Tardigrade.pdf

# General trends and properties of p-block elements The p-block elements exhibit diverse chemical properties due to variations in their electronic configurations, particularly the influence of inner d and f electrons in heavier elements. ### 1.1 General electronic configuration and oxidation states The general valence shell electronic configuration of p-block elements is $ns^2np^x$, where $x$ ranges from 1 to 6. These elements span groups 13 to 18 of the periodic table. The maximum oxidation state displayed by a p-block element is typically equal to the total number of valence electrons (sum of $s$ and $p$ electrons). However, other oxidation states, often differing by two from the group oxidation state, are also observed. For lighter elements in groups 13, 14, and 15, the group oxidation state is the most stable. For heavier elements in these groups, the oxidation state two units less than the group oxidation state becomes progressively more stable, a phenomenon often attributed to the "inert pair effect" [1](#page=1). **Table 11.1 General Electronic Configuration and Oxidation States of p-Block Elements** [2](#page=2). | Group | General electronic configuration | First member of the group | Group oxidation state | Other oxidation states | | :---- | :------------------------------- | :------------------------ | :-------------------- | :--------------------- | | 13 | $ns^2np^1$ | B | +3 | +1 | | 14 | $ns^2np^2$ | C | +4 | +2, -4 | | 15 | $ns^2np^3$ | N | +5 | +3, -3 | | 16 | $ns^2np^4$ | O | +6 | +4, +2, -2 | | 17 | $ns^2np^5$ | F | +7 | +5, +3, +1, -1 | | 18 | $ns^2np^6$ | He | +8 | +6, +4, +2 | > **Tip:** The presence of non-metals and metalloids is exclusive to the p-block [2](#page=2). ### 1.2 Trends in physical and chemical properties #### 1.2.1 Non-metallic to metallic character The non-metallic character decreases down a group in the p-block, with the heaviest element in each group being the most metallic. Non-metals generally possess higher ionization enthalpies and electronegativities compared to metals. Consequently, non-metals tend to form anions, while metals form cations. Compounds formed between highly reactive non-metals and highly reactive metals are typically ionic due to significant electronegativity differences. Conversely, compounds formed between non-metals are largely covalent due to smaller electronegativity differences. A clear illustration of this trend is the acidic nature of non-metal oxides and the basic nature of metal oxides [2](#page=2). #### 1.2.2 Influence of d and f electrons The first member of a p-block group often exhibits distinct properties compared to heavier members. This difference arises from two main factors: atomic size and the presence of d and f orbitals in the inner core of heavier elements. Second-period elements are limited to a maximum covalence of four due to the absence of d orbitals in their valence shell. In contrast, elements from the third period onwards possess vacant d orbitals, allowing them to expand their covalence beyond four. For instance, aluminum can form the ion $[AlF_6]^{3-}$ [2](#page=2). The availability of d orbitals also influences the ability to form $\pi$ bonds. Lighter p-block elements effectively form $p\pi - p\pi$ multiple bonds with themselves and other second-row elements (e.g., C=C, C=O). Heavier elements form weaker $\pi$ bonds, often involving d orbitals ($d\pi - p\pi$ or $d\pi - d\pi$), which contribute less to molecular stability. However, heavier elements can achieve higher coordination numbers than their lighter counterparts in the same oxidation state. For example, in the +5 oxidation state, nitrogen forms NO$_3^-$ (3-coordinate), while phosphorus forms PO$_4^{3-}$ (4-coordinate) where d orbitals are involved in $\pi$ bonding [2](#page=2). ### 1.3 Group 13 elements: The boron family This group displays a significant variation in properties, with boron being a non-metal, aluminum a metal with some non-metallic similarities, and gallium, indium, thallium, and nihonium being primarily metallic [3](#page=3). #### 1.3.1 Electronic configuration and atomic radii The general electronic configuration is $ns^2np^1$. The inner core configurations differ significantly due to the presence of d and f electrons in heavier members. Atomic radii are expected to increase down the group due to the addition of electron shells. However, a deviation occurs where the atomic radius of Gallium (Ga) is smaller than that of Aluminum (Al). This is attributed to the poor screening effect of the 10 d-electrons in Ga, which fail to adequately counteract the increased nuclear charge [3](#page=3). #### 1.3.2 Ionization enthalpy Ionization enthalpy values do not decrease smoothly down the group. The discontinuity between Al and Ga, and between In and Tl, is due to the poor screening effect of d and f electrons which do not compensate for the increased nuclear charge. The order of ionization enthalpies is $\Delta_1 H < \Delta_2 H < \Delta_3 H$, and the sum of the first three ionization enthalpies is very high, favoring covalent bond formation for Boron [3](#page=3). #### 1.3.3 Electronegativity Electronegativity decreases from B to Al and then increases marginally for heavier elements due to discrepancies in atomic size [4](#page=4). #### 1.3.4 Physical properties Boron is a hard, black, non-metallic solid with a very high melting point due to strong crystalline lattice forces. The other members are soft metals with lower melting points and high electrical conductivity. Gallium has a remarkably low melting point (303 K) but a high boiling point (2676 K), making it useful for high-temperature measurements. Density increases down the group from boron to thallium [4](#page=4). #### 1.3.5 Chemical properties **Oxidation states and reactivity:** Boron's high ionization enthalpy sum favors covalent compound formation, while Aluminium can form $Al^{3+}$ ions. For heavier elements (Ga, In, Tl), the $ns^2$ electrons are held tightly by the increased effective nuclear charge (inert pair effect), leading to the observation of both +1 and +3 oxidation states. The +1 oxidation state becomes progressively more stable down the group: Al < Ga < In < Tl. Thallium predominantly exhibits the +1 oxidation state, while its +3 state is highly oxidizing [4](#page=4). In the +3 oxidation state, compounds are generally covalent and electron-deficient, acting as Lewis acids by accepting electron pairs. For example, $BCl_3$ readily accepts a lone pair from $NH_3$. Aluminium chloride forms a dimer, $Al_2Cl_6$. Many trivalent compounds hydrolyze in water to form species like $[M(OH)_4]^-$ or $[M(H_2O)_6]^{3+}$ [5](#page=5). **Reactivity towards air:** Crystalline boron is unreactive. Aluminium forms a protective oxide layer. Boron trioxide ($B_2O_3$) is acidic, aluminum and gallium oxides are amphoteric, and indium and thallium oxides are basic [5](#page=5). **Reactivity towards acids and alkalies:** Boron is unreactive towards acids and alkalies. Aluminium dissolves in mineral acids and aqueous alkalies, exhibiting amphoteric character [6](#page=6). **Reactivity towards halogens:** These elements react with halogens to form trihalides ($EX_3$), except for Thallium which does not readily form $TII_3$ [6](#page=6). #### 1.3.6 Some important compounds of boron * **Borax ($Na_2B_4O_7 \cdot 10H_2O$):** Dissolves in water to give an alkaline solution. Upon heating, it forms metaborate ($NaBO_2$) and boric anhydride ($B_2O_3$). Borax bead test is used to identify transition metals due to the characteristic colors of their metaborates [6](#page=6). * **Orthoboric acid ($H_3BO_3$):** A weak monobasic acid that acts as a Lewis acid by accepting hydroxide ions from water. On heating, it forms metaboric acid ($HBO_2$) and then boric oxide ($B_2O_3$) [6](#page=6) [7](#page=7). * **Diborane ($B_2H_6$):** A colorless, toxic gas that is highly flammable in air. It reacts with water to form boric acid and with Lewis bases to form borane adducts ($BH_3 \cdot L$). Reaction with ammonia yields borazine ($B_3N_3H_6$), also known as "inorganic benzene". Diborane has a unique structure with two bridging hydrogen atoms forming three-center-two-electron bonds (banana bonds). It forms hydridoborate salts like $[BH_4]^-$ which are used as reducing agents [7](#page=7) [8](#page=8). #### 1.3.7 Uses of boron and aluminium Boron fibers are used in bullet-proof vests and aircraft materials. Boron isotopes are used in nuclear reactors. Borax and boric acid are used in making heat-resistant glass, fiberglass, and as fluxes. Aluminium is used extensively in industry for its high tensile strength and conductivity, forming alloys for various applications like packing, utensils, and transportation [8](#page=8). ### 1.4 Group 14 elements: The carbon family This group consists of carbon, silicon, germanium, tin, lead, and flerovium. Carbon is a unique element due to its versatile bonding capabilities [8](#page=8). #### 1.4.1 Electronic configuration and atomic radii The valence shell electronic configuration is $ns^2np^2$. There is a significant increase in covalent radius from C to Si, followed by a smaller increase from Si to Pb due to the presence of filled d and f orbitals in heavier members [9](#page=9). #### 1.4.2 Ionization enthalpy The first ionization enthalpy of Group 14 elements is higher than that of Group 13 elements. Ionization enthalpies generally decrease down the group, with minor variations due to the shielding effect of d and f electrons and increasing atomic size [9](#page=9). #### 1.4.3 Electronegativity Elements in this group are slightly more electronegative than Group 13 elements. Electronegativity values are similar from Si to Pb [9](#page=9). #### 1.4.4 Physical properties All members of Group 14 are solids. Carbon and silicon are non-metals, germanium is a metalloid, and tin and lead are soft metals with low melting points. Their melting and boiling points are generally higher than those of Group 13 elements [10](#page=10). #### 1.4.5 Chemical properties **Oxidation states and reactivity:** The common oxidation states are +4 and +2, with carbon also exhibiting negative oxidation states. Compounds in the +4 oxidation state are generally covalent due to the high sum of ionization enthalpies. The tendency to exhibit the +2 oxidation state increases down the group (Ge < Sn < Pb) due to the inert pair effect. Carbon and silicon predominantly show the +4 oxidation state. In the +2 state, tin is a reducing agent, while lead compounds in the +2 state are stable and in the +4 state are strong oxidizing agents. Elements heavier than carbon can expand their covalence beyond four due to the availability of d orbitals, leading to hydrolysis of their halides and complex formation [10](#page=10). **Reactivity towards oxygen:** All members form oxides ($MO$ and $MO_2$) upon heating in oxygen. Higher oxidation state oxides are generally more acidic. $CO_2$, $SiO_2$, and $GeO_2$ are acidic, while $SnO_2$ and $PbO_2$ are amphoteric. Among monoxides, CO is neutral, GeO is acidic, and SnO and PbO are amphoteric [10](#page=10). **Reactivity towards water:** Carbon, silicon, and germanium are unaffected by water. Tin decomposes steam, and lead is generally unaffected due to a protective oxide film [10](#page=10). **Reactivity towards halogens:** They form halides of formula $MX_2$ and $MX_4$. Most $MX_4$ compounds are covalent and tetrahedral, except for $SnF_4$ and $PbF_4$ which are ionic. $PbI_4$ is unstable. Stability of dihalides increases down the group ($GeX_2 < SnX_2 < PbX_2$). Most tetrachlorides hydrolyze in water, with $CCl_4$ being an exception due to the absence of d orbitals [10](#page=10). #### 1.4.6 Important anomalous behaviour of carbon Carbon differs from other members due to its small size, high electronegativity, high ionization enthalpy, and the absence of d orbitals. This limits its maximum covalence to four. Carbon uniquely forms stable $p\pi - p\pi$ multiple bonds with itself and other small, electronegative atoms (e.g., C=C, C=O). It exhibits catenation, the ability to form chains and rings, due to the strength of C-C bonds. The tendency for catenation decreases down the group: C >> Si > Ge ~ Sn [11](#page=11). #### 1.4.7 Allotropes of carbon Carbon exists in various allotropic forms: * **Diamond:** A hardest crystalline form where each carbon atom is $sp^3$ hybridized and tetrahedrally bonded to four other carbon atoms, forming a rigid 3D network of strong covalent bonds [11](#page=11) [12](#page=12). * **Graphite:** Has a layered structure where carbon atoms are arranged in planar hexagonal rings. Within layers, $sp^2$ hybridized carbon atoms form sigma bonds, and the fourth electron is delocalized, allowing for electrical conductivity along the sheets. Layers are held by weak van der Waals forces, making graphite soft and slippery [12](#page=12). * **Fullerenes:** Cage-like molecules, with C$_{60}$ (Buckminsterfullerene) being the most famous, resembling a soccer ball. They have a smooth structure without 'dangling' bonds and exhibit aromatic character [12](#page=12) [13](#page=13). Other impure forms include carbon black, coke, and charcoal [13](#page=13). #### 1.4.8 Some important compounds of carbon and silicon * **Oxides of Carbon:** * **Carbon Monoxide (CO):** A colorless, odorless, and toxic gas, it is a powerful reducing agent used in metal extraction. It forms a stable complex with hemoglobin, leading to its toxicity [13](#page=13) [14](#page=14). * **Carbon Dioxide ($CO_2$):** An acidic gas that forms carbonic acid in water. It is essential for photosynthesis and maintains blood pH. Increased atmospheric $CO_2$ contributes to the greenhouse effect. Solid $CO_2$ (dry ice) is used as a refrigerant. It has a linear structure with sp hybridization of the carbon atom [14](#page=14). * **Silicon Dioxide ($SiO_2$):** Exists in various crystalline forms like quartz. It is a 3D network solid where each silicon is tetrahedrally bonded to four oxygen atoms, and each oxygen is bonded to two silicon atoms. It is highly unreactive but attacked by HF and NaOH [14](#page=14) [15](#page=15). * **Silicones:** Organosilicon polymers with a $(-R_2SiO-)$ repeating unit. They are water-repellent, thermally stable, and used as sealants, greases, and electrical insulators [15](#page=15) [16](#page=16). * **Silicates:** Minerals with the basic structural unit being the $SiO_4^{4-}$ tetrahedron. These units can link to form chains, rings, sheets, or 3D networks. Zeolites, a type of aluminosilicate, are used as catalysts and ion exchangers [16](#page=16). --- # Group 13 elements: The Boron family The Group 13 elements, also known as the boron family, exhibit a significant variation in properties, ranging from non-metallic boron to predominantly metallic gallium, indium, and thallium, with aluminum being a metal that shares chemical similarities with boron. ### 2.1 Occurrence and abundance Boron is a relatively rare element, found primarily as orthoboric acid ($H_3BO_3$) and borax ($Na_2B_4O_7 \cdot 10H_2O$). Aluminum is the most abundant metal and the third most abundant element in Earth's crust, with bauxite ($Al_2O_3 \cdot 2H_2O$) and cryolite ($Na_3AlF_6$) as its major minerals. Gallium, indium, and thallium are less abundant. Nihonium (Nh) is a synthetically prepared radioactive element with atomic number 113 [3](#page=3). ### 2.1.1 Electronic configuration The general outer electronic configuration of Group 13 elements is $ns^2 np^1$. Boron and aluminum have noble gas cores, while gallium and indium have a noble gas core plus 10 d-electrons, and thallium has a noble gas core plus 14 f-electrons and 10 d-electrons. These complex electronic structures influence their properties and chemistry [3](#page=3). ### 2.1.2 Atomic radii As one moves down the group, atomic radius generally increases due to the addition of electron shells. However, a notable deviation occurs with gallium, whose atomic radius (135 pm) is smaller than that of aluminum (143 pm) (#page=3, page=4). This is attributed to the poor screening effect of the 10 d-electrons in gallium, which fail to adequately shield the outer electrons from the increased nuclear charge [3](#page=3) [4](#page=4). ### 2.1.3 Ionization enthalpy Ionization enthalpy values do not decrease smoothly down the group. The decrease from boron to aluminum is due to increasing atomic size. The discontinuities observed between aluminum and gallium, and between indium and thallium, are due to the ineffective screening by d- and f-electrons, which cannot fully compensate for the increased nuclear charge. The first three ionization enthalpies ($ \Delta H_1 < \Delta H_2 < \Delta H_3 $) are generally high for these elements, particularly for boron, affecting their chemical behavior (#page=3, page=4) [3](#page=3) [4](#page=4). ### 2.1.4 Electronegativity Electronegativity initially decreases from boron to aluminum and then marginally increases for gallium, indium, and thallium. This variation is linked to the discrepancies in atomic sizes [4](#page=4). ### 2.1.5 Physical properties Boron is a non-metallic, black, and extremely hard solid with numerous allotropic forms and a high melting point due to its strong crystalline lattice. The other elements in the group are soft metals with lower melting points and good electrical conductivity. Gallium has an unusually low melting point (303 K), allowing it to exist as a liquid in summer, and a high boiling point (2676 K), making it suitable for high-temperature measurements. The density of the elements increases down the group from boron to thallium [4](#page=4). ### 2.1.6 Chemical properties #### Oxidation state and trends in chemical reactivity Boron, due to its high sum of the first three ionization enthalpies, primarily forms covalent compounds rather than +3 ions (#page=4, page=5). Aluminum, with a lower sum of ionization enthalpies, readily forms $Al^{3+}$ ions and is highly electropositive (#page=4, page=5). Down the group, the inert pair effect, caused by the poor shielding of d- and f-electrons, leads to the increasing stability of the +1 oxidation state relative to the +3 state, with Tl in the +1 oxidation state being predominant (#page=4, page=5). Compounds in the +1 oxidation state are generally more ionic than those in the +3 state [4](#page=4) [5](#page=5). In the trivalent state, boron compounds like $BF_3$ are electron-deficient, possessing only six electrons around the central atom, and thus act as Lewis acids, readily accepting electron pairs (#page=5, page=6). This tendency to act as a Lewis acid decreases with increasing atomic size down the group. For example, $BCl_3$ reacts with $NH_3$ to form an adduct. Aluminum chloride ($AlCl_3$) stabilizes by forming a dimer, $Al_2Cl_6$ [5](#page=5) [6](#page=6). Most trivalent compounds are covalent and undergo hydrolysis in water, forming tetrahedral $[M(OH)_4]^-$ species, where the central atom (M) is $sp^3$ hybridized. Aluminum chloride in acidified aqueous solutions forms octahedral $[Al(H_2O)_6]^{3+}$ ions, where Al is $sp^3d^2$ hybridized, involving its 3d orbitals [5](#page=5). > **Tip:** The inability of boron to expand its octet beyond four is due to the absence of d orbitals, limiting its maximum covalence to 4, unlike heavier elements which can exceed this limit (#page=6, page=8) [6](#page=6) [8](#page=8). ##### Reactivity towards air Crystalline boron is unreactive. Aluminum forms a protective oxide layer on its surface, preventing further reaction. Both amorphous boron and aluminum react with air and nitrogen at high temperatures to form their respective oxides ($B_2O_3$, $Al_2O_3$) and nitrides (BN). The nature of the oxides varies: $B_2O_3$ is acidic, $Al_2O_3$ and $Ga_2O_3$ are amphoteric, and $In_2O_3$ and $Tl_2O_3$ are basic [5](#page=5). ##### Reactivity towards acids and alkalies Boron does not react with acids or alkalies even at moderate temperatures. Aluminum, however, dissolves in mineral acids and aqueous alkalies, exhibiting amphoteric behavior. It reacts with dilute HCl to liberate dihydrogen gas and with aqueous alkali to form sodium tetrahydroxoaluminate(III) and dihydrogen. Concentrated nitric acid passivates aluminum by forming a protective oxide layer [5](#page=5). ##### Reactivity towards halogens Group 13 elements react with halogens to form trihalides ($EX_3$), except for $T1I_3$ [6](#page=6). > **Example:** White fumes around a bottle of anhydrous aluminum chloride are due to its partial hydrolysis with atmospheric moisture, releasing HCl gas, which appears white [6](#page=6). ### 2.2 Anomalous properties and important compounds of boron #### 2.2.1 Borax Borax, $Na_2B_4O_7 \cdot 10H_2O$, is a key boron compound. Its correct formula is $Na_2[B_4O_5(OH)_4 \cdot 8H_2O$ (#page=6, page=8). It dissolves in water to yield an alkaline solution due to the formation of orthoboric acid and sodium hydroxide. Upon heating, borax loses water, swells, and on further heating, forms a transparent liquid that solidifies into a glass-like material known as borax bead. This bead test is used to identify transition metals due to the characteristic colors of their metaborate compounds [6](#page=6) [8](#page=8). #### 2.2.2 Orthoboric acid Orthoboric acid, $H_3BO_3$, is a white, crystalline solid with a soapy touch. It is sparingly soluble in cold water but highly soluble in hot water. It can be prepared by acidifying an aqueous solution of borax or by the hydrolysis of most boron compounds. $H_3BO_3$ has a layered structure where planar $BO_3$ units are linked by hydrogen bonds. It is a weak monobasic acid and acts as a Lewis acid by accepting a hydroxyl ion from water, rather than releasing a proton directly. On heating, it forms metaboric acid ($HBO_2$) and then boric oxide ($B_2O_3$) [6](#page=6) [7](#page=7). > **Tip:** Boric acid is considered a weak acid because it accepts $OH^-$ ions from water to complete its octet, thereby releasing $H^+$ ions, rather than releasing its own protons [7](#page=7). #### 2.2.3 Diborane ($B_2H_6$) Diborane is the simplest known boron hydride. It can be prepared by treating boron trifluoride with lithium aluminum hydride, by oxidizing sodium borohydride with iodine, or industrially by reacting $BF_3$ with sodium hydride. Diborane is a colorless, highly toxic gas that is spontaneously flammable in air and burns in oxygen, releasing significant energy. It readily hydrolyzes in water to form boric acid and dihydrogen. Diborane undergoes cleavage reactions with Lewis bases to form borane adducts ($BH_3 \cdot L$). Reaction with ammonia yields borazine ($B_3N_3H_6$), also known as "inorganic benzene" [7](#page=7). The structure of diborane consists of four terminal hydrogen atoms and two boron atoms lying in a plane, with two bridging hydrogen atoms above and below this plane. The four terminal B-H bonds are normal two-center-two-electron bonds, while the two B-H-B bridge bonds are three-center-two-electron bonds, also referred to as banana bonds (#page=7, page=8). Each boron atom is $sp^3$ hybridized in diborane [7](#page=7) [8](#page=8). Boron also forms tetrahydridoborate ions, $[BH_4]^-$, which are important reducing agents in organic synthesis [8](#page=8). ### 2.3 Uses of boron and aluminum and their compounds Boron, due to its hardness, high melting point, low density, and low electrical conductivity, finds applications in bullet-proof vests and light composite materials for aircraft. The boron-10 isotope is used in the nuclear industry as neutron shields and control rods. Borax and boric acid are used in the manufacture of heat-resistant glasses, glass-wool, and fiberglass. Borax also serves as a flux for soldering and as a component in glazes and medicinal soaps. Orthoboric acid is used as a mild antiseptic [8](#page=8). Aluminum is a bright, silvery-white metal with high tensile strength, electrical, and thermal conductivity, making it twice as conductive as copper on a