sigmoidal kinetics due to allosteric regulation of multisubunit proteins [7](#page=7). * Reversible covalent modifications, such as phosphorylation, are also key mechanisms for enzyme regulation, particularly in signal transduction pathways [7](#page=7). ### 4.6 Summary of Biomolecule Interactions and Properties The structure of biomolecules dictates their function through various types of bonding and interactions: * **Covalent bonds** are formed by the sharing of electrons and are central to the structure of biomolecules like carbohydrates, lipids, nucleotides, and proteins [4](#page=4). * **Non-covalent interactions**, including hydrogen bonds, ionic interactions, London dispersion forces, and van der Waals forces, are crucial for stabilizing the three-dimensional structures of biomolecules and mediating their interactions with each other. Hydrogen bonds, specifically, are important in DNA structure and protein secondary structures [5](#page=5) [7](#page=7). * **Electronegativity** differences between atoms determine bond polarity and can lead to the formation of **dipoles** within molecules. The overall polarity of a molecule, influenced by its geometry, dictates its solubility and interactions (e.g., H$_2$O is polar, CO$_2$ is nonpolar despite polar bonds) [4](#page=4) [5](#page=5). * **Chirality** and **stereoisomerism** are fundamental properties of many biomolecules, especially carbohydrates and amino acids, leading to different biological activities for different stereoisomers (e.g., L-amino acids are found in proteins, D-glucose is a common sugar). Techniques like Fischer and Haworth projections are used to depict these spatial arrangements [4](#page=4) [6](#page=6) [7](#page=7). * The arrangement of electrons in **atomic orbitals** and their hybridization (e.g., sp$^3$, sp$^2$) explain molecular geometry and bonding capabilities, such as in the tetrahedral structure of methane [4](#page=4). --- # Enzyme kinetics and regulation This section delves into the principles governing the rates of enzyme-catalyzed reactions and the mechanisms by which these rates are controlled. ### 5.1 Enzyme kinetics Enzyme kinetics studies the rates of enzyme-catalyzed reactions, which are influenced by factors such as substrate concentration, enzyme concentration, temperature, and pH. #### 5.1.1 Michaelis-Menten kinetics Michaelis-Menten kinetics describes the relationship between the initial reaction velocity ($V_0$) and substrate concentration ($[S]$) for enzyme-catalyzed reactions. It is based on the formation of an enzyme-substrate complex (ES) followed by its breakdown into product (P) and free enzyme (E). The fundamental steps are: 1. $E + S \rightleftharpoons ES$ (formation of enzyme-substrate complex) 2. $ES \rightarrow E + P$ (breakdown to product) The Michaelis-Menten equation is: $$V_0 = \frac{V_{max}[S]}{K_m + [S]}$$ Where: * $V_0$ is the initial reaction velocity. * $V_{max}$ is the maximum reaction velocity, achieved when the enzyme is saturated with substrate. * $[S]$ is the substrate concentration. * $K_m$ (Michaelis constant) is the substrate concentration at which the reaction velocity is half of $V_{max}$. It is a measure of the enzyme's affinity for its substrate; a lower $K_m$ indicates higher affinity. **K_cat (Turnover Number):** This represents the maximum number of substrate molecules an enzyme can convert into product per unit time when saturated with substrate. It is a measure of the enzyme's catalytic efficiency. $K_{cat} = \frac{V_{max}}{[E]_T}$, where $[E]_T$ is the total enzyme concentration. **Catalytic Efficiency:** This is often expressed as the ratio $k_{cat}/K_m$. A higher value indicates a more efficient enzyme. **Lineweaver-Burk Plot (Double Reciprocal Plot):** This is a linear transformation of the Michaelis-Menten equation, plotting $1/V_0$ against $1/[S]$. It is useful for determining $V_{max}$ and $K_m$ graphically. $$\frac{1}{V_0} = \frac{K_m}{V_{max}[S]} + \frac{1}{V_{max}}$$ The plot yields a straight line with a y-intercept of $1/V_{max}$ and an x-intercept of $-1/K_m$. #### 5.1.2 Enzyme inhibition Enzyme inhibitors are substances that bind to enzymes and reduce their activity. * **Reversible Inhibitors:** These bind non-covalently and their effects can be reversed. * **Competitive Inhibitors:** * Bind to the active site, competing with the substrate. * Increase $K_m$ (higher substrate concentration is needed to reach half $V_{max}$), but $V_{max}$ remains unchanged. * The effect can be overcome by increasing substrate concentration. * **Uncompetitive Inhibitors:** * Bind only to the enzyme-substrate (ES) complex, preventing product release. * Decrease both $K_m$ and $V_{max}$ by the same factor. * Cannot be overcome by increasing substrate concentration. * **Non-competitive Inhibitors:** * Bind to a site distinct from the active site (allosteric site) and can bind to both free enzyme and the ES complex. * Do not affect substrate binding ( $K_m$ remains unchanged), but they reduce the enzyme's catalytic efficiency, thus decreasing $V_{max}$. * Effectively reduce the concentration of active enzyme. * **Mixed Inhibitors:** Bind to both the free enzyme and the ES complex, but with different affinities. They affect both $K_m$ and $V_{max}$. * **Irreversible Inhibitors:** These bind covalently to the enzyme, permanently inactivating it. ### 5.2 Enzyme regulation Enzymes are regulated to control metabolic pathways and cellular processes. #### 5.2.1 Allosteric regulation Allosteric regulation involves the binding of effector molecules (activators or inhibitors) to an allosteric site on the enzyme, which is different from the active site. This binding induces a conformational change in the enzyme, altering its catalytic activity. * **Allosteric activators** increase enzyme activity. * **Allosteric inhibitors** decrease enzyme activity. * Allosteric enzymes, particularly those with multiple subunits, often exhibit sigmoidal kinetics rather than the hyperbolic Michaelis-Menten kinetics, showing a cooperative binding of substrate. #### 5.2.2 Reversible covalent modifications Enzyme activity can be modulated by the reversible addition or removal of small chemical groups, such as phosphate (phosphorylation), acetyl groups (acetylation), or methyl groups (methylation). Phosphorylation, catalyzed by kinases, is a common regulatory mechanism involved in signal transduction pathways, often activating or inactivating enzymes. #### 5.2.3 Feedback inhibition In feedback inhibition, the end product of a metabolic pathway acts as an allosteric inhibitor for an enzyme earlier in the pathway, preventing the overproduction of the product. #### 5.2.4 Enzyme levels regulation The amount of an enzyme in a cell can also be regulated at the transcriptional or translational level, influencing the overall rate of a metabolic pathway (translational