Senior High School Chemistry: Comprehensive Year 2 Study Notes

THERMODYNAMICS AND ENERGY CHANGES

• Fundamental Definitions     - 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 energy of a system, calculated as: $H = U + PV$, where $U$ is internal energy, $P$ is pressure, and $V$ is volume.     - Enthalpy Change (\Delta H): Defined as: $\Delta H = H_{products} - H_{reactants}$.     - Hess’s Law of Constant Heat Summation: States that the total enthalpy change of a chemical reaction is equal to the sum of all individual enthalpy changes, regardless of the pathway taken.     - Lattice Energy: The energy required to separate 1 mole of an ionic solid into gaseous ions.

• Types of Chemical Systems     - Open System: Can exchange both heat and matter with the surroundings (e.g., heating water in an open beaker).     - Closed System: Can exchange heat but not matter with the surroundings (e.g., heating water in a sealed beaker).     - Isolated System: Cannot exchange heat or matter with surroundings (e.g., well-insulated thermos or calorimeter). Note: No system is truly ideal; some heat/matter leakage always occurs.

• Exothermic and Endothermic Reactions     - Exothermic Reactions: Release heat to the surroundings. The enthalpy of the system decreases ($\Delta H$ is negative). Example: Methane combustion: $CH_4(g) + 2O_2(g) \rightarrow CO_2(g) + 2H_2O(l)$, $\Delta H = -890.4\,kJ/mol$.     - Endothermic Reactions: Absorb heat from surroundings. The enthalpy of the system increases ($\Delta H$ is positive). Example: Decomposition of calcium carbonate: $CaCO_3(s) \rightarrow CaO(s) + CO_2(g)$, $\Delta H = +177.8\,kJ/mol$.

• Standard Enthalpy Changes (\Delta H^\circ)     - Standard Conditions: Pressure of $1\,atm$ ($101.3\,kPa$) and Temperature of $298\,K$ ($25\,^\circ C$).     - Standard State: The most stable form of an element/compound at standard conditions (e.g., $O_2(g)$, $H_2O(l)$, $C(graphite)$).     - Standard Enthalpy of Formation (\Delta H_f^\circ): Enthalpy change when one mole of a compound is formed from its elements in their standard states. Elements in their standard states have $\Delta H_f^\circ = 0$.     - Standard Enthalpy of Reaction (\Delta H_{rxn}^\circ): Calculated as: $\Delta H_{rxn}^\circ = \sum n\Delta H_f^\circ(products) - \sum m\Delta H_f^\circ(reactants)$.     - Standard Enthalpy of Combustion (\Delta H_c^\circ): Enthalpy change when one mole of a substance is burned completely in oxygen. Example: $H_2(g) + \frac{1}{2}O_2(g) \rightarrow H_2O(l)$, $\Delta H_c^\circ = -285.8\,kJ/mol$.     - Standard Enthalpy of Neutralisation (\Delta H_n^\circ): Enthalpy change when 1 mole of water is produced from acid-alkali reaction. For strong acids/bases, this value is approximately $\Delta H = -57.4\,kJ/mol$.     - Standard Enthalpy of Solution (\Delta H_{soln}^\circ): Heat change when one mole of ionic substance dissolves in excess water.     - Standard Enthalpy of Hydration (\Delta H_{hyd}^\circ): Energy evolved when one mole of gaseous ions is surrounded and stabilized by water molecules (infinite dilution).

• Experimental Determination and Calorimetry     - Energy Transfer Formula: $Q = m_{water} \times c \times \Delta T$, where $c = 4.18\,J/g^\circ C$.     - Heat of Combustion Calculation: Per mole = $\frac{Q}{mass\,burned} \times Molar\,Mass$.

• Hess's Law Manipulations and Born-Haber Cycles     - Rules for Equations: Reversing an equation negates the sign of $\Delta H$. Multiplying coefficients by a factor requires multiplying $\Delta H$ by the same factor.     - Born-Haber Cycle for LiF: Shows steps: Sublimation of $Li(s)$ ($+155.2\,kJ/mol$), Dissociation of $\frac{1}{2}F_2(g)$ ($+75.3\,kJ/mol$), Ionisation of $Li(g)$ ($+520\,kJ/mol$), Electron affinity of $F(g)$ ($-328\,kJ/mol$), and Formation of $LiF(s)$ ($-594.1\,kJ/mol$). Calculated Lattice Energy $x = -1016.6\,kJ/mol$.

