CHM 102: Chemistry of Selected Metals and Non-Metals

General Principles of Periodic Classification and Elemental Properties

• Periods and Groups:     * Horizontal rows in the periodic table are referred to as periods.     * Each period concludes with a noble gas.     * Vertical columns are known as groups.     * Groups are categorized into four distinct blocks: $s$, $p$, $d$, and $f$.

• Electronic Configuration Patterns:     * All elements within a specific group possess the same number of electrons in their outer shell.     * The group number indicates the number of electrons present in the outer shell (e.g., Group 1 contains 1 electron; Group 4 contains 4 electrons).     * Similar outer shell electron counts result in similar chemical characteristics among group members.     * Elements located in the same period share an equal number of electron shells.

• Classification of Elements:     * Metals: Located on the left side of the periodic table. They easily lose valence electrons to achieve an electron configuration resembling the nearest noble gas. Metals form basic oxides and display positive oxidation states. They rarely combine chemically among themselves but can form alloys.     * Non-metals: Located on the right side of the periodic table. Their atoms ionize by gaining electrons to resemble the nearest noble gas configuration. They can combine with themselves, other non-metals, and metals. In binary compounds with metals, they show negative oxidation states. With other non-metals, oxidation states may be positive or negative depending on relative electronegativity.     * Metalloids (Semimetals): Found between metals and non-metals; they display characteristics of both groups.

• Specific Group Characteristics:     * s-block Metals: Found in Groups I and II.     * Reducing Agents: Metals at the bottom of Group I are the most powerful reducing agents.     * Alkali Metals (Group I): React with water to produce strong alkalis.     * Alkaline Earth Metals (Group II): Tend to make weaker alkalis.     * Halogens (Group VII): Show the most complete set of non-metal properties. Fluorine, at the top of the group, is the most powerful oxidizing agent.

The Chemistry of Sodium ($Na$)

• Extraction:     * Sodium is extracted via the electrolysis of a concentrated solution of sodium chloride ($NaCl$/brine).     * Sodium is discharged at the steel cathode.     * Chlorine is released at the graphite anode.

• Physical Properties:     * Appearance: Silvery soft solid with a metallic lustre.     * Melting Point: $98^\circ C$.     * Boiling Point: $883^\circ C$.     * Conductivity: Good conductor of heat and electricity.     * Density: $0.97\,g\,cm^{-3}$.

• Chemical Properties:     * Reaction with Air: Tarnishes rapidly due to oxidation by atmospheric oxygen.     * $4Na_{(s)} + O_{2(g)} \rightarrow 2Na_2O_{(s)}$     * Sodium oxide reacts with water vapor to form sodium hydroxide: $Na_2O_{(s)} + H_2O_{(g)} \rightarrow 2NaOH_{(aq)}$     * Sodium hydroxide absorbs atmospheric $CO_2$ to form hydrated sodium trioxocarbonate IV: $2NaOH_{(aq)} + CO_{2(g)} \rightarrow Na_2CO_{3(s)} + H_2O_{(l)}$     * When heated in plentiful air, it burns with a golden yellow flame to form sodium peroxide: $2Na_{(s)} + O_{2(g)} \rightarrow Na_2O_{2(s)}$     * Storage: Always stored under paraffin oil, toluene, or naphthalene to prevent atmospheric oxidation.     * Reaction with Cold Water: Reacts vigorously, releasing heat, hydrogen, and sodium hydroxide: $2Na_{(s)} + 2H_2O_{(l)} \rightarrow 2NaOH_{(aq)} + H_{2(g)}$     * Reaction with Acids: Reacts explosively with dilute acids (dangerous laboratory reaction): $2Na_{(s)} + 2HCl_{(aq)} \rightarrow 2NaCl_{(aq)} + H_{2(g)}$     * Reaction with Ammonia: Forms sodamide and hydrogen: $2Na_{(s)} + 2NH_{3(g)} \rightarrow 2NaNH_{2(s)} + H_{2(g)}$     * Reaction with Non-metals: Directly combines on heating with most non-metals except Boron, Carbon, and Nitrogen.     * Reaction with Sulphur: $2Na_{(s)} + S_{(s)} \rightarrow Na_2S_{(s)}$     * Reaction with Phosphorus: $3Na_{(s)} + P_{(s)} \rightarrow Na_3P_{(s)}$ (Trisodium Phosphide III).

• Flame Test for Sodium Ions ($Na^+$):     * Sodium compounds impart a golden-yellow color to a non-luminous flame.     * Confirmation: If the golden-yellow color is not visible through a blue glass, the presence of sodium ions is confirmed.