regulation). ### 5.3 Factors affecting enzyme activity #### 5.3.1 pH Enzymes have an optimal pH at which their activity is maximal. Deviations from the optimal pH can alter the ionization state of amino acid residues in the active site or elsewhere in the enzyme, affecting substrate binding and catalysis, or even leading to denaturation. #### 5.3.2 Temperature Enzymes have an optimal temperature for activity. At temperatures below the optimum, reaction rates increase with temperature due to increased kinetic energy. However, at temperatures above the optimum, enzyme activity decreases sharply as the enzyme denatures (loses its three-dimensional structure) due to the disruption of non-covalent bonds. --- # Metabolic pathways: glycolysis, Krebs cycle, and oxidative phosphorylation This section details central metabolic pathways, focusing on glycolysis, the Krebs cycle, and oxidative phosphorylation, explaining their steps, energy production, and regulation. ### 6.1 Glycolysis Glycolysis is a fundamental metabolic pathway that breaks down glucose into pyruvate. It occurs in the cytosol and does not require oxygen, making it an anaerobic process. The overall reaction is [6](#page=6) [7](#page=7): Glucose + 2 NAD$^+$ + 2 ADP + 2 P$_\text{i}$ $\rightarrow$ 2 Pyruvate + 2 NADH + 2 H$^+$ + 2 ATP + 2 H$_2$O [6](#page=6). Glycolysis can be divided into three main stages: #### 6.1.1 Investment Phase This phase requires an input of energy to prepare glucose for cleavage. It involves three reactions that convert glucose into fructose-1,6-bisphosphate, consuming two ATP molecules. * **Reaction 1: Phosphorylation of glucose** * Glucose is phosphorylated to glucose-6-phosphate by hexokinase. * This reaction is irreversible and traps glucose within the cell, as the addition of a phosphate group makes it negatively charged and unable to cross the cell membrane via glucose transporters [7](#page=7). * Reaction: Glucose + ATP $\rightarrow$ Glucose-6-phosphate + ADP * **Reaction 2: Isomerization** * Glucose-6-phosphate (an aldose) is isomerized to fructose-6-phosphate (a ketose) by phosphoglucose isomerase [7](#page=7). * This rearrangement is necessary because the subsequent cleavage requires the carbonyl group to be at carbon 2, which is characteristic of a ketose. The ring structure of fructose-6-phosphate also provides a free hydroxyl group at C1, facilitating further phosphorylation [7](#page=7). * Reaction: Glucose-6-phosphate $\rightleftharpoons$ Fructose-6-phosphate * **Reaction 3: Second Phosphorylation** * Fructose-6-phosphate is phosphorylated to fructose-1,6-bisphosphate by phosphofructokinase. * This is another irreversible, energy-consuming step, requiring ATP [7](#page=7). * Reaction: Fructose-6-phosphate + ATP $\rightarrow$ Fructose-1,6-bisphosphate + ADP #### 6.1.2 Lysis Phase This stage splits the six-carbon fructose-1,6-bisphosphate molecule into two three-carbon molecules. * **Cleavage:** Fructose-1,6-bisphosphate is cleaved into dihydroxyacetone phosphate and glyceraldehyde-3-phosphate by aldolase. This reaction is reversible [7](#page=7). * **Isomerization:** Dihydroxyacetone phosphate is converted into glyceraldehyde-3-phosphate by triosephosphate isomerase. This is also a reversible reaction, and the equilibrium favors glyceraldehyde-3-phosphate due to its consumption in the subsequent steps of glycolysis [7](#page=7). * Following these steps, two molecules of glyceraldehyde-3-phosphate proceed through the rest of glycolysis. #### 6.1.3 Payoff Phase In this phase, the three-carbon units are converted into pyruvate, generating ATP and NADH. This phase occurs twice for each initial glucose molecule. * **Reaction 1: Oxidation and Phosphorylation** * Glyceraldehyde-3-phosphate is oxidized and phosphorylated to 1,3-bisphosphoglycerate by glyceraldehyde-3-phosphate dehydrogenase. This reaction involves the reduction of NAD$^+$ to NADH. The enzyme couples a highly exergonic reaction with an endergonic reaction, utilizing a high-energy intermediate [7](#page=7). * Reaction: Glyceraldehyde-3-phosphate + NAD$^+$ + P$_\text{i}$ $\rightleftharpoons$ 1,3-Bisphosphoglycerate + NADH + H$^+$ * **Reaction 2: ATP Production (Substrate-Level Phosphorylation)** * 1,3-Bisphosphoglycerate is converted to 3-phosphoglycerate, with the transfer of a phosphate group to ADP to form ATP. This is an example of substrate-level phosphorylation [7](#page=7). * Reaction: 1,3-Bisphosphoglycerate + ADP $\rightleftharpoons$ 3-Phosphoglycerate + ATP * **Reaction 3: Phosphate Group Rearrangement** * 3-Phosphoglycerate is isomerized to 2-phosphoglycerate by phosphoglycerate mutase, moving the phosphate group from carbon 3 to carbon 2 [7](#page=7). * Reaction: 3-Phosphoglycerate $\rightleftharpoons$ 2-Phosphoglycerate * **Reaction 4: Dehydration** * 2-Phosphoglycerate is dehydrated to phosphoenolpyruvate (PEP) by enolase, removing a water molecule [7](#page=7). * Reaction: 2-Phosphoglycerate $\rightleftharpoons$ Phosphoenolpyruvate + H$_2$O * **Reaction 5: Second ATP Production (Substrate-Level Phosphorylation)** * Phosphoenolpyruvate is converted to pyruvate by pyruvate kinase, producing ATP through substrate-level phosphorylation. This reaction is irreversible [7](#page=7). * Reaction: Phosphoenolpyruvate + ADP $\rightarrow$ Pyruvate + ATP **Net yield of glycolysis per glucose molecule:** 2 ATP, 2 NADH, and 2 pyruvate molecules. * **Anaerobic conditions:** If oxygen is absent, pyruvate is converted into lactate (in humans) or ethanol (in yeast) through fermentation to regenerate NAD$^+$ [7](#page=7). * **Aerobic conditions:** Pyruvate is transported into the mitochondrial matrix and converted to acetyl-CoA, which then enters the Krebs cycle. ### 6.2 Krebs Cycle (Citric Acid Cycle or TCA Cycle) The Krebs cycle is a series of reactions occurring in the mitochondrial matrix that oxidizes acetyl-CoA, producing ATP, NADH, FADH$_2$, and carbon dioxide. **Overview of the Krebs Cycle:** The cycle begins with the condensation of acetyl-CoA (a two-carbon molecule) with oxaloacetate (a four-carbon molecule) to form citrate (a six-carbon molecule). Through a series of enzymatic steps, citrate is oxidized and rearranged, releasing carbon dioxide and generating high-energy electron carriers (NADH and FADH$_2$) and ATP (or GTP). The cycle regenerates oxaloacetate, allowing it to accept another acetyl-CoA molecule. **Key steps and outputs:** 1. **Citrate formation:** Acetyl-CoA + Oxaloacetate + H$_2$O $\rightarrow$ Citrate + CoA [7](#page=7). * Catalyzed by citrate synthase. 2. **Isomerization:** Citrate $\rightarrow$ Isocitrate [7](#page=7). * Catalyzed by aconitase. This step repositions the hydroxyl group to prepare for decarboxylation. 3. **Oxidative decarboxylation:** Isocitrate + NAD$^+$ $\rightarrow$ $\alpha$-ketoglutarate + NADH + H$^+$ + CO$_2$ [7](#page=7). * Catalyzed by isocitrate dehydrogenase. This is the first of four oxidative decarboxylations in the cycle. 4. **Oxidative decarboxylation:** $\alpha$-ketoglutarate + NAD$^+$ + CoA $\rightarrow$ Succinyl-CoA + NADH + H$^+$ + CO$_2$ [7](#page=7). * Catalyzed by the $\alpha$-ketoglutarate dehydrogenase complex. This reaction forms a high-energy thioester bond in succinyl-CoA. 5. **Substrate-level phosphorylation:** Succinyl-CoA + ADP + P$_\text{i}$ $\rightarrow$ Succinate + ATP + CoA [7](#page=7). * Catalyzed by succinyl-CoA synthase. The energy from the thioester bond hydrolysis drives ATP synthesis. GTP can be formed instead of ATP depending on the enzyme isoform. 6. **Oxidation:** Succinate + FAD $\rightarrow$ Fumarate + FADH$_2$ [7](#page=7). * Catalyzed by succinate dehydrogenase, which is embedded in the inner mitochondrial membrane and is also part of Complex II of the electron transport chain. 7. **Hydration:** Fumarate + H$_2$O $\rightarrow$ L-Malate [7](#page=7). * Catalyzed by fumarase. This reaction creates a new stereocenter, specifically L-malate. 8. **Oxidation:** L-Malate + NAD$^+$ $\rightarrow$ Oxaloacetate + NADH + H$^+$ [7](#page=7). * Catalyzed by malate dehydrogenase. This reaction is energetically unfavorable but is driven forward by the consumption of oxaloacetate in the next cycle and NADH in the electron transport chain. **Net yield per acetyl-CoA molecule entering the cycle:** 3 NADH, 1 FADH$_2$, 1 ATP (or GTP), and 2 CO$_2$. Since one glucose molecule yields two pyruvate molecules, and thus two acetyl-CoA molecules, the Krebs cycle turns twice per glucose molecule. **Alternative metabolic substrates:** Glycerol and fatty acids can be broken down to intermediates that enter glycolysis or the Krebs cycle, respectively. Amino acids can also be catabolized to pyruvate, acetyl-CoA, or other Krebs cycle intermediates [7](#page=7). ### 6.3 Oxidative Phosphorylation Oxidative phosphorylation is the primary process for ATP production in aerobic respiration, occurring on the inner mitochondrial membrane. It involves two main components: the electron transport chain (ETC) and ATP synthase [7](#page=7). #### 6.3.1 Electron Transport Chain (ETC) The ETC is a series of protein complexes (Complex I-IV) embedded in the inner mitochondrial membrane that accept electrons from NADH and FADH$_2$ produced during glycolysis and the Krebs cycle [7](#page=7). * **Electron carriers:** NADH and FADH$_2$ donate electrons, which are passed down the chain through a series of redox reactions. * NADH donates electrons to Complex I [7](#page=7). * FADH$_2$ donates electrons to Complex II [7](#page=7). * **Proton pumping:** As electrons move through the complexes, energy is released and used to pump protons (H$^+$) from the mitochondrial matrix into the intermembrane space. This creates an electrochemical gradient, known as the proton-motive force [7](#page=7). * Complex I pumps 4 H$^+$ per NADH [7](#page=7). * Complex II does not pump protons directly but transfers electrons to ubiquinone [7](#page=7). * Complex III pumps 4 H$^+$ per pair of electrons from ubiquinol (QH$_2$) [7](#page=7). * Complex IV pumps 4 H$^+$ and uses oxygen as the final electron acceptor, reducing it to water [7](#page=7). * **Electron flow summary:** * NADH $\rightarrow$ Complex I $\rightarrow$ Ubiquinone $\rightarrow$ Complex III $\rightarrow$ Cytochrome c $\rightarrow$ Complex IV $\rightarrow$ O$_2$ [7](#page=7). * FADH$_2$ $\rightarrow$ Complex II $\rightarrow$ Ubiquinone $\rightarrow$ Complex III $\rightarrow$ Cytochrome c $\rightarrow$ Complex IV $\rightarrow$ O$_2$ [7](#page=7). #### 6.3.2 ATP Synthase ATP synthase is a molecular machine that utilizes the proton-motive force to synthesize ATP. * **Structure:** It consists of two main parts: F$_0$ (embedded in the membrane) and F$_1$ (catalytic headpiece extending into the matrix) [7](#page=7). * **Mechanism:** Protons flow from the intermembrane space back into the matrix through a channel in the F$_0$ component. This proton flow drives the rotation of a rotor within ATP synthase, causing conformational changes in the F$_1$ headpiece. These changes facilitate the binding of ADP and inorganic phosphate (P$_\text{i}$) and their subsequent phosphorylation to ATP [7](#page=7). * **ATP yield:** The number of ATP molecules produced per NADH and FADH$_2$ molecule varies. Generally, one NADH molecule yields approximately 2.5 ATP, and one FADH$_2$ molecule yields approximately 1.5 ATP. NADH produced in the cytosol during glycolysis requires shuttle systems (malate-aspartate or glycerol phosphate shuttle) to transfer its electrons into the mitochondria, which can affect the final ATP yield [7](#page=7). **Overall ATP production:** This process generates the vast majority of ATP during aerobic respiration. Without oxygen, the ETC cannot function, and ATP production ceases, leading to cell death. **Regulation and Uncoupling:** * Uncouplers, like FCCP, can dissipate the proton gradient by facilitating proton re-entry into the matrix, uncoupling electron transport from ATP synthesis. This increases oxygen consumption but reduces ATP production [7](#page=7). * The Warburg effect describes increased glycolysis even in the presence of oxygen in cancer cells, potentially providing rapid ATP and building blocks for proliferation [7](#page=7). --- ## 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 orbiting electrons. | | Nucleus | The central part of an atom, composed of protons and neutrons, which carries a positive charge due to the protons. | | Proton | A subatomic particle found in the nucleus of an atom, carrying a positive charge (+1) and having a mass comparable to a neutron. | | Neutron | A subatomic particle found in the nucleus of an atom, carrying no charge (0), and having a mass nearly equal to a proton. | | Electron | A subatomic particle with a negative charge (-1) that orbits the nucleus of an atom. Its mass is significantly smaller than that of a proton or neutron. | | Atomic number (Z) | The number of protons in the nucleus of an atom, which uniquely identifies an element and determines the number of electrons in a neutral atom. | | Atomic mass number (A) | The total number of protons and neutrons in an atom's nucleus. | | Unified atomic mass unit (Da or U) | A standard unit of mass used for atoms and molecules, defined as 1/12 of the mass of a neutral atom of carbon-12. It is approximately equal to the mass of a proton or neutron. | | Periodic table | An arrangement of chemical elements ordered by their atomic number, electron configuration, and recurring chemical properties, organized into groups (columns) and periods (rows). | | Group (in periodic table) | A vertical column in the periodic table, where elements typically share the same number of valence electrons and exhibit similar chemical properties. | | Period (in periodic table) | A horizontal row