CHEMICAL KINETICS AND REACTION RATES

• Rate Principles     - Rate of Reaction: Change in concentration of reactant or product per unit time. Unit: $mol\,dm^{-3}\,s^{-1}$.     - Mathematical Rate for $aA + bB \rightarrow cC + dD$: $Rate = -\frac{1}{a} \frac{\Delta[A]}{\Delta t} = -\frac{1}{b} \frac{\Delta[B]}{\Delta t} = \frac{1}{c} \frac{\Delta[C]}{\Delta t} = \frac{1}{d} \frac{\Delta[D]}{\Delta t}$.     - Initial Rate: Rate at $t = 0$.     - Instantaneous Rate: Rate at a specific instant, determined by the tangent slope of the concentration-time curve.     - Average Rate: $\frac{\Delta[Concentration]}{\Delta t}$.

• Factors Affecting Reaction Rate     - Temperature: Higher temperature increases average kinetic energy; a greater proportion of molecules exceed Activation Energy ($E_a$).     - Concentration/Pressure: More particles in a given volume increases collision frequency.     - Surface Area: More contact points in solids lead to more frequent effective collisions.     - Catalyst: Lowers $E_a$ by providing an alternative reaction pathway; it is not consumed.

• Collision Theory and Maxwell-Boltzmann     - Collision Theory Requirements: Particles must collide with (1) energy $\ge E_a$ and (2) correct orientation.     - Maxwell-Boltzmann Distribution: At higher temperatures (T_2 > T_1), the peak shift right and lowers; the area under the curve beyond $E_a$ increases significantly.

• Rate Equations and Reaction Order     - Rate Law: $Rate = k[A]^x[B]^y$, where $k$ is the rate constant, and $x, y$ are orders determined experimentally.     - Zero Order: Rate is independent of concentration ($k$ units: $mol\,dm^{-3}\,s^{-1}$).     - First Order: Rate doubles when concentration doubles ($k$ units: $s^{-1}$). Half-life $t_{1/2} = \frac{0.693}{k}$.     - Second Order: Rate quadruples when concentration doubles ($k$ units: $mol^{-1}\,dm^3\,s^{-1}$). Half-life $t_{1/2} = \frac{1}{k[A]_0}$.     - Rate-Determining Step (RDS): The slowest elementary step in a reaction mechanism.

DYNAMIC EQUILIBRIUM IN CHEMICAL SYSTEMS

• The Nature of Equilibrium     - Reversible Reactions: Products can convert back to reactants (e.g., $N_2 + 3H_2 \rightleftharpoons 2NH_3$).     - Dynamic Equilibrium: A state in a closed system where forward and reverse rates are equal, and concentrations remain constant.

• Equilibrium Constants     - Law of Mass Action: For $aA + bB \rightleftharpoons cC + dD$, $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$.     - Partial Pressure Constant ($K_p$): For gases, $K_p = \frac{P_C^c \times P_D^d}{P_A^a \times P_B^b}$.     - Relation Formula: $K_p = K_c(RT)^{\Delta n}$, where $\Delta n = (c+d) - (a+b)$.     - Solubility Product ($K_{sp}$): For sparingly soluble salt $A_xB_y(s) \rightleftharpoons xA^{y+}(aq) + yB^{x-}(aq)$, $K_{sp} = [A^{y+}]^x[B^{x-}]^y$. Pure solids/liquids are excluded from expressions.

• Le Chatelier’s Principle     - Concentration: Adding reactants shifts equilibrium to use them up (forward).     - Pressure/Volume: Increasing pressure (decreasing volume) shifts equilibrium to the side with fewer gas moles.     - Temperature: Increasing temperature shifts in the endothermic direction (\Delta H > 0) to absorb heat.     - Catalyst: Reaches equilibrium faster but does not shift the position or change $K$.

• Industrial Applications     - Haber Process: $N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$, $\Delta H = -92\,kJ/mol$. Optimized by high pressure, iron catalyst, and removing ammonia. Lower temperature favors yield but higher (compromised) temperature favors rate.     - Contact Process: $2SO_2(g) + O_2(g) \rightleftharpoons 2SO_3(g)$, $\Delta H = -196\,kJ/mol$. Uses Vanadium(V) oxide ($V_2O_5$) catalyst.     - Catalytic Converters: Reduces $NO$ to $N_2$ and oxidizes $CO$ to $CO_2$ using platinum catalysts.

THE CHEMISTRY OF ACIDS, BASES, AND SALTS

• Acid-Base Theories     - Arrhenius: Acids produce $H^+/H_3O^+$ in water; Bases produce $OH^-$ in water. Limit: Only applies to aqueous solutions.     - Brønsted-Lowry: Acid is a proton ($H^+$) donor; Base is a proton acceptor. Conjugate pairs differ by a single $H^+$. Water is amphiprotic.     - Lewis: Acid is an electron-pair acceptor (e.g., $BF_3$, $Al^{3+}$); Base is an electron-pair donor (e.g., $NH_3$, $H_2O$). Forms coordinate covalent bonds.

• Properties and Applications     - Physical: Acids (sour, pH < 7, turn blue litmus red); Bases (bitter, slippery, pH > 7, turn red litmus blue).     - Chemical: Acids + Carbonates $\rightarrow$ Salt + $H_2O$ + $CO_2$; Bases + Ammonium salts + heat $\rightarrow$ Salt + $H_2O$ + $NH_3(g)$.     - Neutralisation In Life: Antacids ($Mg(OH)_2$, $Al(OH)_3$) for reflux; Agriculture (lime/$CaCO_3$) for soil acidity.

• Salts and Solubility     - Classification:         - Acidic Salts: Partial replacement of acid hydrogens (e.g., $NaHSO_4$).         - Double Salts: Two different cations (e.g., Potash alum).         - Complex Salts: Central metal with ligands (e.g., $[Cr(NH_3)_6]Cl_3$).         - Deliquescent: Absorb enough air moisture to dissolve (e.g., $NaOH$, $CaCl_2$).         - Efflorescent: Lose water of crystallization to air (e.g., $Na_2CO_3 \cdot 10H_2O$).

• Acid-Base Titration     - Key Terms: Titrant (burette, known conc); Analyte (flask, unknown conc); Equivalence Point (molar equivalence); Endpoint (indicator colour change).     - Direct Titration Calculation: Mole ratio $\frac{n_{base}}{n_{acid}} = \frac{C_bV_b}{C_aV_a}$.     - Back Titration: Used when direct measurement is difficult (e.g., Ecotrin aspirin analysis). Add excess known reagent, then titrate the excess.     - Double-Indicator Titration: Uses phenolphthalein and methyl orange to analyze mixtures of $Na_2CO_3$ and $NaHCO_3$.

PERIODICITY AND PROPERTIES OF ELEMENTS

• Period 3 Trends     - Metallic Character: Decreases from $Na$ to $Ar$. $Na, Mg, Al$ are metals; $Si$ is a metalloid; $P, S, Cl, Ar$ are non-metals.     - Melting/Boiling Points: Increase from $Na$ to $Si$ ($Si$ has a giant covalent network), then decrease. Trend for non-metals: S_8 > P_4 > Cl_2 > Ar (based on Van der Waals forces/molecular size).     - Electrical Conductivity: Decreases across. $Al$ is the best conductor (3 delocalized electrons). $Si$ is a semiconductor.

• Period 3 Compounds     - Oxides: $Na_2O, MgO$ (Basic, ionic); $Al_2O_3$ (Amphoteric); $SiO_2, P_4O^{10}, SO_3$ (Acidic, covalent).     - Chlorides: $NaCl, MgCl_2$ (Ionic, neutral in water); $AlCl_3, SiCl_4, PCl_5$ (Covalent, hydrolysis in water to produce acidic $HCl$ fumes).     - