A Level Inorganic Chemistry Study Guide: Period 3, Group 2, Group 17, and Nitrogen & Sulfur

Chemical Periodicity of Period 3 Elements

• Atomic Arrangement and Periodicity: Elements in the periodic table are organized by increasing atomic number in horizontal rows called periods and vertical columns called groups. Periodicity refers to the repeating patterns in chemical and physical properties observed across these periods.

• Atomic Radius Definition: The distance between the nucleus and the outermost electron. It is measured by taking two atoms of the same element, measuring the distance between their nuclei, and halving it ($\text{distance} / 2$). In metals, this is designated as the metallic radius; in non-metals, it is the covalent radius.

• Trend in Atomic Radius: Across Period 3, the atomic radius decreases.     * Reasoning: As you move from left to right, the number of protons (nuclear charge) and electrons increases by one.     * Shielding: Since the number of principal quantum shells remains the same, the shielding effect is constant.     * Attraction: The increasing nuclear charge pulls the electrons more strongly toward the nucleus, reducing the size of the atom.

• Ionic Radius:     * Cations ($Na^{+}$ to $Si^{4+}$ ): Positive ions are smaller than their parent atoms because they lose their valence electrons and have less shielding. The radius decreases across the metal ions due to increasing nuclear charge affecting the remaining electrons in the second principal quantum shell.     * Anions ($P^{3-}$ to $Cl^{-}$ ): Negative ions are larger than their parent atoms because they gain electrons in the third principal quantum shell. This increases electron-electron repulsion while the nuclear charge remains constant, causing the electron cloud to spread out.     * Trend: Across the anions ($P^{3-}$ to $Cl^{-}$ ), the ionic radius decreases because the nuclear charge increases and fewer electrons are gained ($P$ gains 3, $S$ gains 2, $Cl$ gains 1).

Structure and Bonding Across Period 3

• Metallic Elements ($Na$, $Mg$, $Al$):     * These form giant metallic lattices with positive ions held by a 'sea' of delocalised electrons.     * Bonding Strength: Metallic bonding strengthens from $Na$ to $Al$. $Na$ donates $1$ electron, $Mg$ donates $2$, and $Al$ donates $3$ to the delocalised pool.     * Electrostatic Forces: The forces between a $3+$ ion ($Al^{3+}$) and more delocalised electrons are significantly stronger than those between a $1+$ ion ($Na^{+}$) and fewer electrons. This result is a higher melting point for aluminium compared to sodium.

• Giant Molecular Structure ($Si$): Silicon has the highest melting point in Period 3 ($\approx 1410^{\circ}\text{C}$ or $1610^{\circ}\text{C}$ for its oxide $SiO_2$ in specific contexts). Each silicon atom is held in a network of strong covalent bonds.

• Simple Molecular Structures ($P_4$, $S_8$, $Cl_2$, $Ar$): These are non-metals with strong intramolecular covalent bonds but weak intermolecular instantaneous dipole-induced dipole (van der Waals) forces.     * Melting Point Trend: Melting points decrease from phosphorus to argon. However, sulfur ($S_8$) has a higher melting point than phosphorus ($P_4$) because sulfur exists as a larger molecule with stronger van der Waals forces.

• Electrical Conductivity:     * Trend: Increases from $Na$ to $Al$ due to the increasing number of delocalised valence electrons available for charge carriage.     * Data Values: $Na = 0.213$, $Mg = 0.224$, $Al = 0.382$, $Si = 2 \times 10^{-10}$.     * Non-conductors: $Si$ is a metalloid (semimetal) and lacks sufficient free electrons. $P$, $S$, $Cl$, and $Ar$ have no delocalised electrons and cannot conduct electricity.

Period 3 Oxides and Hydroxides: Acid-Base Behavior

• Oxidation States: Oxygen is more electronegative than Period 3 elements. In oxides, oxygen has an oxidation state of $-2$. Period 3 elements exhibit positive oxidation states:     * $Na_2O$ ($+1$)     * $MgO$ ($+2$)     * $Al_2O_3$ ($+3$)     * $SiO_2$ ($+4$)     * $P_4O_{10}$ ($+5$)     * $SO_2$ ($+4$) and $SO_3$ ($+6$).

• Reactions with Water:     * $Na_2O$: Reacts vigorously, forming $NaOH$ ($pH \text{ 14}$, strongly alkaline).     * $MgO$: Reacts slowly, forming $Mg(OH)_2$ ($pH \text{ 11}$, weakly alkaline due to low solubility).     * $Al_2O_3$ and $SiO_2$: Insoluble in water. $Al_2O_3$ protects the metal beneath it from further corrosion.     * $P_4O_{10}$: Reacts to form phosphoric acid ($H_3PO_4$, acidic).     * $SO_2$ / $SO_3$: Reacts to form sulfurous acid ($H_2SO_3$) or sulfuric acid ($H_2SO_4$), creating highly acidic solutions.