in the periodic table, where elements indicate the number of electron shells an atom possesses. | | Ionization energy | The minimum energy required to remove an electron from a neutral atom or molecule in its gaseous state. | | Electron affinity | The energy change that occurs when an electron is added to a neutral atom or molecule to form a negative ion. | | Atomic radius | The distance from the center of an atom's nucleus to its outermost electron shell. | | Electronegativity | A measure of the tendency of an atom to attract a bonding pair of electrons towards itself. | | Isotope | Atoms of the same element (same atomic number) that have different numbers of neutrons, resulting in different atomic mass numbers and consequently different atomic masses. | | Bohr model | An early model of the atom describing electrons orbiting the nucleus in specific energy levels or orbits. It was largely superseded by the quantum mechanical model. | | Quantum mechanical model | A model of the atom that describes electrons as having wave-like properties and existing in probability distributions called atomic orbitals, rather than fixed orbits. | | Schrödinger wave equation | A fundamental equation in quantum mechanics that describes the wave-like behavior of subatomic particles, including electrons. | | Wave function | A mathematical function ($\psi$) that describes the quantum state of an atomic or subatomic system. The square of the wave function gives the probability density of finding a particle in a particular location. | | Atomic orbital | A region in space around an atom's nucleus where there is a high probability (typically 90%) of finding an electron. Orbitals have distinct shapes, sizes, and orientations. | | Quantum numbers | A set of numbers that describe the properties of electrons in atoms, including their energy, angular momentum, orientation in space, and spin. The four quantum numbers are: principal (n), orbital (l), magnetic (m_L), and spin (m_S). | | Principal quantum number (n) | Describes the energy level and size of an atomic orbital. Higher values of n indicate higher energy and greater distance from the nucleus. It can be any positive integer (1, 2, 3,...). | | Orbital quantum number (l) | Describes the shape of an atomic orbital and the subshell it belongs to. For a given n, l can take values from 0 to n-1. l=0 corresponds to s orbitals (spherical), l=1 to p orbitals (dumbbell-shaped), l=2 to d orbitals, and l=3 to f orbitals. | | Magnetic quantum number (m_L) | Describes the orientation of an atomic orbital in space relative to an external magnetic field. For a given l, m_L can take integer values from -l to +l, including 0. | | Spin quantum number (m_S) | Describes the intrinsic angular momentum of an electron, often visualized as its spin. It can only take two values: +1/2 or -1/2. | | Pauli exclusion principle | A principle stating that no two electrons in an atom can have the same set of four quantum numbers. This means that an atomic orbital can hold a maximum of two electrons, and they must have opposite spins. | | Degenerate atomic orbitals | Atomic orbitals within the same subshell that have the same energy but different spatial orientations (e.g., the three p orbitals: px, py, pz). | | Aufbau principle | A principle that states that electrons fill atomic orbitals starting from the lowest energy levels and proceeding to higher ones. | | Hund's rule | A rule stating that within a subshell, electrons will occupy each orbital singly before pairing up, and these singly occupied orbitals will have parallel spins. | | Molecule | An electrically neutral entity formed when two or more atoms are chemically bonded together. | | Ionic bond | A chemical bond formed by the electrostatic attraction between oppositely charged ions, typically formed between a metal and a nonmetal due to a large difference in electronegativity. | | Cation | A positively charged ion, formed when an atom loses one or more electrons. | | Anion | A negatively charged ion, formed when an atom gains one or more electrons. | | Covalent bond | A chemical bond formed by the sharing of one or more pairs of electrons between atoms, typically occurring between nonmetals with similar electronegativities. | | Polar covalent bond | A covalent bond in which the electron pair is shared unequally between two atoms due to a difference in electronegativity, creating partial positive and negative charges on the atoms. | | Molecular formula | A representation of a molecule that lists the types and numbers of atoms present using chemical symbols. | | Empirical formula | The simplest whole-number ratio of atoms in a compound. | | Hill notation | A system for writing chemical formulas, particularly for organic compounds, where carbon atoms are listed first, followed by hydrogen, and then other elements in alphabetical order. Exceptions exist for ionic compounds, oxides, acids, and hydroxides. | | Structural formula | A representation of a molecule that shows the order in which atoms are connected and the arrangement of chemical groups. | | Skeletal formula | A simplified representation of organic molecules where carbon atoms and hydrogen atoms bonded to carbon are implied, and lines represent bonds. Heteroatoms are explicitly shown. | | Valence shell electron pair repulsion (VSEPR) theory | A theory used to predict the geometry of molecules based on the repulsion between electron pairs (both bonding and lone pairs) in the valence shell of the central atom. Electron groups arrange themselves to be as far apart as possible. | | Sigma bond ($\sigma$) | A type of covalent bond formed by the direct, end-to-end overlap of atomic orbitals along the internuclear axis. | | Pi bond ($\pi$) | A type of covalent bond formed by the sideways overlap of atomic orbitals (typically p orbitals) above and below the internuclear axis. Pi bonds are generally weaker than sigma bonds and restrict rotation around the bond axis. | | Hybridization | The mixing of atomic orbitals within an atom to form new hybrid orbitals that are better suited for bonding, having different shapes, energies, and orientations than the original atomic orbitals (e.g., sp, sp², sp³). | | sp³ hybridization | The hybridization of one s orbital and three p orbitals to form four equivalent sp³ hybrid orbitals, which are arranged in a tetrahedral geometry. | | sp² hybridization | The hybridization of one s orbital and two p orbitals to form three equivalent sp² hybrid orbitals, arranged in a trigonal planar geometry, leaving one unhybridized p orbital. | | sp hybridization | The hybridization of one s orbital and one p orbital to form two equivalent sp hybrid orbitals, arranged in a linear geometry, leaving two unhybridized p orbitals. | | Molecular orbital theory | A theory that describes chemical bonding in terms of molecular orbitals, which