Comprehensive Chemistry Study Notes for Senior High School - Year 2

CHEMISTRY FOR SENIOR HIGH SCHOOLS: ENERGY CHANGES

• Contributors and Context: This material is produced by the Ministry of Education, Ghana, and the Ghana Association of Science Teachers (GAST). Authors include Alhassan Kuvidana Mustapha, Christian Dzikunu, Delali Robert Akplai, and Yirwelle Santus.

• Key Definitional Concepts:     - Atomisation energy: The energy required to break all the bonds in one mole of a substance to form individual atoms in the gas phase under standard conditions.     - Bond dissociation energy: The energy required to break a specific chemical bond within a compound.     - Bond enthalpy: The average energy needed to split one mole of a specific type of covalent bond in a gaseous molecule.     - Enthalpy (H): A thermodynamic quantity describing the total energy of a system. It is defined mathematically as:     $H = U + PV$     - Enthalpy change (\Delta H): Compare the enthalpy of products to reactants:     $\Delta H = H_{\text{products}} - H_{\text{reactants}}$     - Hess’s Law of constant heat summation: The total enthalpy change of a chemical reaction is equal to the sum of all individual enthalpy changes.     - Lattice energy: The energy required to separate one mole of an ionic solid into gaseous ions.

• Thermodynamic Systems:     1. Open system: Can exchange both heat and matter with surroundings (e.g., water in an open beaker).     2. Closed system: Can exchange heat but not matter (e.g., water in a covered beaker).     3. Isolated system: Cannot exchange heat or matter (e.g., a well-insulated thermos/calorimeter).

• Exothermic and Endothermic Reactions:     1. Exothermic: Release heat; surroundings get warmer. $\Delta H_{\text{rxn}} < 0$ (negative). Example: Combustion of methane:     $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l) \quad \Delta H = -890.4\,kJ\,mol^{-1}$     2. Endothermic: Absorb heat; surroundings get cooler. $\Delta H_{\text{rxn}} > 0$ (positive). Example: Decomposition of calcium carbonate:     $CaCO_3(s) \rightarrow CaO(s) + CO_2(g) \quad \Delta H = +177.8\,kJ\,mol^{-1}$

• Standard Enthalpy Changes (\Delta H^{\circ}):     - Standard Conditions: Pressure at $1\,atm$ ($101.3\,kPa$) and Temperature at $298\,K$ ($25\,^{\circ}C$).     - Standard state: Most stable form of an element/compound at room temperature and $1\,atm$.     - Standard enthalpy of reaction: $\Delta H_{\text{rxn}}^{\circ} = \sum n\Delta H_f^{\circ}(\text{products}) - \sum m\Delta H_f^{\circ}(\text{reactants})$     - Standard enthalpy of formation (\Delta H_f^{\circ}): Enthalpy change when one mole of a compound is formed from its elements in standard states. Elements in their standard states (e.g., $O_2(g)$, $C(\text{graphite})$) have $\Delta H_f^{\circ} = 0$.     - Standard enthalpy of combustion (\Delta H_c^{\circ}): Enthalpy change when one mole of a substance is burned completely in oxygen. Always negative.     - Standard enthalpy of neutralisation (\Delta H_n^{\circ}): Enthalpy change when acid and alkali react to produce $1\,mol$ of water. For strong acids/bases, this is approximately $-57.4\,kJ\,mol^{-1}$.     - Standard enthalpy of solution (\Delta H_{\text{soln}}^{\circ}): Enthalpy change when one mole of ionic substance dissolves in excess water.     - Standard enthalpy of hydration (\Delta H_{\text{hyd}}^{\circ}): Energy evolved when one mole of gaseous ions is surrounded and stabilized by water molecules.

• Experimental Determination:     - The formula used is:     $Q = m_{\text{water}} \times c \times \Delta T$     - Where $c = 4.18\,J\,g^{-1}\,^{\circ}C^{-1}$ (specific heat capacity of water).     - Calorimetry Procedure: Measure initial/final mass of fuel and initial/final temperature of water.