• Acid-Base Nature:     * Basic Oxides: $Na_2O$ and $MgO$. They contains ionic bonds. The $O^{2-}$ ions react with water: $O^{2-}(aq) + H_2O(l) \rightarrow 2OH^{-}(aq)$.     * Amphoteric Oxide: $Al_2O_3$. It reacts with both acids (e.g., $HCl$) and bases (e.g., $NaOH$).     * Acidic Oxides: $SiO_2$, $P_4O_{10}$, $SO_2$, $SO_3$. These are covalently bonded. They react with bases to form salts.

• Hydroxide Properties:     * $NaOH$: Strong base.     * $Mg(OH)_2$: Basic, used in indigestion remedies to neutralize stomach acid: $Mg(OH)_2(s) + 2HCl(aq) \rightarrow MgCl_2(aq) + 2H_2O(l)$.     * $Al(OH)_3$: Amphoteric. Reacts with hydrochloric acid ($AlCl_3$) and sodium hydroxide ($NaAl(OH)_4$).

Period 3 Chlorides and Their Reactions with Water

• Bonding and Structure:     * $NaCl$, $MgCl_2$: Giant ionic structures. They dissolve in water to form neutral ($pH \text{ 7}$) or slightly acidic ($pH \text{ 6.5}$) solutions as ions become hydrated.     * $AlCl_3$: Exists as a giant ionic lattice or a covalent dimer $Al_2Cl_6$. It hydrolyses in water. The highly charged $Al^{3+}$ ion polarizes water molecules, splitting an $H-O$ bond to release $H^{+}$ ions, making the solution acidic and producing white fumes of $HCl$ gas.     * $SiCl_4$ and $PCl_5$: Simple molecular structures with covalent bonds. Both undergo rapid hydrolysis.         * $SiCl_4(l) + 2H_2O(l) \rightarrow SiO_2(s) + 4HCl(g)$         * $PCl_5(s) + 4H_2O(l) \rightarrow H_3PO_4(aq) + 5HCl(g)$     * Solutions from hydrolysis of non-metal chlorides have very low pH (approx $pH \text{ 1-2}$).

Group 2 Elements: Reactions and Thermal Decomposition

• Reactions of Group 2 Metals:     * With Oxygen: $2M(s) + O_2(g) \rightarrow 2MO(s)$.     * With Water: $M(s) + 2H_2O(l) \rightarrow M(OH)_2(s) + H_2(g)$. Beryllium ($Be$) does not react with water. Magnesium ($Mg$) reacts very slowly with cold water but vigorously with steam: $Mg(s) + H_2O(g) \rightarrow MgO(s) + H_2(g)$. Reactivity increases down the group.     * With Dilute Acid: $M(s) + 2HCl(aq) \rightarrow MCl_2(aq) + H_2(g)$. Reaction with sulfuric acid ($H_2SO_4$) forms sulfates ($MSO_4$). Note that $SrSO_4$ and $BaSO_4$ are insoluble.

• Thermal Decomposition:     * Carbonates: $XCO_3(s) \xrightarrow{\text{HEAT}} XO(s) + CO_2(g)$.     * Nitrates: $2X(NO_3)_2(s) \xrightarrow{\text{HEAT}} 2XO(s) + 4NO_2(g) + O_2(g)$. Brown fumes ($NO_2$) are observed.     * Trend in Stability: Thermal stability increases down the group. Smaller ions at the top ($Mg^{2+}$) have higher charge density and polarising power, distorting the carbonate/nitrate anions and weakening the $C-O$ or $N-O$ bonds, making them easier to decompose.

Physical and Chemical Trends of Group 2

• Atomic Radius: Increases down the group as extra principal quantum shells are added.

• Ionization Energy: Both first and second ionization energies decrease down the group. Increased shielding and distance from the nucleus outweigh the increase in nuclear charge, making it easier to lose electrons (hence reactivity increases).

• Melting Point: Generally decreases down the group. The distance between the positive nuclei and the sea of delocalised electrons increases, weakening the metallic bond.

• Solubility Trends:     * Hydroxides: Solubility increases down the group ($Mg(OH)_2$ is least soluble; $Ba(OH)_2$ is most soluble). Consequently, solutions become more alkaline down the group.     * Sulfates: Solubility decreases down the group ($MgSO_4$ is soluble; $BaSO_4$ is an insoluble white precipitate).

Group 17: Physical and Chemical Properties of Halogens

• Physical Properties:     * Colours: Get darker down the group (Fluorine: pale yellow gas; Chlorine: green gas; Bromine: red-brown liquid; Iodine: shiny grey-black solid/purple vapour).     * Volatility: Decreases down the group. Boiling points increase because the increasing number of electrons lead to stronger instantaneous dipole-induced dipole (van der Waals) forces.