are formed by the combination of atomic orbitals from different atoms. | | Fajans' rule | A set of guidelines used to predict the degree of covalent character in an ionic bond, based on factors such as the charge density of the cation and the polarizability of the anion. | | Moles (mol) | The SI unit for the amount of substance, defined as containing exactly $6.02214076 \times 10^{23}$ elementary entities (Avogadro's number). | | Avogadro's constant ($N_A$ or L) | The number of constituent particles (atoms, molecules, ions, etc.) that are contained in the amount of substance given by one mole. Its value is approximately $6.022 \times 10^{23}$ mol$^{-1}$. | | Molar mass | The mass of one mole of a substance, typically expressed in grams per mole (g mol$^{-1}$). It is numerically equivalent to the atomic or molecular weight in daltons. | | Concentration | The amount of a solute dissolved in a specific amount of solvent or solution. Common units include molarity (mol/L), mass concentration (g/L), and volume concentration (%). | | Molar concentration (Molarity) | The concentration of a solution expressed as the number of moles of solute per liter (or cubic decimeter) of solution (mol dm$^{-3}$ or M). | | Mass concentration | The concentration of a solution expressed as the mass of solute per unit volume of solution (e.g., g/L). | | Volume concentration | The concentration of a solution expressed as the volume of solute per unit volume of solution, often as a percentage. | | Dilution | The process of reducing the concentration of a solute in a solution by adding more solvent. | | Dilution factor (DF) | A ratio representing how many times a solution has been diluted, often expressed as the ratio of the stock concentration to the final concentration or the final volume to the stock volume. | | Dilution ratio | A ratio comparing the volume of the sample to the volume of the solvent added, e.g., 1:9 means 1 part sample to 9 parts solvent. | | Acid | A substance that can donate a proton (H$^+$) in solution (Brønsted-Lowry definition). | | Base | A substance that can accept a proton (H$^+$) in solution (Brønsted-Lowry definition). | | Water dissociation | The process by which water molecules can split into hydrogen ions (H$^+$) and hydroxide ions (OH$^-$), or more accurately, form hydronium ions (H$_3$O$^+$) and hydroxide ions (OH$^-$). | | Acid dissociation constant (K$_a$) | An equilibrium constant that quantifies the strength of an acid in solution; it represents the ratio of dissociated ions to the undissociated acid at equilibrium. A higher K$_a$ indicates a stronger acid. | | Ionic product of water (K$_w$) | The product of the molar concentrations of hydrogen ions ([H$^+$]) and hydroxide ions ([OH$^-$]) in pure water at a given temperature. At 25$^\circ$C, K$_w$ is approximately $1 \times 10^{-14}$ M$^2$. | | pH | A measure of the acidity or alkalinity of a solution, defined as the negative logarithm (base 10) of the hydrogen ion concentration: pH = -log$_{10}$[H$^+$]. A pH less than 7 is acidic, greater than 7 is basic, and equal to 7 is neutral. | | pKa | The negative logarithm (base 10) of the acid dissociation constant (K$_a$): pKa = -log$_{10}$(K$_a$). It is a measure of acid strength; a lower pKa indicates a stronger acid. | | pH buffer | An aqueous solution that resists significant changes in pH upon the addition of small amounts of acid or base. It typically consists of a weak acid and its conjugate base, or a weak base and its conjugate acid. | | Strong acid | An acid that completely dissociates into its ions in aqueous solution. | | Weak acid | An acid that only partially dissociates in aqueous solution, establishing an equilibrium between the undissociated acid and its ions. | | Henderson-Hasselbalch equation | An equation used to calculate the pH of a buffer solution: pH = pKa + log$_{10}$([A$^-$]/[HA]), where [A$^-$] is the concentration of the conjugate base and [HA] is the concentration of the weak acid. | | 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. | | Structural isomerism | Isomers that have the same molecular formula but differ in the connectivity of their atoms (i.e., the order in which atoms are bonded). | | Stereoisomerism | Isomers that have the same molecular formula and the same connectivity of atoms but differ in the three-dimensional arrangement of atoms in space. | | Optical isomer (Enantiomer) | A pair of stereoisomers that are non-superimposable mirror images of each other. They have identical physical properties in achiral environments but differ in their interaction with plane-polarized light and with other chiral molecules. | | Stereocenter (Chiral center) | An atom (typically carbon) in a molecule bonded to four different atoms or groups, such that swapping any two groups creates a new stereoisomer. | | Diastereoisomers | Stereoisomers that are not mirror images of each other. They differ in the configuration at one or more, but not all, stereocenters. | | Epimer | A type of diastereoisomer that differs in the configuration at only one stereocenter. | | Anomer | A specific type of epimer formed when a sugar cyclizes into a ring structure. Anomers differ in configuration at the anomeric carbon. | | Electronegativity | The intrinsic property of an atom to attract electrons towards itself in a chemical bond. | | Dipole | A separation of positive and negative electric charges within a molecule, often due to differences in electronegativity and molecular geometry. | | Dipole moment | A measure of the magnitude and direction of a dipole within a molecule. | | Non-covalent interactions | Weak attractive or repulsive forces between molecules or between different parts of a large molecule that do not involve the sharing of electrons. Examples include hydrogen bonds, dipole-dipole interactions, London dispersion forces, and ionic interactions. | | Hydrogen bond | A strong type of dipole-dipole interaction that occurs between a hydrogen atom bonded to a highly electronegative atom (like O, N, or F) and another electronegative atom with a lone pair of electrons. | | London dispersion forces | Weak, temporary attractive forces that arise from transient fluctuations in electron distribution around atoms or molecules, inducing temporary dipoles that attract each other. They exist between all molecules. | | van der Waals forces | A general term for weak intermolecular forces, including London dispersion forces and dipole-dipole interactions. | | Thermodynamics | The study of energy, its transformations, and its relation to matter and macroscopic properties. It focuses on heat, work, temperature, and energy transfer. | | System | In thermodynamics, the part of the universe that is being studied, separated from the surroundings by a boundary. | | Surroundings | In thermodynamics, everything outside the system that can interact with it. | | Boundary | The surface separating the system from its surroundings. | | First law of thermodynamics | States that energy cannot be created or destroyed, only transferred or changed from one form to another. The total energy of an isolated system remains constant. | | Potential energy | Stored energy that an object possesses due to its position, state, or composition (e.g., chemical energy stored in bonds). | | Thermodynamic temperature | A measure of the average kinetic energy of the particles within a system, expressed in Kelvin (K). | | Heat | Energy transferred between a system and its surroundings due to a temperature difference. It flows spontaneously from a region of higher temperature to lower temperature. | | Enthalpy (H) | The total heat content of a chemical system, representing the sum of its internal energy and the product of its pressure and volume. The enthalpy change ($\Delta$H) at constant pressure is equal to the heat exchanged. | | Bond energy | The average amount of energy required to break one mole of a particular type of chemical bond in the gaseous state. It is also the energy released when that bond is formed. | | Enthalpy change ($\Delta$H) | The change in enthalpy during a process (e.g., a chemical reaction), calculated as the difference between the enthalpy of the products and the enthalpy of the reactants ($\Delta$H = H$_{products}$ - H$_{reactants}$). A negative $\Delta$H indicates an exothermic reaction (heat released), and a positive $\Delta$H indicates an endothermic reaction (heat absorbed). | | Hess's Law | States that the total enthalpy change for a chemical reaction is independent of the pathway taken, meaning it is the same whether the reaction occurs in one step or in multiple steps. | | Entropy (S) | A thermodynamic measure of the dispersal or randomness of energy within a system. It is often associated with the number of possible microstates a system can occupy. Entropy change ($\Delta$S) is measured in J mol$^{-1}$ K$^{-1}$. | | Second law of thermodynamics | States that the total entropy of an isolated system (or the universe) can only increase over time; spontaneous processes tend to move towards a state of greater disorder or energy dispersal. | | Gibbs free energy (G) | A thermodynamic potential that combines enthalpy and entropy to determine the spontaneity of a process at constant temperature and pressure. It is defined as G = H - TS. The change in Gibbs free energy ($\Delta$G) indicates spontaneity: $\Delta$G < 0 (spontaneous/exergonic), $\Delta$G > 0 (non-spontaneous/endergonic), $\Delta$G = 0 (at equilibrium). | | Chemical equilibrium | A state in a reversible reaction where the rate of the forward reaction equals the rate of the reverse reaction, resulting in no net change in the concentrations of reactants and products. | | Equilibrium constant (K) | A ratio of the product concentrations to reactant concentrations at equilibrium, raised to the power of their stoichiometric coefficients. It indicates the extent to which a reaction proceeds towards products. | | Van't Hoff Equation | Relates the change in the equilibrium constant (K) of a chemical reaction to the change in temperature (T) and the standard enthalpy change ($\Delta$H$^\circ$) and entropy change ($\Delta$S$^\circ$) of the reaction. | | Chemical kinetics | The study of the rates of chemical reactions and the factors that influence them, such as concentration, temperature, and catalysts. | | Transition state | A high-energy, unstable intermediate arrangement of atoms formed during a chemical reaction, where old bonds are breaking and new bonds are forming. It represents the peak of the activation energy barrier. | | Activation energy (E$_a$) | The minimum amount of energy required for reactant molecules to overcome the energy barrier and proceed to form products in a chemical reaction. | | Catalyst | A substance that increases the rate of a chemical reaction without being consumed in the process. Catalysts achieve this by providing an alternative reaction pathway with a lower activation energy. | | Electron pushing | A convention using curved arrows to depict the movement of electron pairs or single electrons during a chemical reaction, illustrating the mechanism. | | Nucleophile | An electron-rich species (atom, molecule, or ion) that readily donates an electron pair to form a chemical bond, typically with an electron-poor species (electrophile). | | Electrophile | An electron-poor species that readily accepts an electron pair to form a chemical bond, typically from an electron-rich species (nucleophile). | | Reaction mechanism | A step-by-step sequence of elementary reactions by which an overall chemical change occurs. | | Addition reaction | A reaction in which two or more molecules combine to form a larger molecule, often occurring at double or triple bonds. | | Elimination reaction | A reaction in which atoms or groups of atoms are removed from a molecule, often resulting in the formation of a double or triple bond. | | Substitution reaction | A reaction in which an atom or functional group in a molecule is replaced by another atom or functional group. | | Redox reaction | A reaction involving the transfer of electrons between chemical species, resulting in a change in their oxidation states. Oxidation is the loss of electrons, and reduction is the gain of electrons. | | SN1 reaction | A unimolecular nucleophilic substitution reaction that proceeds in two steps, typically involving the formation of a carbocation intermediate. | | SN2 reaction | A bimolecular nucleophilic substitution reaction that occurs in a single step, where the nucleophile attacks the substrate simultaneously as the leaving group departs. | | Biomolecules | Organic molecules that are essential for life, such as carbohydrates, lipids, proteins, and nucleic acids. | | Carbohydrates | Organic compounds composed of carbon, hydrogen, and oxygen, often with the general formula (CH$_2$O)$_n$. They serve as energy sources, structural components, and in cell recognition. | | Monosaccharide | The simplest form of carbohydrate, a single sugar unit that cannot be hydrolyzed into smaller carbohydrates (e.g., glucose, fructose). | | Aldose | A monosaccharide that contains an aldehyde functional group (R-CHO). | | Ketose | A monosaccharide that contains a ketone functional group (R-CO-R'). | | Fischer projection | A two-dimensional representation of a three-dimensional molecule, particularly useful for showing the stereochemistry around chiral centers. Horizontal lines represent bonds coming out of the plane, and vertical lines represent bonds going into the plane. | | Cyclization (of carbohydrates) | The process by which the open-chain form of a carbohydrate molecule reacts with itself to form a stable ring structure, typically involving the carbonyl group and a hydroxyl