• Born-Haber Cycle Example (LiF):     1. Sublimation: $Li(s) \rightarrow Li(g) \quad \Delta H^{\circ} = +155.2\,kJ\,mol^{-1}$     2. Dissociation: $\frac{1}{2}F_2(g) \rightarrow F(g) \quad \Delta H^{\circ} = +75.3\,kJ\,mol^{-1}$     3. Ionization: $Li(g) \rightarrow Li^{+}(g) + e^{-} \quad \Delta H^{\circ} = +520\,kJ\,mol^{-1}$     4. Electron Affinity: $F(g) + e^{-} \rightarrow F^{-}(g) \quad \Delta H^{\circ} = -328\,kJ\,mol^{-1}$     5. Lattice Energy: $Li^{+}(g) + F^{-}(g) \rightarrow LiF(s) \quad \Delta H^{\circ} = -1016.6\,kJ\,mol^{-1}$

• Bond Energy Calculation:     - $\Delta H_{\text{rxn}}^{\circ} = \sum \Delta H^{\circ}(\text{bonds broken}) - \sum \Delta H^{\circ}(\text{bonds formed})$

CHEMICAL KINETICS

• Reaction Rate Definition: The change in concentration (moles/mass) of a reactant or product per unit time.     $\text{Rate} = \frac{\Delta [A]}{\Delta t}$     - For a reaction $aA + bB \rightarrow cC + dD$:     $\text{Rate} = -\frac{1}{a} \frac{\Delta [A]}{\Delta t} = -\frac{1}{b} \frac{\Delta [B]}{\Delta t} = \frac{1}{c} \frac{\Delta [C]}{\Delta t} = \frac{1}{d} \frac{\Delta [D]}{\Delta t}$
• Types of Rates:     - Initial rate: Rate at beginning ($t = 0$).     - Average rate: Speed over a specific time interval ($\frac{C_2 - C_1}{t_2 - t_1}$).     - Instantaneous rate: Slope of the tangent to the concentration-time curve at a specific point.
• Factors Affecting Reaction Rates:     1. Temperature: Increases kinetic energy and frequency of effective collisions.     2. Concentration: More particles per volume leads to higher collision probability.     3. Surface Area: Larger area for solids (e.g., powders vs. blocks) increases rate.     4. Catalyst: Lowers activation energy ($E_a$) by providing an alternative pathway.     5. Pressure: Relevant for gases; increases collision frequency.
• Collision Theory: Requires particles to collide with sufficient energy ($\ge E_a$) and correct orientation.
• Maxwell-Boltzmann Distribution: Shows the distribution of kinetic energies. At higher temperatures, the curve flattens and shifts right, increasing the fraction of molecules exceeding $E_a$.
• Rate Equations and Order:     - $\text{Rate} = k[A]^x[B]^y$     - Zero Order: Rate is constant ($\text{Rate} = k$); independent of concentration.     - First Order: Doubling concentration doubles rate ($\text{Rate} = k[A]^1$).     - Second Order: Doubling concentration quadruples rate ($\text{Rate} = k[A]^2$).
• Half-Life ($t_{1/2}$):     - For First Order: $t_{1/2} = \frac{0.693}{k}$     - For Second Order: $t_{1/2} = \frac{1}{k[A]_0}$
• Rate-Determining Step (RDS): The slowest step in a reaction mechanism that controls the overall speed.

DYNAMIC EQUILIBRIUM

• Reversible Reactions (\rightleftharpoons): Can proceed in both forward and backward directions.
• Dynamic Equilibrium Characteristics:     - Occurs in a closed system.     - Rate of forward reaction = Rate of reverse reaction.     - Concentrations of reactants and products remain constant.
• Equilibrium Constants:     - Concentration ($K_c$) for $aA + bB \rightleftharpoons cC + dD$: $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$     - Pressure ($K_p$) for gases: $K_p = \frac{P_C^c P_D^d}{P_A^a P_B^b}$     - Relationship: $K_p = K_c(RT)^{\Delta n}$     - Where $\Delta n = (\text{sum of gas moles of products}) - (\text{sum of gas moles of reactants})$.
• Phase Equilibria:     - Homogeneous: All species in the same phase.     - Heterogeneous: Species in different phases. Note: Pure solids and liquids are excluded from equilibrium constant expressions.
• Solubility Product ($K_{sp}$): For $\text{A}_x\text{B}_y(s) \rightleftharpoons x\text{A}^{y+}(aq) + y\text{B}^{x-}(aq)$:     $K_{sp} = [\text{A}^{y+}]^x [\text{B}^{x-}]^y$
• Le Chatelier’s Principle: If a system at equilibrium is disturbed, the system shifts to counteract the disturbance.     1. Concentration: Add reactant \rightarrow Shift right; Remove product \rightarrow Shift right.     2. Pressure: Increase pressure \rightarrow Shift toward side with fewer gas molecules.     3. Temperature: Increase temp \rightarrow Shift in endothermic direction.     4. Catalyst: No effect on equilibrium position; only reaches equilibrium faster.
• Industrial Applications:     - Haber Process: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad \Delta H = -92\,kJ\,mol^{-1}$ (Optimized by high pressure, moderate temp, Iron catalyst).     - Contact Process: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g) \quad \Delta H = -196\,kJ\,mol^{-1}$ (Optimized by $V_2O_5$ catalyst).

