CHM 101 Dr. Onipede's Inorganic Chemistry 2024 Notes PDF

# CHM 101 Dr. ONIPEDE's Part ## Recommended Texts: - Concise Inorganic Chemistry by J.D. Lee - Inorganic Chemistry by Catherine E. Housecroft and Alan G. Sharpe - Modern Inorganic Chemistry by C. Chamber and A.K. Holliday ## Comparative Chemistry of Group IA, Group IIA, and Group 14 ### Group IA (...

# CHM 101 Dr. ONIPEDE's Part ## Recommended Texts: - Concise Inorganic Chemistry by J.D. Lee - Inorganic Chemistry by Catherine E. Housecroft and Alan G. Sharpe - Modern Inorganic Chemistry by C. Chamber and A.K. Holliday ## Comparative Chemistry of Group IA, Group IIA, and Group 14 ### Group IA (Alkali Metals) - **Hydrogen [H] (1s1)** - **Lithium [Li] (1s22s1)** - **Sodium [Na] (1s22s22p63s1)** - **Potassium [K] (1s22s22p63s23p64s1)** - **Rubidium [Rb] (1s22s22p63s23p63d104s24p65s1)** - **Cesium [Cs] (1s22s22p63s23p63d104s24p64d105s25p66s¹)** - **Francium [Fr] ([Rn] 7s¹)** Rn: 1s22s22p63s23p63d104s24p64d105s25p64f145d106s26p6 **General Properties:** - All are silvery-white metals, except Cesium (Cs), which is golden yellow. - Excellent conductors of electricity. - Soft and highly reactive. - Low melting point and softness due to the elements contributing only one electron to the molecular orbital. - Have one electron in their outer valence shell, forming univalent ionic and colorless compounds. - Their oxides and hydroxides are strong bases. - Their oxosalts are very stable. - They have the largest atomic size in their period. - They have the smallest ionic size in their period, due to losing their valence electron and losing the outermost shell during ionization. - They have the lowest ionization energy in their respective period. **Reactivity:** - They are very reactive and form simple ionic compounds which are soluble in water. - Reactive with air and water, so they are stored under inert solvents (hydrocarbons). - Found combined in nature due to their reactivity. - Form alloys with themselves (e.g., Na/K) and with other metals (e.g., Na/Hg). - Tarnish rapidly in air, forming a layer of oxide, peroxide, and dioxide. - At ambient temperature, all group IA metals adopt a body-centered cubic structure, but lithium forms a hexagonal close-packed structure at low temperatures. **Differences Between Lithium and other Group IA Metals:** - Lithium has a higher melting and boiling point than other group IA elements. - Lithium is harder. - Lithium reacts less readily with oxygen, forming the normal oxide. It forms peroxide only with great difficulty, and higher oxides are unstable. - Lithium hydroxide is less basic than other hydroxides in the group, so many of its salts are less stable. - Only lithium forms nitride (Li3N) in group IA, but other members form oxides. - Only lithium reacts directly with carbon forming carbide, but other group IA members form oxides. - Lithium forms more complexes than other group IA metals. Ammoniated salts like [Li(NH3) 4] I exist as solids. - Li2CO3, Li3PO4, and LiF are insoluble in water, while LiOH has low solubility. Other group IA metal salts are soluble. - Halides and alkyls of lithium are more covalent than their sodium counterparts. - Compounds with lithium are heavier hydrated. ### Group IIA (Alkaline Earth Metals) - **Beryllium [Be] (1s22s2)** - **Magnesium [Mg] (1s22s22p63s2)** - **Calcium [Ca] (1s22s22p63s23p64s2)** - **Strontium [Sr] (1s22s22p63s23p63d104s24p65s2)** - **Barium [Ba] (1s22s22p63s23p63d104s24p64d105s25p66s2)** - **Radium [Ra] ([Rn]7s2)** Rn: 1s22s22p63s23p63d104s24p64d105s25p64f145d106s26p6 **General Properties:** - All are silvery in color, except beryllium (Be) and magnesium (Mg), which are gray. - Highly reactive metals, but less reactive than Group IA counterparts. - They are divalent, forming colorless ionic compounds. - Oxides and hydroxides are less basic than their group IA counterparts - Their oxosalts (carbonates, sulfates, and nitrates) are less stable to heat. - They have smaller atomic and ionic sizes than their group IA counterparts. - They have higher densities than group IA metals. - Group IIA metals have two valence electrons which participate in bonding forming harder and more cohesive substances. - Beryllium has the highest melting and boiling point of the group and shows some notable differences from other members. **Reactivity:** - Be reacts with steam, but Ca, Sr, and Ba react with cold water, forming hydroxides and hydrogen. Mg reacts with hot water, in the same way. - Be(OH) 2 is amphoteric, but other group IIA metal hydroxides are basic. - All group IIA metals react with acids, forming hydrogen gas and liberating heat. - All react with oxygen forming oxides. - All burn in dry air, forming oxides and nitrides. - Thermal decomposition of oxosalts (CO3, NO3, OH, and SO4) also gives oxides. - All form sulfates, but their solubility decreases down the group (Be > Mg >> Ca > Sr > Ba). - All form nitrates by reacting with nitric acid. Be forms normal nitrate and basic nitrate salts. - All react with hydrogen forming hydrides, except Be. - All react with halogens forming halides. - All react with dinitrogen to form ionic nitrides (M3N2), except Be. Be nitride is volatile. - All react with carbon at high temperatures forming carbides. - BeO reacts with water to form methane. Other group IIA