group. | | Hemiacetal/Hemiketal | The functional group formed when an aldehyde or ketone reacts with an alcohol, resulting in a molecule with both an ether and a hydroxyl group attached to the same carbon atom. | | Anomeric carbon | The carbon atom in a cyclic monosaccharide that was originally the carbonyl carbon (aldehyde or ketone) in the open-chain form. This carbon becomes a new stereocenter upon cyclization. | | Haworth projection | A way to represent the cyclic structure of carbohydrates, showing the ring in perspective and indicating the stereochemistry of substituents. | | Disaccharide | A carbohydrate formed by the covalent linkage of two monosaccharide units through a glycosidic bond. | | Glycosidic bond/linkage | A type of covalent bond that links monosaccharide units together to form disaccharides and polysaccharides, or links a sugar to another molecule. | | Polysaccharide | A complex carbohydrate composed of long chains of monosaccharide units linked by glycosidic bonds (e.g., starch, cellulose, glycogen). | | Lipids | A diverse group of hydrophobic molecules that are insoluble in water, including fatty acids, triglycerides, phospholipids, and steroids. They function in energy storage, cell membranes, and signaling. | | Fatty acid | A long hydrocarbon chain with a carboxylic acid group (-COOH) at one end. They are the building blocks of many lipids. | | Saturated fatty acid | A fatty acid that contains only single bonds between carbon atoms in its hydrocarbon chain. | | Unsaturated fatty acid | A fatty acid that contains one or more double bonds between carbon atoms in its hydrocarbon chain. | | Triacylglycerol | A type of lipid formed from glycerol esterified with three fatty acids. It serves as a major form of energy storage in adipose tissue. | | Ester bond | A covalent bond formed between a carboxylic acid group and a hydroxyl group, with the release of a water molecule. | | Glycerophospholipid | A type of phospholipid in which glycerol is esterified to two fatty acids and a phosphate group. The phosphate group is typically linked to another polar molecule, forming the hydrophilic head. | | Phospholipid | A lipid molecule that contains a phosphate group, forming a key component of cell membranes. They are amphiphilic, with a hydrophilic head and hydrophobic tails. | | Amphiphilic molecule | A molecule that contains both hydrophilic (water-attracting) and hydrophobic (water-repelling) regions. | | Sphingolipid | A type of lipid that contains a sphingoid backbone, often found in cell membranes, particularly in the nervous system. | | Sterol lipid | Lipids characterized by a steroid nucleus, a four-ring structure. Cholesterol is a prominent example, important in cell membranes and as a precursor for steroid hormones. | | Hydrophobic effect | The tendency of nonpolar molecules or parts of molecules to aggregate in aqueous solution to minimize their contact with water, driven by the increase in entropy of the water molecules. | | Micelle | A spherical aggregate of amphiphilic molecules in aqueous solution, with hydrophobic tails oriented inward and hydrophilic heads facing outward. | | Vesicle/Lipid bilayer | A bilayer of amphiphilic molecules (like phospholipids) arranged with hydrophobic tails inward and hydrophilic heads outward, forming a stable structure that encloses an aqueous compartment. | | Nucleotide | The basic building block of nucleic acids (DNA and RNA), consisting of a nucleoside (a sugar linked to a nitrogenous base) and one or more phosphate groups. | | Nucleic acids | Polymers of nucleotides, deoxyribonucleic acid (DNA) and ribonucleic acid (RNA), which carry genetic information. | | DNA double helix | The characteristic coiled structure of DNA, formed by two antiparallel polynucleotide strands held together by hydrogen bonds between complementary bases (A-T, G-C). | | Amino acid | An organic molecule containing both an amino group (-NH$_2$) and a carboxyl group (-COOH). They are the building blocks of proteins. | | Peptide bond | A covalent bond formed between the carboxyl group of one amino acid and the amino group of another, releasing a molecule of water. This bond links amino acids together in a polypeptide chain. | | Primary structure (of protein) | The linear sequence of amino acids in a polypeptide chain. | | Secondary structure (of protein) | Localized folding patterns of the polypeptide chain, stabilized by hydrogen bonds between backbone atoms, forming structures like alpha-helices and beta-sheets. | | Tertiary structure (of protein) | The overall three-dimensional shape of a single polypeptide chain, determined by interactions between amino acid side chains, including hydrogen bonds, ionic interactions, hydrophobic interactions, and disulfide bridges. | | Quaternary structure (of protein) | The arrangement of multiple polypeptide subunits in a functional protein complex. | | Chaperones | Proteins that assist in the proper folding of other proteins, preventing misfolding and aggregation. | | Bioanalysis | Techniques used to study the structure and function of biological molecules using biochemical and biophysical methods. | | Protein purification | Processes used to isolate a specific protein from a complex mixture, often involving chromatography and electrophoresis. | | Chromatography | A separation technique used to separate components of a mixture based on their differential partitioning between a stationary phase and a mobile phase. Common types include size exclusion, ion exchange, and affinity chromatography. | | SDS-PAGE (Sodium Dodecyl Sulfate-Polyacrylamide Gel Electrophoresis) | A technique used to separate proteins based primarily on their molecular weight (size) by denaturing them with SDS and applying an electric field through a polyacrylamide gel. | | Enzyme | A biological catalyst, typically a protein, that accelerates the rate of biochemical reactions by lowering the activation energy without being consumed in the process. | | Michaelis-Menten kinetics | A model describing the rate of enzyme-catalyzed reactions as a function of substrate concentration. It relates the reaction velocity (v) to substrate concentration ([S]) through the Michaelis constant (K$_m$) and maximum velocity (V$_{max}$). | | K$_m$ (Michaelis constant) | The substrate concentration at which the reaction rate is half of its maximum velocity (V$_{max}$). It is an indicator of the enzyme's affinity for its substrate; a lower K$_m$ suggests higher affinity. | | V$_{max}$ (Maximum velocity) | The maximum rate of an enzyme-catalyzed reaction, achieved when the enzyme is saturated with substrate. | | k$_{cat}$ (Turnover number) | The catalytic rate constant of an enzyme, representing the number of substrate molecules converted to product per enzyme molecule per unit time when the enzyme is saturated with substrate. | | Catalytic