weight-to-weight basis. It is widely used in industry, forming alloys with various metals. Aluminum and its alloys are shaped into various forms for packaging, utensils, construction, and transportation industries. However, the use of aluminum and its compounds for domestic purposes has decreased due to concerns about their toxicity [8](#page=8). --- # Group 14 elements: The Carbon family The group 14 elements, comprising carbon, silicon, germanium, tin, and lead, share a common valence shell electronic configuration of $ns^2 np^2$. These elements exhibit a fascinating range of properties, from the non-metallic versatility of carbon to the metallic characteristics of tin and lead, with silicon and germanium acting as metalloids [8](#page=8) [9](#page=9). ### 3.1 Electronic configuration The general valence shell electronic configuration for group 14 elements is $ns^2 np^2$. Flerovium (Fl), with atomic number 114, has the configuration $[Rn 5f^{14} 6d^{10} 7s^2 7p^2$, but its chemistry is not yet established due to its synthetic nature and short half-life [9](#page=9). ### 3.2 Atomic and physical properties #### 3.2.1 Covalent radius There is a significant increase in covalent radius from carbon to silicon. Beyond silicon, the increase in covalent radius from silicon to lead is less pronounced. This gradual increase is attributed to the presence of completely filled $d$ and $f$ orbitals in the heavier elements, which shield the valence electrons less effectively than anticipated [9](#page=9). #### 3.2.2 Ionization enthalpy The first ionization enthalpies of group 14 elements are higher than those of their corresponding group 13 counterparts. As expected, ionization enthalpy generally decreases down the group. However, there is a slight decrease from silicon to germanium to tin, followed by a small increase from tin to lead. This anomaly is a result of the poor shielding effect of the intervening $d$ and $f$ orbitals and the increase in atomic size [9](#page=9). #### 3.2.3 Electronegativity Due to their relatively small atomic sizes, group 14 elements are slightly more electronegative than group 13 elements. The electronegativity values remain relatively constant for silicon, germanium, and tin, with a slight increase for lead [9](#page=9). #### 3.2.4 Physical properties All members of group 14 are solids. Carbon and silicon are non-metals, germanium is a metalloid, and tin and lead are soft metals with low melting points. The melting and boiling points of group 14 elements are considerably higher than those of the corresponding elements in group 13 [10](#page=10). ### 3.3 Chemical properties #### 3.3.1 Oxidation states and trends in chemical reactivity Group 14 elements possess four valence electrons, leading to common oxidation states of +4 and +2. Carbon can also exhibit negative oxidation states. Compounds in the +4 oxidation state are generally covalent due to the high ionization enthalpies required to remove all four valence electrons. The tendency to exhibit the +2 oxidation state increases down the group from germanium to lead, attributed to the inert pair effect where the $ns^2$ electrons are reluctant to participate in bonding [10](#page=10) [17](#page=17). * **Carbon and Silicon:** Primarily exhibit the +4 oxidation state [10](#page=10). * **Germanium:** Forms stable compounds in the +4 state and a few in the +2 state [10](#page=10). * **Tin:** Forms stable compounds in both +4 and +2 oxidation states, with Sn(II) acting as a reducing agent [10](#page=10). * **Lead:** Exhibits stable compounds in the +2 oxidation state, while Pb(IV) compounds are strong oxidizing agents [10](#page=10) [17](#page=17). The maximum covalence of carbon is limited to four due to the absence of $d$ orbitals. However, heavier elements like silicon, germanium, tin, and lead can expand their covalence beyond four by utilizing their vacant $d$ orbitals, enabling them to form complexes [11](#page=11). #### 3.3.2 Reactivity towards oxygen All group 14 elements react with oxygen when heated to form oxides of the general formulas MO (monoxide) and $MO_2$ (dioxide). $SiO$ is only stable at high temperatures. Higher oxidation state oxides are generally more acidic than lower oxidation state oxides [10](#page=10). * **Dioxides:** $CO_2$, $SiO_2$, and $GeO_2$ are acidic. $SnO_2$ and $PbO_2$ are amphoteric [10](#page=10). * **Monoxides:** $CO$ is neutral. $GeO$ is acidic. $SnO$ and $PbO$ are amphoteric [10](#page=10). #### 3.3.3 Reactivity towards water Carbon, silicon, and germanium are unaffected by water. Tin decomposes steam to produce tin dioxide and dihydrogen gas [10](#page=10): $$Sn + 2H_2O \xrightarrow{\Delta} SnO_2 + 2H_2$$ Lead is unaffected by water, likely due to the formation of a protective oxide film [10](#page=10). #### 3.3.4 Reactivity towards halogens Group 14 elements can form halides of the general formulas $MX_2$ and $MX_4$, where X is a halogen. Except for carbon, other members react directly with halogens under suitable conditions. Most $MX_4$ halides are covalent and tetrahedral in shape due to $sp^3$ hybridization of the central atom. $SnF_4$ and $PbF_4$ are exceptions, being ionic. $PbI_4$ does not exist due to insufficient energy released to form the four unpaired electrons required for bonding [10](#page=10). Dihalides ($MX_2$) are also formed by heavier members from germanium to lead, with the stability of dihalides increasing down the group. For example, $GeX_4$ is more stable than $GeX_2$, while $PbX_2$ is more stable than $PbX_4$ [10](#page=10). Except for $CCl_4$, other tetrachlorides are readily hydrolyzed by water. This hydrolysis occurs because the central atom can accept electron pairs from water molecules into its $d$ orbitals. For instance, $SiCl_4$ hydrolysis involves the silicon atom accepting a lone pair from water, eventually leading to the formation of silicic acid, $Si(OH)_4$ [10](#page=10) [11](#page=11). $$SiCl_4 + 4H_2O \rightarrow Si(OH)_4 + 4HCl$$ > **Tip:** While $SiF_6^{2-}$ is known, $SiCl_6^{2-}$ is not commonly formed. This is due to the limited size of the silicon atom being unable to accommodate six large chloride ions and weaker interaction between the lone pair of chloride ions and $Si^{4+}$ [11](#page=11). ### 3.4 Anomalous behaviour of carbon Carbon, the first member of group 14, displays unique properties compared to the rest of the group. This is due to its smaller atomic size, higher electronegativity, higher ionization enthalpy, and the absence of $d$ orbitals in its valence shell [11](#page=11). * **Limited Covalence:** Carbon can accommodate a maximum of four pairs of electrons, limiting its covalence to four [11](#page=11). * **Multiple Bond Formation:** Carbon has a unique ability to form $p\pi-p\pi$ multiple bonds with itself and other small, electronegative atoms (e.g., $C=C$, $C\equiv C$, $C=O$, $C \equiv N$). Heavier elements cannot form such bonds effectively due to the larger size and diffuse nature of their atomic orbitals [11](#page=11). * **Catenation:** Carbon atoms readily link with each other through covalent bonds to form chains and rings, a property known as catenation. This is facilitated by the strength of the C-C bond. The tendency for catenation decreases down the group in the order $C >> Si > Ge \approx Sn$, with lead showing no catenation. The bond enthalpies illustrate this trend [11](#page=11): * C-C: 348 kJ mol$^{-1}$ * Si-Si: 297 kJ mol$^{-1}$ * Ge-Ge: 260 kJ mol$^{-1}$ * Sn-Sn: 240 kJ mol$^{-1}$ [11](#page=11). ### 3.5 Allotropes of carbon Carbon exists in several allotropic forms, both crystalline and amorphous [11](#page=11). #### 3.5.1 Diamond Diamond has a giant covalent structure where each carbon atom is $sp^3$ hybridized and bonded to four other carbon atoms in a tetrahedral arrangement. The C-C bond length is 154 pm. The strong, directional covalent bonds throughout the lattice make diamond the hardest known substance [12](#page=12). > **Example:** Diamond is used as an abrasive for