• Uses of Sodium:     * Manufacturing: Sodium peroxide, sodamide, and sodium cyanide.     * Petrol: Production of tetra ethyl lead IV (anti-knock agent).     * Lighting: Sodium vapor lamps (bright-yellow light for highways, airports, and streets).     * Metallurgy: Reducing agent in the extraction of titanium.     * Nuclear: Liquid sodium used as a coolant in reactors.     * Organic Chemistry: Reducing agent (mixtures of sodium/ethanol or sodium amalgam/water).

• Important Sodium Compounds:     * Sodium oxide ($Na_2O$)     * Sodium peroxide ($Na_2O_2$)     * Sodium hydroxide ($NaOH$)     * Sodium chloride ($NaCl$)     * Sodium tetraoxosulphate VI ($Na_2SO_4$)     * Sodium hydrogen trioxocarbonate IV ($NaHCO_3$)

The Chemistry of Aluminium ($Al$)

• Natural Occurrence:     * Most abundant metal and third most plentiful element in the earth's crust.     * Major Ore: Bauxite (contains up to $60\%$ Aluminium oxide, $Al_2O_3$, known as alumina).     * Other Sources: Cryolite, Kaolin, Corundum, and Mica.

• Extraction Process (Two Stages):     * Stage 1: Purification of Bauxite:         * Bauxite is heated with concentrated $NaOH$ under pressure to form sodium aluminate III: $Al_2O_{3(s)} + 2NaOH_{(aq)} + 3H_2O \rightarrow 2NaAl(OH)<em>{4(aq)}$         * Filtration occurs: Residue contains insoluble iron III oxide and trioxosilicates IV; filtrate is sodium aluminate III.         * The filtrate is seeded with pure aluminium hydroxide crystals to induce precipitation: $NaAl(OH)</em>{4(s)} \rightarrow Al(OH)<em>{3(aq)} + NaOH</em>{(aq)}$         * Precipitate is filtered, dried, and heated: $2Al(OH)<em>{3(s)} \rightarrow Al_2O</em>{3(s)} + 3H_2O_{(l)}$     * Stage 2: Electrolysis of Alumina:         * Cathode: Thick lining of graphite.         * Anode: Graphite rod.         * Electrolyte: Solution of pure alumina in molten cryolite ($Na_3AlF_6$).         * Cryolite role: Lowers the melting temperature to approximately $1000^\circ C$.         * Oxygen is released at the anodes; metallic aluminium is deposited at the cathode.

• Physical Properties:     * Appearance: Silvery-white metal.     * Mechanical: Ductile and malleable (foils, wires, sheets).     * Conductivity: Good conductor of heat and electricity.     * Tensile Strength: Moderate.     * Melting Point: $660^\circ C$.     * Boiling Point: $2450^\circ C$.

• Chemical Properties:     * Reaction with Air: Forms a thin, continuous protective layer of $Al_2O_3$, making it corrosion-free. Burns at $800^\circ C$ to form oxide and nitride.     * $4Al_{(s)} + 3O_{2(g)} \rightarrow 2Al_2O_3$     * $2Al + N_{2(g)} \rightarrow 2AlN_{(s)}$     * Reaction with Non-metals: Combines with halogens, sulphur, nitrogen, phosphorus, and carbon on heating.     * $2Al_{(s)} + 3Cl_{2(g)} \rightarrow 2AlCl_{3(s)}$     * Reaction with Acids:         * Reacts slowly with dilute $HCl$, rapidly with concentrated $HCl$: $2Al_{(s)} + 6HCl_{(aq)} \rightarrow 2AlCl_{3(aq)} + 3H_{2(g)}$         * No reaction with dilute $H_2SO_4$; hot concentrated $H_2SO_4$ releases sulphur IV oxide ($Al$ acting as a reducing agent): $2Al_{(s)} + 6H_2SO_4 \rightarrow Al_2(SO_4)<em>{3(aq)} + 3SO</em>{2(g)}$         * No reaction with $HNO_3$ (any concentration) due to a protective oxide layer.     * Reaction with Alkali: Dissolves in $NaOH$ or $KOH$ to form aluminate III and hydrogen: $2Al_{(s)} + 2NaOH_{(aq)} + 6H_2O_{(l)} \rightarrow 2NaAl(OH)<em>{4(aq)} + 3H</em>{2(g)}$     * Thermite Process: Reduces iron III oxide to molten iron: $2Al_{(s)} + Fe_2O_{3(s)} \rightarrow Al_2O_{3(s)} + 2Fe_{(s)}$

• Test for Aluminium Ions ($Al^{3+}$):     * Adding $NaOH$: Formation of white gelatinous precipitate ($Al(OH)_3$) which dissolves in excess $NaOH$ to form $NaAl(OH)_4$.     * Confirmatory test: Adding aqueous ammonia forms a white gelatinous precipitate that is insoluble in excess ammonia.