efficiency | A measure of an enzyme's effectiveness, often expressed as the ratio k$_{cat}$/K$_m$. A higher value indicates a more efficient enzyme. | | Enzyme inhibitors | Molecules that bind to enzymes and reduce their activity, often by blocking substrate binding or altering enzyme conformation. | | Competitive inhibitor | An inhibitor that binds to the active site of an enzyme, competing with the substrate. It increases K$_m$ but does not affect V$_{max}$. | | Uncompetitive inhibitor | An inhibitor that binds only to the enzyme-substrate complex, reducing V$_{max}$ and decreasing K$_m$. | | Non-competitive inhibitor | An inhibitor that binds to an allosteric site on the enzyme (not the active site), affecting enzyme activity but not substrate binding. It reduces V$_{max}$ but does not affect K$_m$. | | Allosteric regulation | Regulation of enzyme activity by the binding of molecules (allosteric effectors) to a site other than the active site (the allosteric site), which alters the enzyme's conformation and activity. | | Cofactor | A non-protein chemical compound or metallic ion that is required for an enzyme's biological activity. | | Coenzyme | A type of cofactor that is an organic non-protein compound, often derived from vitamins, that binds to an enzyme and participates in its catalytic activity. | | Glycolysis | A metabolic pathway that breaks down glucose into two molecules of pyruvate, producing ATP and NADH in the process. It occurs in the cytoplasm and does not require oxygen. | | Gluconeogenesis (GNG) | The metabolic process by which glucose is synthesized from non-carbohydrate precursors (such as pyruvate, lactate, amino acids, and glycerol), primarily occurring in the liver. | | 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 later steps of glycolysis where ATP and NADH are produced as glucose is converted into pyruvate. | | Substrate-level phosphorylation | The direct transfer of a phosphate group from a high-energy substrate molecule to ADP, forming ATP, without the involvement of an electron transport chain. | | 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. | | Hyperglycaemia | A condition characterized by abnormally high blood glucose levels. | | Oxidative decarboxylation | A reaction that involves both oxidation (loss of electrons) and the removal of a carboxyl group as carbon dioxide (CO$_2$). For example, the conversion of pyruvate to acetyl-CoA. | | Krebs cycle (TCA cycle) | A series of metabolic reactions that occur in the mitochondrial matrix, oxidizing acetyl-CoA to CO$_2$, generating ATP, and producing high-energy electron carriers (NADH and FADH$_2$) for oxidative phosphorylation. | | Electron transport chain (ETC) | A series of protein complexes embedded in the inner mitochondrial membrane that transfer electrons from NADH and FADH$_2$ to oxygen, using the released energy to pump protons across the membrane and create a proton gradient. | | Oxidative phosphorylation | The process by which ATP is synthesized through a series of redox reactions in the ETC and ATP synthase, driven by the proton gradient established across the inner mitochondrial membrane. | | ATP synthase | An enzyme complex located in the inner mitochondrial membrane that uses the energy from a proton gradient to catalyze the synthesis of ATP from ADP and inorganic phosphate. | | Proton-motive force | The electrochemical potential energy stored in the proton gradient across a membrane, comprising both a chemical potential (pH difference) and an electrical potential (charge difference). | | Warburg effect | The observation that cancer cells predominantly use glycolysis to produce ATP even in the presence of oxygen, leading to increased lactate production and an acidic tumor microenvironment. | | Endothermic reaction | A reaction that absorbs heat from its surroundings, resulting in products with higher potential energy than the reactants ($\Delta$H > 0). | | Exothermic reaction | A reaction that releases heat into its surroundings, resulting in products with lower potential energy than the reactants ($\Delta$H < 0). | | Catalyst | A substance that increases the rate of a chemical reaction without being consumed itself, by providing an alternative reaction pathway with a lower activation energy. | | Nucleophile | An electron-rich species that donates an electron pair to form a bond. | | Electrophile | An electron-poor species that accepts an electron pair to form a bond. | | Carbohydrates | Organic compounds composed of carbon, hydrogen, and oxygen, often in the ratio (CH$_2$O)$_n$, serving as energy sources and structural components. | | Fatty acids | Long hydrocarbon chains with a terminal carboxylic acid group, forming the building blocks of lipids. | | Lipids | A diverse group of hydrophobic molecules including fats, oils, phospholipids, and steroids, crucial for energy storage, cell membranes, and signaling. | | Membranes | Lipid bilayers forming the boundary of cells and organelles, regulating the passage of substances. | | Nucleotides | The monomers of nucleic acids (DNA and RNA), consisting of a sugar, a phosphate group, and a nitrogenous base. | | Proteins | Macromolecules composed of amino acid subunits linked by peptide bonds, performing a vast array of functions in cells. | | Enzyme | Biological catalysts that accelerate biochemical reactions by lowering activation energy. | | K$_m$ | The Michaelis-Menten constant, representing the substrate concentration at which the reaction rate is half of V$_{max}$. | | V$_{max}$ | The maximum rate of an enzyme-catalyzed reaction when saturated with substrate. | | k$_{cat}$ | The turnover number, indicating how many substrate molecules an enzyme can convert per second at saturation. | | Glycolysis | The metabolic pathway converting glucose to pyruvate, yielding ATP and NADH. | | Gluconeogenesis (GNG) | The synthesis of glucose from non-carbohydrate precursors. | | Investment phase (of glycolysis) | The initial part of glycolysis where ATP is consumed to prepare glucose for cleavage. | | Payoff phase (of glycolysis) | The latter part of glycolysis where ATP and NADH are generated. | | Substrate level phosphorylation | Direct ATP synthesis via phosphate transfer from a high-energy substrate to ADP. | | Normoglycaemia | Normal blood glucose levels. | | Hypoglycaemia | Low blood glucose levels. | | Oxidative decarboxylation | A reaction involving both oxidation and the removal of carbon dioxide. | | Krebs/TCA cycle | A central metabolic pathway in cellular respiration that oxidizes acetyl-CoA, producing ATP, NADH, and FADH$_2$. | | Electron transport chain (ETC) | A series of protein complexes that transfer electrons, generating a proton gradient to drive ATP synthesis. | | Oxidative phosphorylation | ATP synthesis driven by the ETC and proton gradient across the inner mitochondrial membrane. |