sharpening tools, in dyes, and for making filaments in electric light bulbs [12](#page=12). > **Tip:** Diamond's high melting point, despite being covalent, is due to the extensive three-dimensional network of strong C-C bonds that require significant energy to break [12](#page=12). #### 3.5.2 Graphite Graphite has a layered structure. Within each layer, carbon atoms form planar hexagonal rings with a C-C bond length of 141.5 pm. Each carbon atom is $sp^2$ hybridized, forming sigma bonds with three neighboring carbon atoms. The fourth electron is delocalized in a $\pi$ system across the entire sheet, allowing graphite to conduct electricity along the layers. The layers are held together by weak van der Waals forces (340 pm separation) [12](#page=12). Graphite cleaves easily between layers, making it soft and slippery. This property makes it useful as a dry lubricant in high-temperature machinery where oils are unsuitable [12](#page=12). #### 3.5.3 Fullerenes Fullerenes are a class of allotropes formed by heating graphite in an electric arc in the presence of inert gases. The most common fullerene is $C_{60}$, known as Buckminsterfullerene, which has a soccer ball-like shape. It consists of 20 six-membered rings and 12 five-membered rings, with all carbon atoms being $sp^2$ hybridized and forming three sigma bonds. The remaining electrons are delocalized, giving the molecule aromatic character. Fullerenes are considered pure allotropes as they lack 'dangling' bonds [12](#page=12) [13](#page=13). > **Example:** Other fullerenes like $C_{70}$ and those with even numbers of carbon atoms up to 350 also exist [12](#page=12). Graphite is the thermodynamically most stable allotrope of carbon, with its standard enthalpy of formation ($ \Delta H_f^\circ $) taken as zero. The standard enthalpies of formation for diamond and $C_{60}$ are 1.90 kJ mol$^{-1}$ and 38.1 kJ mol$^{-1}$, respectively [13](#page=13). Amorphous forms of carbon, such as carbon black, coke, and charcoal, are impure forms of graphite or fullerenes [13](#page=13). #### 3.5.4 Uses of carbon allotropes * **Graphite fibers:** Used in high-strength, lightweight composites for products like tennis rackets and aircraft [13](#page=13). * **Graphite:** Used as electrodes in batteries and electrolysis due to its conductivity. Graphite crucibles are resistant to dilute acids and alkalis [13](#page=13). * **Activated charcoal:** Highly porous and used for adsorbing poisonous gases, in water filters, and in air conditioning for odor control [13](#page=13). * **Carbon black:** Used as a black pigment and as a filler in automobile tires [13](#page=13). * **Coke:** Used as a fuel and a reducing agent in metallurgy [13](#page=13). * **Diamond:** A precious gemstone used in jewelry, measured in carats [13](#page=13). ### 3.6 Important compounds of carbon and silicon #### 3.6.1 Oxides of carbon Two key oxides are carbon monoxide ($CO$) and carbon dioxide ($CO_2$) [13](#page=13). * **Carbon monoxide ($CO$)**: Prepared by the limited oxidation of carbon, dehydration of formic acid, or reaction of steam with hot coke (forming water gas, a mixture of $CO$ and $H_2$) [13](#page=13). $$2C(s) + O_2(g) \xrightarrow{\Delta} 2CO(g)$$ $$HCOOH \xrightarrow{373K, conc. H_2SO_4} H_2O + CO$$ $$C(s) + H_2O(g) \xrightarrow{473-1273K} CO(g) + H_2(g) \quad (\text{Water gas})$$ Water gas and producer gas (formed when air is used instead of steam, $2C(s) + O_2(g) + 4N_2(g) \xrightarrow{1273K} 2CO(g) + 4N_2(g)$) are important industrial fuels. $CO$ is a colorless, odorless, and nearly insoluble gas, acting as a powerful reducing agent used in metal extraction [13](#page=13). $$Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(s) + 3CO_2(g)$$ $$ZnO(s) + CO(g) \rightarrow Zn(s) + CO_2(g)$$ The $CO$ molecule has one sigma and two $\pi$ bonds. The lone pair on carbon allows it to act as a donor, forming metal carbonyls. $CO$ is highly poisonous because it forms a complex with hemoglobin that is about 300 times more stable than the oxygen-hemoglobin complex, preventing oxygen transport [13](#page=13). * **Carbon dioxide ($CO_2$)**: Prepared by the complete combustion of carbon and carbon-containing fuels or by the reaction of dilute acids with carbonates [14](#page=14). $$C(s) + O_2(g) \rightarrow CO_2(g)$$ $$CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$$ $$CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)$$ $CO_2$ forms carbonic acid ($H_2CO_3$) in water, a weak dibasic acid crucial for maintaining blood pH [14](#page=14). $$H_2CO_3(aq) + H_2O(l) \rightleftharpoons HCO_3^-(aq) + H_3O^+(aq)$$ $$HCO_3^-(aq) + H_2O(l) \rightleftharpoons CO_3^{2-}(aq) + H_3O^+(aq)$$ Photosynthesis removes atmospheric $CO_2$ for carbohydrate synthesis [14](#page=14). $$6CO_2 + 12H_2O \xrightarrow{Chlorophyll, hv} C_6H_{12}O_6 + 6O_2 + 6H_2O$$ Solid $CO_2$ (dry ice) is used as a refrigerant and in carbonating soft drinks. Its density makes it useful as a fire extinguisher. The $CO_2$ molecule is linear, with the carbon atom undergoing $sp$ hybridization, forming two sigma bonds and two $\pi$ bonds with oxygen atoms [14](#page=14). $$: \ddot{O}-C \equiv O: \leftrightarrow \underset{+}{:} \ddot{O}=C=\ddot{O}: \leftrightarrow :O \equiv C-\ddot{O}: \quad (\text{Resonance structures})$$ Increased atmospheric $CO_2$ from fossil fuel combustion is a concern for the greenhouse effect [14](#page=14). #### 3.6.2 Silicon dioxide ($SiO_2$) $SiO_2$, commonly known as silica, is a covalent, three-dimensional network solid. Each silicon atom is tetrahedrally bonded to four oxygen atoms, and each oxygen atom is bonded to two silicon atoms, forming eight-membered rings. This structure results in a very high Si-O bond enthalpy, making silica largely unreactive towards halogens, dihydrogen, and most acids and metals. However, it is attacked by HF and NaOH [14](#page=14) [15](#page=15). $$SiO_2 + 2NaOH \rightarrow Na_2SiO_3 + H_2O$$ $$SiO_2 + 4HF \rightarrow SiF_4 + 2H_2O$$ Quartz is used for its piezoelectric properties in clocks and communication systems. Silica gel is a drying agent and a support for chromatography. Kieselghur, an amorphous form, is used in filtration [15](#page=15). #### 3.6.3 Silicones Silicones are organosilicon polymers with the repeating unit $(-R_2SiO-)$. They are synthesized from alkyl or aryl substituted silicon chlorides, $R_nSiCl_{(4-n)}$. Hydrolysis of dimethyl dichlorosilane, $(CH_3)_2SiCl_2$, followed by condensation polymerization, yields straight-chain silicone polymers [15](#page=15). $$2CH_3Cl + Si \xrightarrow{Cu \text{ powder}, 570K} (CH_3)_2SiCl_2$$ $$(CH_3)_2SiCl_2 + 2H_2O \rightarrow (CH_3)_2Si(OH)_2 + 2HCl$$ $$n(CH_3)_2Si(OH)_2 \xrightarrow{-H_2O} \left[ -O-Si(CH_3)_2- \right]_n \quad (\text{Silicone})$$ The polymer chain length can be controlled by adding terminating agents like $(CH_3)_3SiCl$. Silicones are water-repellent due to the non-polar alkyl groups and possess high thermal stability, dielectric strength, and resistance to oxidation and chemicals. They are used as sealants, greases, electrical insulators, for waterproofing, and in surgical and cosmetic applications [16](#page=16). > **Tip:** Silicones are hydrophobic (water-repellant) polymers where alkyl or phenyl groups surround the silicon-oxygen backbone [16](#page=16). #### 3.6.4 Silicates Silicates are a major class of minerals, with the basic structural unit being the $SiO_4^{4-}$ tetrahedron. In silicates, these units can be discrete or linked by sharing oxygen atoms, forming chains, rings, sheets, or three-dimensional networks. The negative charge is neutralized by metal cations. Examples include feldspar, zeolites, mica, and asbestos. Glass and cement are important man-made silicates [16](#page=16). #### 3.6.5 Zeolites Zeolites are aluminosilicates formed when aluminum atoms substitute for some silicon atoms in the $SiO_2$ network, creating an overall negative charge balanced by cations like $Na^+$, $K^+$, or $Ca^{2+}$. Zeolites like ZSM-5 are used as catalysts