ACIDS, BASES AND SALTS

• Acid-Base Theories:     1. Arrhenius: Acids produce $H^{+}$ in water; Bases produce $OH^{-}$ in water.     2. Br\o nsted-Lowry: Acids are proton ($H^{+}$) donors; Bases are proton acceptors.     3. Lewis: Acids are electron-pair acceptors; Bases are electron-pair donors.
• Conjugate Pairs: Differ by a single proton ($H^{+}$). Example: $H_2O$ (base) and $H_3O^{+}$ (conjugate acid).
• Physical Properties:     - Acids: Sour taste, turn blue litmus red, $pH < 7$, conduct electricity.     - Bases: Bitter taste, slippery feel, turn red litmus blue, $pH > 7$.
• Neutralisation Applications:     - Antacids: Bases like $Mg(OH)_2$ or $Al(OH)_3$ neutralize stomach $HCl$.     - Agriculture: Calcium carbonate ($CaCO_3$) reduces soil acidity.
• Types of Salts:     - Normal: All acidic hydrogens replaced (e.g., $NaCl$).     - Acidic: Partial replacement of hydrogens (e.g., $NaHSO_4$).     - Double: Two different cations (e.g., potash alum).     - Deliquescent: Absorb enough water to dissolve (e.g., $NaOH$, $CaCl_2$).     - Efflorescent: Lose water of crystallization to air (e.g., $Na_2CO_3 \cdot 10H_2O$).
• Titration Methods:     - Back Titration: React analyte with excess known reagent, then titrate the remainder.         - Aspirin Calculation: Two crushed tablets mixed with $25\,cm^3$ of $1.00\,mol\,dm^{-3}\,NaOH$. Excess titrated with $0.050\,mol\,dm^{-3}\,H_2SO_4$. Result: $1.33\,g$ of acetylsalicylic acid per tablet.     - Double-Indicator Titration: Uses phenolphthalein and methyl orange to analyze mixtures like $Na_2CO_3 + NaHCO_3$.

TRENDS IN PERIOD 3 ELEMENTS

• Elements: $Na$, $Mg$, $Al$, $Si$, $P$, $S$, $Cl$, $Ar$.
• Physical Trends:     - Metallic Character: Decreases left to right.     - Melting/Boiling Points: Increase from $Na$ to $Si$ ($Si$ is highest due to giant covalent structure), then decrease ($S_8 > P_4 > Cl_2 > Ar$ due to molecular size/Van der Waals forces).     - Electrical Conductivity: Highest for metals ($Al > Mg > Na$), semiconductor ($Si$), non-metals are insulators.
• Chemical Reactions:     - With Water: $Na$ reacts violently; $Mg$ reacts slowly with cold/vigorously with steam.     - Oxides:         - Basic: $Na_2O$, $MgO$.         - Amphoteric: $Al_2O_3$.         - Acidic: $SiO_2$, $P_4O_{10}$, $SO_2$.     - Chlorides:         - Ionic ($NaCl$, $MgCl_2$): Neutral $pH$ in water.         - Covalent ($AlCl_3$, $SiCl_4$, $PCl_5$): Hydrolyse in water to produce acidic solutions (release $HCl$ gas).

PROPERTIES OF THE HALOGENS (GROUP 17)

• Physical States: $F_2$ (Pale yellow gas), $Cl_2$ (Yellow-green gas), $Br_2$ (Reddish-brown liquid), $I_2$ (Shiny purple-black solid).
• Trends:     - Boiling Point: Increases down the group as Van der Waals forces strengthen.     - Oxidising Strength: Decreases down the group ($F_2 > Cl_2 > Br_2 > I_2$).     - Electronegativity: Decreases down the group.
• Hydrogen Halides (HX):     - Acid Strength: Increases down the group ($HF \ll HCl < HBr < HI$). $HF$ is weak due to high bond enthalpy ($565\,kJ\,mol^{-1}$).     - Thermal Stability: Decreases down the group ($HF$ is most stable; $HI$ decomposes easily on heating).
• Uses:     - Chlorination: Water purification ($Cl_2 + H_2O \rightarrow HCl + HClO$).