metal carbides react with water, forming ethyne. ## Differences Between Lithium and other Group IA Metals - Beryllium and aluminum form carbide which forms methane on reaction with water while other groups IIA carbides liberate ethyne on reaction with water. - Be and aluminum form covalent hydrides, halides and oxides other group IIA analogues are predominantly ionic. - BeCl2 and AlCl3 fume in moist air, reacting to give HCl. - Beryllium oxide (BeO) and aluminum oxide (Al2O3) are amphoteric but magnesium oxide is basic. - Beryllium and aluminum form [Be(OH)4]2- and [Al(OH)4] - respectively in the presence of excess OH- ions but Mg does not react with OH- ions. - Beryllium hydride is electron deficient and polymeric with multicenter bonding like aluminum hydride. - Beryllium form many complexes not typical of group IA and group IIA. - Be and aluminum alike are rendered passive by nitric acid. - The standard electrode potential of Be and Al are - 1.85V and 1.66V respectively which are much closer than for Ca, Sr and Ba (- 2.87V, 2.89V, - 2.90V) respectively. - Be salts are extensively hydrolysed. - Be salts are among the most soluble salts known. ### Group 14 Elements - **Carbon [C] (1s22s22p2)** - **Silicon [Si] (1s22s22p63s23p2)** - **Germanium [Ge] (1s22s22p63s23p64s23d104p2)** - **Tin [Sn] (1s22s22p63s23p63d104s24p65s24d105p2)** - **Lead [Pb] (1s22s22p63s23p63d104s24p64d105s25p64f145d106s26p2)** - **Flerovium [FI] (1s22s22p63s23p63d104s24p64d105s24f145d106s26p65f146d107s27p2)** Rn: 1s22s22p63s23p63d104s24p64d105s25p64f145d106s26p6 **General Properties:** - The group members vary from non-metals (C) to metalloids (Si, Ge) to metals (Sn, Pb). - They are the most abundant elements found on earth, constituting the earth structure and biological systems (proteins, carbohydrates, and fats). - All except lead exhibit allotropy. - Silicon is the second most abundant element by weight in the earth's crust, making up 24%. - The ores include diamond, graphite, fullerene, charcoal, peat, quartz; SiO2, zeolite; Na2(Al2Si3O10) 2H2O, cassiterite; SnO2, teallite; PbSnS2, galena; PbS, anglesite; PbSO4, boulangerite; Pb5Sb4S11, winkle's ore (75% Ag, 15% S, 7% Ge). ## Differences Between Carbon, Silicon, and other Group 14 Elements - The first element (carbon) differs from the rest of the elements because of its smaller size, higher electronegativity, greater tendency to form covalent bonds, and more non-metallic character. - Carbon forms multiple bonds, such as C=C, C=C, and C=O. - Carbon has a great ability to form chains due to its high tendency to form multiple bonds. While other members form chains, they are shorter and less strong. - Carbon and silicon only contain s and p orbitals in their electron configuration. Subsequent members have d electrons, which accounts for the graded trend in properties observed in the group. ## Comparative Relationship Between Boron and Silicon: - Boron and silicon form acidic oxides (B2O3 and SiO2). - Boron and silicon form many polymeric oxide structures. - Boron and silicon form flammable gaseous hydrides. ## Reactions of Group 14 Elements - All form tetravalent hydrides, with carbon and silicon forming a series of catenated molecule hydrides. - All form tetrahalides with halogens but lead forms dihalides which are stable. - Silicon and germanium can form pentahalides with fluorine. - Only dihalide of lead is stable. - Tin tetrahalides can form pentahalides and hexahalides in an acidic solution. - All react with oxygen to produce oxides. - All react with nitrogen to form stable compounds. - Carbon forms stable sulfides with sulfur. ## Lewis Octet Rule - Atoms in a molecule tend to bond in such a way that each atom has eight electrons in its valence shell. - **Steps to predict the structures of atoms in a molecule.** 1. Place the atom with the lower group number or, in case of the same group number, the higher period number in the center. Hydrogen can form only one bond, so it's never a central atom. 2. Determine the total number of valence electrons available for all atoms. For polyatomic ions, add one electron for each negative charge or subtract one for each positive charge. 3. Draw a single bond from each surrounding atom to the central atom and subtract two valence electrons for each bond. 4. Distribute the remaining electrons so that each atom has eight electrons. - **The VSEPR Theory** determines the shape of molecules by assuming that electrons in a valence shell are arranged as far away as possible from each other to minimize repulsions. This arrangement minimizes electron-electron repulsions and predicts the shape of molecules. - **Rule 1:** Regions of high electron concentration (bonds and lone pairs) repel one another. - **Rule 2:** There is no distinction between single and multiple bonds. - **Rule 3:** All regions of high electron density are included in a description of the electronic arrangement, but only the positions of atoms are considered when reporting the shape of a molecule. - **Rule 4:** The strength of repulsions are in the order: lone pair-lone pair > lone pair - atom > atom-atom. - **Molecular Geometry:** The shape of a molecule is based on its electron arrangement and the VSEPR rules. Lone pairs significantly influence the shape of a molecule, making bond angles smaller.