in petrochemical industries for cracking hydrocarbons and isomerisation, and hydrated zeolites are used as ion exchangers for water softening [16](#page=16). --- # Important compounds of boron, carbon, and silicon This section explores the chemistry and applications of specific, significant compounds formed by boron, carbon, and silicon. ### 4.1 Compounds of boron Boron forms several important compounds, including borax, orthoboric acid, and diborane [7](#page=7). #### 4.1.1 Borax Borax, with the formula $Na_2B_4O_7 \cdot 10H_2O$, is a key boron compound. Its correct formulation is $Na_2[B_4O_5(OH)_4 \cdot 8H_2O$ as it contains tetranuclear units. Dissolving borax in water yields an alkaline solution due to its hydrolysis [6](#page=6): $$Na_2B_4O_7 + 7H_2O \rightarrow 2NaOH + 4H_3BO_3$$ [6](#page=6). Upon heating, borax initially loses water molecules and then melts into a transparent liquid that solidifies into a glass-like material called borax bead. This borax bead can be used to identify transition metals due to the characteristic colours of their metaborates, formed after further heating [6](#page=6): $$Na_2B_4O_7 \xrightarrow{\Delta} Na_2B_4O_7 \xrightarrow{\Delta} 2NaBO_2 + B_2O_3$$ [6](#page=6). Sodium metaborate Boric anhydride An example of the borax bead test involves heating borax with cobalt oxide ($CoO$) on a platinum wire loop, resulting in a blue colored $Co(BO_2)_2$ bead [6](#page=6). #### 4.1.2 Orthoboric acid Orthoboric acid ($H_3BO_3$) is a white crystalline solid with a soapy touch. It is sparingly soluble in cold water but highly soluble in hot water. It can be prepared by acidifying an aqueous solution of borax [6](#page=6): $$Na_2B_4O_7 + 2HCl + 5H_2O \rightarrow 2NaCl + 4B(OH)_3$$ [7](#page=7). Orthoboric acid is also formed through the hydrolysis of most boron compounds, such as halides and hydrides. It possesses a layered structure where planar $BO_3$ units are linked by hydrogen bonds [7](#page=7). > **Tip:** Boric acid is considered a weak acid because it does not readily release $H^+$ ions. Instead, it acts as a Lewis acid by accepting electrons from a hydroxyl ion ($OH^-$) from water, thereby releasing $H^+$ ions: > $$B(OH)_3 + 2HOH \rightarrow [B(OH)_4]^- + H_3O^+$$ [7](#page=7). On heating above $370K$, orthoboric acid first forms metaboric acid ($HBO_2$), which further decomposes into boric oxide ($B_2O_3$) upon stronger heating [7](#page=7). $$H_3BO_3 \xrightarrow{>370K} HBO_2 \xrightarrow{>420K} B_2O_3$$ [7](#page=7). #### 4.1.3 Diborane ($B_2H_6$) Diborane is the simplest known boron hydride. It can be synthesized by reacting boron trifluoride ($BF_3$) with lithium aluminum hydride ($LiAlH_4$) in diethyl ether [7](#page=7): $$4BF_3 + 3LiAlH_4 \rightarrow 2B_2H_6 + 3LiF + 3AlF_3$$ [7](#page=7). A convenient laboratory preparation involves the oxidation of sodium borohydride ($NaBH_4$) with iodine [7](#page=7): $$2NaBH_4 + I_2 \rightarrow B_2H_6 + 2NaI + H_2$$ [7](#page=7). Industrially, diborane is produced by reacting $BF_3$ with sodium hydride ($NaH$) [7](#page=7): $$2BF_3 + 6NaH \xrightarrow{450K} B_2H_6 + 6NaF$$ [7](#page=7). Diborane is a colorless, highly toxic gas with a boiling point of $180K$. It is spontaneously flammable in air and burns in oxygen to release significant energy [7](#page=7): $$B_2H_6 + 3O_2 \rightarrow B_2O_3 + 3H_2O$$; $\Delta H^\ominus = -1976 \text{ KJ mol}^{-1}$ [7](#page=7). Most higher boranes are also spontaneously flammable in air. Diborane is readily hydrolyzed by water to produce boric acid [7](#page=7): $$B_2H_6(g) + 6H_2O(l) \rightarrow 2B(OH)_3(aq) + 6H_2(g)$$ [7](#page=7). Diborane reacts with Lewis bases (L) to form borane adducts ($BH_3 \cdot L$). For example [7](#page=7): $$B_2H_6 + 2NMe_3 \rightarrow 2BH_3 \cdot NMe_3$$ [7](#page=7). $$B_2H_6 + 2CO \rightarrow 2BH_3 \cdot CO$$ [7](#page=7). Reaction with ammonia initially yields $B_2H_6 \cdot 2NH_3$, which can be formulated as $[BH_2(NH_3)_2]^+ [BH_4]^-$. Further heating leads to the formation of borazine ($B_3N_3H_6$), known as "inorganic benzene" due to its ring structure with alternating $BH$ and $NH$ groups [7](#page=7). The structure of diborane features four terminal hydrogen atoms and two boron atoms in one plane, with two bridging hydrogen atoms above and below this plane. The terminal $B-H$ bonds are normal two-center-two-electron bonds, while the bridge bonds ($B-H-B$) are three-center-two-electron bonds, also referred to as banana bonds. Each boron atom in diborane is $sp^3$ hybridized [7](#page=7) [8](#page=8). Boron also forms hydridoborates, with the most important being the tetrahedral $[BH_4]^-$ ion. Lithium and sodium tetrahydridoborates (borohydrides) are synthesized by reacting metal hydrides with diborane in diethyl ether [8](#page=8): $$2MH + B_2H_6 \rightarrow 2 M^+ [BH_4]^-$$ (M = Li or Na) [8](#page=8). Both $LiBH_4$ and $NaBH_4$ are employed as reducing agents in organic synthesis and serve as precursors for other metal borohydrides [8](#page=8). #### 4.1.4 Uses of boron compounds Borax and boric acid are primarily used in the manufacture of heat-resistant glasses (like Pyrex), glass-wool, and fibreglass. Borax also functions as a flux for soldering metals, in heat-resistant glazes for earthenware, and as a component in medicinal soaps. An aqueous solution of orthoboric acid serves as a mild antiseptic [8](#page=8). ### 4.2 Compounds of carbon Carbon forms numerous compounds, with two significant oxides being carbon monoxide ($CO$) and carbon dioxide ($CO_2$) [13](#page=13). #### 4.2.1 Carbon monoxide ($CO$) Carbon monoxide ($CO$) is produced by the direct oxidation of carbon in a limited supply of oxygen or air [13](#page=13): $$2C(s) + O_2(g) \xrightarrow{\text{limited } O_2} 2CO(g)$$ [13](#page=13). Pure $CO$ can be prepared on a small scale by dehydrating formic acid ($HCOOH$) with concentrated sulfuric acid ($H_2SO_4$) at $373K$ [13](#page=13): $$HCOOH \xrightarrow{373K, \text{conc. } H_2SO_4} H_2O + CO$$ [13](#page=13). On a commercial scale, $CO$ is obtained by passing steam over hot coke, producing water gas (synthesis gas) ] [13](#page=13): $$C(s) + H_2O(g) \xrightarrow{473-1273K} CO(g) + H_2(g)$$ [13](#page=13). Water gas Using air instead of steam produces producer gas, a mixture of $CO$ and nitrogen ($N_2$) [13](#page=13): $$2C(s) + O_2(g) + 4N_2(g) \xrightarrow{1273K} 2CO(g) + 4N_2(g)$$ [13](#page=13). Producer gas Both water gas and producer gas are important industrial fuels as they release heat upon further combustion to form carbon dioxide [13](#page=13). Carbon monoxide is a colorless, odorless, and nearly insoluble gas in water. It is a potent reducing agent, capable of reducing most metal oxides except those of alkali and alkaline earth metals, aluminum, and a few transition metals. This property is utilized in the extraction of metals from their oxide ores [13](#page=13) [14](#page=14): $$Fe_2O_3(s) + 3CO(g) \rightarrow 2Fe(s) + 3CO_2(g)$$ [14](#page=14). $$ZnO(s) + CO(g) \rightarrow Zn(s) + CO_2(g)$$ [14](#page=14). The $CO$ molecule contains one sigma and two pi bonds between carbon and oxygen (:C=O:). The lone pair on the carbon atom allows $CO$ to act as a donor ligand, forming metal carbonyls with certain metals upon heating. The toxicity of $CO$ stems from its ability to form a complex with hemoglobin that is approximately 300 times more stable than the oxygen-hemoglobin complex, thus impeding oxygen transport in the blood and potentially leading to death [14](#page=14). #### 4.2.2 Carbon dioxide ($CO_2$) Carbon dioxide ($CO_2$) is formed by the complete combustion of carbon and carbon-containing fuels in excess air [14](#page=14): $$C(s) + O_2(g) \xrightarrow{\text{excess } O_2} CO_2(g)$$ [14](#page=14). $$CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(g)$$ [14](#page=14). In the laboratory, $CO_2$ is conveniently prepared by reacting dilute hydrochloric acid ($HCl$) with calcium carbonate ($CaCO_3$) [14](#page=14): $$CaCO_3(s) + 2HCl(aq) \rightarrow CaCl_2(aq) + CO_2(g) + H_2O(l)$$ [14](#page=14). Commercially, it is obtained by heating limestone [14](#page=14). $CO_2$ is a colorless and odorless gas with low solubility in water, making it significant in biochemical and geochemical processes. It forms carbonic acid ($H_2CO_3$) in water, a weak dibasic acid that dissociates in two steps [14](#page=14): $$H_2CO_3(aq) + H_2O(l) \rightleftharpoons HCO_3^-(aq) + H_3O^+(aq)$$ [14](#page=14). $$HCO_3^-(aq) + H_2O(l) \rightleftharpoons CO_3^{2-}(aq) + H_3O^+(aq)$$ [14](#page=14). The $H_2CO_3/HCO_3^-$ buffer system is crucial for maintaining blood pH between $7.26$ and $7.42$. Being acidic, $CO_2$ reacts with alkalies to form metal carbonates [14](#page=14). Photosynthesis, the process by which green plants convert atmospheric $CO_2$ into carbohydrates like glucose, removes $CO_2$ from the atmosphere [14](#page=14): $$\text{Chlorophyll}, hv$$ $$6CO_2 + 12H_2O \rightarrow C_6H_{12}O_6 + 6O_2 + 6H_2O$$ [14](#page=14). Unlike $CO$, $CO_2$ is not poisonous. However, increased combustion of fossil fuels and limestone decomposition for cement production are raising atmospheric $CO_2$ levels, potentially intensifying the greenhouse effect and global temperatures [14](#page=14). Solid carbon dioxide, known as dry ice, is formed by rapidly expanding liquefied $CO_2$ and is used as a refrigerant for ice cream and frozen foods. Gaseous $CO_2$ is used to carbonate soft drinks and as a fire extinguisher due to its density and non-combustible nature. A substantial amount of $CO_2$ is also used in the manufacture of urea [14](#page=14). In the $CO_2$ molecule, the carbon atom undergoes $sp$ hybridization, forming two sigma bonds with oxygen atoms through the overlap of $sp$ orbitals with oxygen's $p$ orbitals. The remaining two electrons on carbon participate in $p\pi-p\pi$ bonding with oxygen, resulting in a linear shape with equal $C-O$ bond lengths ($115 \text{ pm}$) and no net dipole moment. Its resonance structures are [14](#page=14): $$: \stackrel{-}{O}-C=\stackrel{+}{O}: \leftrightarrow:\stackrel{+}{O}=C=\stackrel{-}{O}: \leftrightarrow:\stackrel{-}{O}=C-\stackrel{+}{O}:$$ [14](#page=14). ### 4.3 Compounds of silicon Silicon forms a vast array of compounds, with silicon dioxide ($SiO_2$) and silicates being particularly abundant, constituting about $95\%$ of the Earth's crust [14](#page=14). #### 4.3.1 Silicon dioxide ($SiO_2$) Silicon dioxide, commonly known as silica, exists in various crystallographic forms, including quartz, cristobalite, and tridymite. Silica is a covalent, three-dimensional network solid where each silicon atom is tetrahedrally bonded to four oxygen atoms, and each oxygen atom is bonded to another silicon atom. The entire crystal can be viewed as a giant molecule with eight-membered rings of alternating silicon and oxygen atoms [14](#page=14) [15](#page=15). Due to the very high $Si-O$ bond enthalpy, silica is generally unreactive, resisting attacks from halogens, hydrogen, most acids, and metals, even at elevated temperatures. However, it reacts with hydrofluoric acid ($HF$) and sodium hydroxide ($NaOH$) ] [15](#page=15): $$SiO_2 + 2NaOH \rightarrow Na_2SiO_3 + H_2O$$ [15](#page=15). $$SiO_2 + 4HF \rightarrow SiF_4 + 2H_2O$$ [15](#page=15). Quartz is utilized for its piezoelectric properties in accurate clocks, radio, television broadcasting, and mobile communications. Silica gel is used as a drying agent and a support for chromatography and catalysts. Kieselghur, an amorphous form of silica, finds use in filtration plants [15](#page=15). #### 4.3.2 Silicones Silicones are organosilicon polymers characterized by the repeating unit $\{R_2SiO\}$. They are synthesized from alkyl or aryl substituted silicon chlorides ($R_nSiCl_{4-n}$). For instance, the reaction of methyl chloride with silicon in the presence of copper catalyst at $570K$ produces various methyl-substituted chlorosilanes, primarily dimethyl dichlorosilane ($(CH_3)_2SiCl_2$) ] [15](#page=15): $$2CH_3Cl + Si \xrightarrow{Cu \text{ powder}, 570K} (CH_3)_2SiCl_2 + \text{other products}$$ [15](#page=15). Hydrolysis of dimethyl dichlorosilane, followed by condensation polymerization, yields straight-chain silicone polymers [15](#page=15): $$(CH_3)_2SiCl_2 + 2H_2O \rightarrow (CH_3)_2Si(OH)_2 + 2HCl$$ [15](#page=15). $$n(CH_3)_2Si(OH)_2 \xrightarrow{\text{Polymerisation, } -H_2O} \left[ \begin{array}{c} CH_3 \\ | \\ -O-Si- \\ | \\ CH_3 \end{array} \right]_n$$ [15](#page=15). Silicone The chain length can be controlled by adding blocking agents like trimethylchlorosilane ($(CH_3)_3SiCl$) [15](#page=15). > **Tip:** Silicones are hydrophobic (water-repellent) due to the non-polar alkyl groups surrounding them. They generally exhibit high thermal stability, high dielectric strength, and resistance to oxidation and chemicals [16](#page=16). Silicones have diverse applications, including sealants, greases, electrical insulators, and waterproofing for fabrics. Their biocompatibility also makes them suitable for surgical and cosmetic implants [16](#page=16). #### 4.3.3 Silicates Silicates are a diverse group of minerals, including feldspar, zeolites, mica, and asbestos. The fundamental structural unit of silicates is the tetrahedron, $SiO_4^{4-}$, where a silicon atom is bonded to four oxygen atoms. These units can be discrete or linked together by sharing oxygen atoms at their corners, forming chains, rings, sheets, or three-dimensional networks. The negative charge of the silicate structures is balanced by positively charged metal ions [16](#page=16). Two important man-made silicates are glass and cement [16](#page=16). #### 4.3.4 Zeolites When aluminum atoms substitute some silicon atoms in the three-dimensional silica network, an aluminosilicate structure with an overall negative charge is formed. Cations like $Na^+$, $K^+$, or $Ca^{2+}$ neutralize this charge. Examples include feldspar and zeolites. Zeolites are widely used as catalysts in the petrochemical industry for cracking and isomerization of hydrocarbons, such as ZSM-5, which converts alcohols into gasoline. Hydrated zeolites are also employed as ion exchangers for softening hard water [16](#page=16). --- # Allotropes of carbon and their properties Carbon exhibits numerous allotropic forms, encompassing both crystalline and amorphous structures, which arise from its unique ability to form chains and rings through catenation and its capacity to form $\pi - \pi$ multiple bonds [11](#page=11). ### 5.1 Crystalline allotropes of carbon #### 5.1.1 Diamond Diamond is characterized by a crystalline lattice where each carbon atom is $sp^3$ hybridized. Each carbon atom is covalently bonded to four other carbon atoms in a tetrahedral arrangement. The C-C bond length in diamond is 154 pm. This structure extends into a rigid, three-dimensional network of carbon atoms with directional covalent bonds throughout the lattice. Due to the immense difficulty in breaking these extensive covalent bonds, diamond is the hardest known substance. Its hardness makes it valuable as an abrasive for sharpening tools, in the manufacture of dyes, and for tungsten filaments in light bulbs [11](#page=11). > **Tip:** The three-dimensional network of strong C-C bonds in diamond is responsible for its exceptionally high melting point, despite being a covalent substance [12](#page=12). * **Uses:** * Abrasive for sharpening hard tools [11](#page=11). * Manufacture of tungsten filaments for electric light bulbs [11](#page=11). * Jewellery [13](#page=13). #### 5.1.2 Graphite Graphite possesses a layered structure, with layers held together by weak van der Waals forces. The distance between these layers is 340 pm. Each layer consists of planar hexagonal rings of carbon atoms, with a C-C bond length of 141.5 pm within the layer. Within each layer, every carbon atom is $sp^2$ hybridized and forms three sigma bonds with its neighboring carbon atoms. The fourth electron from each carbon atom forms a $\pi$ bond, resulting in delocalized electrons across