• Uses of Aluminium:     * Packaging: Aluminium foils.     * Water Treatment: $Al^{3+}$ ions used as coagulating agents.     * Electrical: Overhead cables (light weight and high conductivity).     * Household: Cooking utensils.     * Paints/Mirrors: Aluminium powder in oil (high reflectivity).     * Industry: Thermite process for welding/repairs.     * Alloys: Duralumin ($Al, Cu, Mg, Mn$), aluminium bronze ($Cu, Al$), and magnalium ($Al, Mg$).

• Important Compounds: Aluminium oxide ($Al_2O_3$), Aluminium hydroxide ($Al(OH)_3$), Aluminium chloride ($AlCl_3$), and Aluminium tetraoxosulphate VI.

The Chemistry of Nitrogen ($N$)

• Molecular Structure:     * Exists as a diatomic molecule ($N_2$) with a triple covalent bond ($N \equiv N$).     * The high bond energy of the triple bond makes it extremely stable and unreactive under ordinary conditions.

• Natural Occurrence:     * Most common gas, making up nearly $78\%$ of the atmosphere.     * Role: Dilutes oxygen to slow down combustion and metal oxidation.

• Industrial Preparation:     * Obtained by fractional distillation of liquid air.     * Nitrogen gas evolves first at $-196^\circ C$ at standard pressure.     * Oxygen boils later at $-183^\circ C$.

• Physical Properties:     * Appearance: Colourless, odourless, tasteless gas.     * Density: Slightly less dense than air.     * Melting Point: $-210^\circ C$.     * Boiling Point: $-196^\circ C$.     * Solubility: Only slightly soluble in water.     * Oxidation States: Varies from $+3$ to $+5$.

• Chemical Properties:     * With Non-metals: Combines with hydrogen to produce ammonia: $N_{2(g)} + H_{2(g)} \rightarrow 2NH_{3(g)}$     * With Metals: Forms nitrides at high temperatures (e.g., Magnesium Nitride): $3Mg_{(s)} + N_{2(g)} \rightarrow Mg_3N_{2(s)}$     * Hydrolysis of Nitrides: Nitrides react with warm water to release ammonia gas: $Mg_3N_{2(s)} + 6H_2O_{(l)} \rightarrow 3Mg(OH)<em>{2(s)} + 2NH</em>{3(g)}$

• Uses of Nitrogen:     * Chemicals: Manufacture of ammonia, cyanide, and fertilizers.     * Cryogenics: Liquid nitrogen used as a cooling agent.     * Chromatography: Used as a carrier gas due to its inert nature.

The Chemistry of Chlorine ($Cl$)

• Occurrence: Highly reactive, never found free in nature; occurs mostly as chlorides.

• Laboratory Preparation: Oxidation of concentrated $HCl$ acid using strong oxidizing agents like $MnO_2$ or $KMnO_4$.

• Industrial Preparation:     * Electrolysis of brine or molten chlorides (Sodium, Magnesium, or Calcium).     * Specialized cells used: Castner, Kellner, and Solvay cells.     * Stored as liquid under pressure in steel cylinders.

• Physical Properties:     * Appearance: Greenish-yellow gas with a choking, irritating smell.     * Solubility: Moderately soluble in water.     * Density: Denser than air.     * Safety: Highly poisonous.

• Chemical Properties:     * Displacement: Displaces all halogens except fluorine from their acid/salt solutions: $Cl_{2(g)} + 2HI_{(aq)} \rightarrow 2HCl_{(aq)} + I_{2(g)}$     * Combination: Directly reacts with elements to form chlorides (e.g., $2Na + Cl_2 \rightarrow 2NaCl$).     * Oxidizing Power: Powerful oxidizing agent (removes hydrogen and accepts electrons). Example with iron II salts: $2FeCl_2 + Cl_2 \rightarrow 2FeCl_3$     * Hydrogen Affinity: Strong affinity for hydrogen; extracts it from compounds like hydrocarbons: $C_{10}H_{16(g)} + 8Cl_{2(g)} \rightarrow 10C_{(s)} + 16HCl_{(g)}$     * Bleaching Agent: Bleaches dyes and inks (except carbon-based ones) in the presence of water.

• Uses of Chlorine:     * Textiles: Bleaching agent for cotton, linen, and wood pulp.     * Sanitation: Germicide for sterilization.     * Organic Synthesis: Manufacture of $CHCl_3$, $CCl_4$, PVC, $KClO$, and $HCl$.     * Consumer Products: Used in aerosol propellants.