• Bond Strength: Bond enthalpy generally decreases down the group as atomic size increases (weaker attraction between nuclei and the bonding pair). Exception: Fluorine ($F_2$) has a lower bond enthalpy than chlorine because the atoms are so small that the lone pairs experience significant repulsion.

• Oxidising Power: Halogens act as oxidising agents by gaining an electron: $X_2 + 2e^{-} \rightarrow 2X^{-}$.     * Trend: Oxidising power decreases down the group as electronegativity decreases and shielding increases.     * Displacement: A more reactive halogen (e.g., $Cl_2$) will displace a less reactive halide (e.g., $Br^{-}$) from solution: $Cl_2(aq) + 2Br^{-}(aq) \rightarrow 2Cl^{-}(aq) + Br_2(aq)$.

Reactions of Halide Ions and Hydrogen Halides

• Reducing Power of Halide Ions: Increases down the group. Larger ions ($I^{-}$) lose electrons more easily because the outer electrons are further from the nucleus and more shielded.

• Reaction with Concentrated Sulfuric Acid ($H_2SO_4$):     * Chlorides: Only displacement occurs: $NaCl(s) + H_2SO_4(l) \rightarrow NaHSO_4(s) + HCl(g)$ (white fumes).     * Bromides: Displacement followed by redox. $H_2SO_4$ oxidises $HBr$ to $Br_2$ (red-brown gas) and is reduced to $SO_2$.     * Iodides: Strongest reduction. $H_2SO_4$ is reduced to $SO_2$, $S$ (yellow solid), or $H_2S$ (bad egg smell). $HI$ is oxidised to $I_2$ (purple vapour).

• Silver Nitrate Test:     * $Ag^{+}(aq) + X^{-}(aq) \rightarrow AgX(s)$.     * $AgCl$: White precipitate; dissolves in dilute ammonia.     * $AgBr$: Cream precipitate; dissolves in concentrated ammonia.     * $AgI$: Yellow precipitate; insoluble in ammonia.

• Chlorine Reactions:     * Cold Alkali ($15^{\circ}\text{C}$): Disproportionation to chloride ($Cl^{-}$) and chlorate(I) ($ClO^{-}$).     * Hot Alkali ($70^{\circ}\text{C}$): Disproportionation to chloride ($Cl^{-}$) and chlorate(V) ($ClO_3^{-}$).     * Water Purification: $Cl_2 + H_2O \rightleftharpoons HCl + HClO$. Chloric(I) acid ($HClO$) and its ion ($ClO^{-}$) act as sterilising agents to kill bacteria.

The Chemistry of Nitrogen and Its Compounds

• Nitrogen Gas ($N_2$): Makes up $78\%$ of the atmosphere. It is very unreactive due to its strong triple covalent bond ($N \equiv N$) with a bond enthalpy of $1000\,kJ\,mol^{-1}$ and its lack of polarity.

• Ammonia ($NH_3$):     * Preparation: Produced via the Haber Process. In the lab, by heating an ammonium salt with a base: $NH_4Cl + Ca(OH)_2 \xrightarrow{\text{HEAT}} CaCl_2 + 2H_2O + 2NH_3$.     * Basicity: Acts as a weak Brønsted–Lowry base. It uses its lone pair to accept a proton ($H^{+}$).     * Structure: Ammonia is pyramidal. After accepting a proton via a dative covalent bond, the ammonium ion ($NH_4^{+}$) is formed, which is tetrahedral.

Nitrogen Oxides: Pollution, Smog, and Acid Rain

• Formation: Nitrogen oxides ($NO$, $NO_2$) form during lightning or in high-pressure/temperature environments like car engines.

• Catalytic Converters: Reduce pollutants using a catalyst (e.g., platinum) to convert $NO$ and $CO$ into harmless gases: $2CO(g) + 2NO(g) \rightarrow 2CO_2(g) + N_2(g)$.

• Photochemical Smog: Nitrogen oxides (primary pollutants) react with Volatile Organic Compounds (VOCs) in the presence of sunlight (photochemical reaction) to form peroxyacetyl nitrate (PAN, $CH_3CO_2NO_2$), a secondary pollutant that affects lungs and eyes.

• Acid Rain:     * Nitrogen(IV) oxide dissolves in water to form nitric acid: $2NO_2 + H_2O + \frac{1}{2}O_2 \rightarrow 2HNO_3$.     * Catalytic Role: $NO_2$ also catalyses the oxidation of sulfur dioxide ($SO_2$) into sulfur trioxide ($SO_3$), which then forms sulfuric acid ($H_2SO_4$).     * Step 1: $NO_2(g) + SO_2(g) \rightarrow SO_3(g) + NO(g)$.     * Step 2: $NO(g) + \frac{1}{2}O_2(g) \rightarrow NO_2(g)$.