the entire sheet. These mobile electrons enable graphite to conduct electricity along the layers. Due to the weak forces between layers, graphite cleaves easily, making it soft and slippery. This property makes it an excellent dry lubricant for high-temperature machinery where oil is not suitable [12](#page=12). > **Tip:** The conductivity of graphite is attributed to the delocalized electrons in its layered structure [12](#page=12). * **Uses:** * Dry lubricant in machines running at high temperatures [12](#page=12). * Electrodes in batteries and industrial electrolysis due to its conductivity [13](#page=13). * Crucibles made from graphite are inert to dilute acids and alkalies [13](#page=13). * Graphite fibers are used to form high strength, lightweight composites for products like tennis rackets, fishing rods, aircraft, and canoes [13](#page=13). #### 5.1.3 Fullerenes Fullerenes are formed by heating graphite in an electric arc in the presence of inert gases like helium or argon. The resulting sooty material contains mainly C60, along with smaller amounts of C70 and other fullerenes with even numbers of carbon atoms up to 350 or more. Fullerenes are considered pure forms of carbon as they lack "dangling" bonds. They are cage-like molecules. The C60 molecule, known as Buckminsterfullerene, has a shape resembling a soccer ball. It comprises twenty six-membered rings and twelve five-membered rings. A six-membered ring can fuse with either six- or five-membered rings, while a five-membered ring can only fuse with six-membered rings. All carbon atoms in C60 are equivalent and are $sp^2$ hybridized. Each carbon atom forms three sigma bonds with other carbon atoms, and the remaining electron is delocalized in molecular orbitals, imparting aromatic character to the molecule. The C-C distances in C60 are 143.5 pm for single bonds and 138.3 pm for double bonds. Spherical fullerenes are also referred to as bucky balls [12](#page=12) [13](#page=13). > **Example:** The C60 molecule, with its soccer ball-like structure, exemplifies the unique cage-like architecture of fullerenes [12](#page=12). * **Stability and Enthalpy:** Graphite is the most thermodynamically stable allotrope of carbon, with a standard enthalpy of formation ($\Delta_f H^\circ$) of zero. The $\Delta_f H^\circ$ values for diamond and fullerene C60 are 1.90 kJ mol$^{-1}$ and 38.1 kJ mol$^{-1}$, respectively [13](#page=13). ### 5.2 Amorphous allotropes of carbon Other forms of elemental carbon, such as carbon black, coke, and charcoal, are impure forms of graphite or fullerenes [13](#page=13). * **Carbon black:** Obtained by burning hydrocarbons with limited air supply [13](#page=13). * **Charcoal:** Produced by heating wood at high temperatures in the absence of air [13](#page=13). * **Coke:** Produced by heating coal at high temperatures in the absence of air [13](#page=13). > **Tip:** Activated charcoal, a porous form, is effective in adsorbing poisonous gases and is used in water filters and air-conditioning systems for odor control [13](#page=13). * **Uses of Amorphous Forms:** * Carbon black is used as a black pigment in ink and as a filler in automobile tires [13](#page=13). * Coke is used as a fuel and a reducing agent in metallurgy [13](#page=13). * Charcoal (activated) is used for adsorption of poisonous gases, in water filters, and in air-conditioning systems [13](#page=13). ### 5.3 Impact of structure on properties The distinct structures of diamond, graphite, and fullerenes directly influence their physical and chemical properties. Diamond's rigid, three-dimensional covalent network leads to extreme hardness and a high melting point. Graphite's layered structure with delocalized electrons allows for electrical conductivity and makes it soft and slippery, suitable as a lubricant. Fullerenes, with their cage-like structures and delocalized $\pi$ electrons, exhibit unique electronic and chemical properties [11](#page=11) [12](#page=12) [13](#page=13). --- ## 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 | |------|------------| | p-block elements | Elements where the last electron enters the outermost p orbital, comprising groups 13 to 18 of the periodic table. | | Valence shell electronic configuration | The arrangement of electrons in the outermost shell of an atom, specifically the ns np^(1-6) configuration for p-block elements (excluding Helium). | | Oxidation state | A number assigned to an element in a chemical combination which represents the number of electrons lost or gained by an atom of that element in order to form that compound. | | Inert pair effect | The reluctance of the two s-electrons in the valence shell of heavier elements of groups 13-16 to participate in bonding, leading to a preference for oxidation states two units less than the group oxidation state. | | Ionisation enthalpy | The energy required to remove the most loosely bound electron from a neutral gaseous atom in its ground state. | | Electronegativity | A measure of the tendency of an atom to attract a bonding pair of electrons. | | Catenation | The ability of an atom to form long chains or rings by bonding with other atoms of the same element, particularly characteristic of carbon. | | Allotropes | Different structural modifications of an element; the same element existing in two or more different forms in the same physical state, differing in their internal arrangement of atoms. | | Lewis acid | A chemical species that can accept an electron pair to form a covalent bond. | | Lewis base | A chemical species that can donate an electron pair to form a covalent bond. | | Amphoteric oxides | Oxides that can act as either an acid or a base, reacting with both acids and bases. | | Hydrolysis | A chemical reaction in which a molecule of water breaks down into two parts and then combines with a substance, causing the decomposition of the substance. | | Borax | A white crystalline solid, a salt of boric acid and sodium, with the formula Na2B4O7.10H2O, commonly used in detergents, cosmetics, and as a flux. | | Orthoboric acid | A weak monoprotic acid with the formula H3BO3, often referred to as boric acid, known for its antiseptic properties and use in glass manufacturing. | | Diborane | The simplest boron hydride with the chemical formula B2H6, known for its unique three-center two-electron bonding and its reactivity. | | Carbon monoxide | A colorless, odorless, and toxic gas with the formula CO, formed by incomplete combustion of carbon. | | Carbon dioxide | A colorless gas with the formula CO2, essential for photosynthesis and a major greenhouse gas, formed by complete combustion of carbon. | | Silicon dioxide | A compound with the chemical formula SiO2, commonly known as silica, forming the basis of most rocks and sand, and used in glass and ceramics. | | Silicones | Organosilicon polymers characterized by a repeating -[R2SiO]- unit, known for their thermal stability and water-repellent properties. | | Silicates | Minerals containing silicon and oxygen, with the basic structural unit being the tetrahedron SiO4, forming various structures like chains, sheets, and 3D networks. | | Zeolites | Aluminosilicate minerals with a porous structure, widely used as catalysts and ion exchangers. | | Hybridisation | The concept of mixing atomic orbitals into new hybrid orbitals suitable for the pairing of electrons to form chemical bonds in valence bond theory. | | Banana bonds | Three-center two-electron bonds found in molecules like diborane, where a single pair of electrons is delocalized over three atoms. | | Green house effect | The process by which radiation from the sun is trapped by the atmosphere of the Earth, leading to an increase in global temperature. | | Water gas | A mixture of carbon monoxide and hydrogen produced by passing steam over hot coke, used as a fuel and a source of hydrogen. | | Producer gas | A mixture of carbon monoxide, nitrogen, and hydrogen produced by passing air over hot coke, used as a fuel. |