STRUCTURE, BONDING AND MOLECULAR GEOMETRY

• Electronegativity: Pauling scale (0.7 to 4.0). Fluorine is highest (4.0).
• Bond Types:     - Non-polar Covalent: Difference $< 0.5$.     - Polar Covalent: Difference $0.5$ to $1.7$.     - Ionic: Difference $> 1.7$.
• VSEPR Theory Predicted Shapes:     - Linear: $CO_2$ ($180^{\circ}$, 2 charge centers).     - Trigonal Planar: $BF_3$ ($120^{\circ}$, 3 charge centers).     - Tetrahedral: $CH_4$ ($109.5^{\circ}$, 4 charge centers).     - Trigonal Pyramidal: $NH_3$ ($107^{\circ}$ due to 1 lone pair).     - Bent: $H_2O$ ($104.5^{\circ}$ due to 2 lone pairs).
• Sigma (\sigma) and Pi (\pi) Bonds:     - Sigma (\sigma): Head-on overlap; found in all single bonds; allows free rotation.     - Pi (\pi): Sideways overlap of p-orbitals; found in double/triple bonds; restricts rotation.
• Hybridisation:     - $sp^3$: 4 equivalent orbitals (Tetrahedral, e.g., Methane).     - $sp^2$: 3 equivalent orbitals + 1 unhybridised p-orbital (Trigonal planar, e.g., Ethene).     - $sp$: 2 equivalent orbitals + 2 unhybridised p-orbitals (Linear, e.g., Ethyne).

ORGANIC COMPOUNDS

• Alkanes ($C_n H_{2n+2}$):     - Saturated, unreactive except for Halogenation (requires UV light) and Combustion.     - Cracking: Breaking long chains into shorter alkanes and alkenes.
• Alkenes ($C_n H_{2n}$):     - Unsaturated (double bond). Suffix "-ene".     - Addition Reactions: Hydrogenation (Nickel catalyst, $130\,^{\circ}C$), Halogenation, Hydration (Steam + acid catalyst to make alcohols).     - Test: Bromine water turns from reddish-brown to colourless.
• Alkynes ($C_n H_{2n-2}$):     - Triple bond. Example: Ethyne ($C_2H_2$). Prepared from Calcium Carbide ($CaC_2 + 2H_2O \rightarrow C_2H_2 + Ca(OH)_2$).
• Benzene ($C_6H_6$):     - Hexagonal ring with delocalized \pi-electrons. Unique stability (Resonance).     - Reactions: Prefers Electrophilic Substitution (Nitration, Halogenation, Friedel-Crafts) over addition to maintain aromaticity.
• Alkanols ($C_n H_{2n+1}OH$):     - Primary (1^\circ): carbon attached to 1 other carbon.     - Secondary (2^\circ): carbon attached to 2 others.     - Tertiary (3^\circ): carbon attached to 3 others.     - Lucas Test: Tertiary alcohols react immediately (cloudiness); Secondary in 5-10 mins; Primary no reaction at room temp.
• Alkanoic Acids ($C_n H_{2n+1}COOH$):     - Contain carboxyl group. Esterification: Alkanoic acid + Alkanol \rightarrow Ester + Water (requires conc. $H_2SO_4$).

QUESTIONS & DISCUSSION

• Energy Changes Activities: Why stir the water in a calorimeter? (To ensure even heat distribution). What causes differences from theoretical values? (Heat loss to surroundings).
• Kinetics Activities: How does surface area affect grain mill safety? (Dust can cause explosions due to high reactivity). Why is smoking forbidden near oxygen bottles? (Oxygen increases the rate of combustion reactions).
• Equilibrium Discussion: Why does $K_{sp}$ remain constant even if solubility changes? ($K_{sp}$ is temperature-dependent, not concentration-dependent).
• Organic Discussion: The use of alcohol breath analysers relies on the redox reaction between ethanol and acidified potassium dichromate ($K_2Cr_2O_7$), which changes from orange to green.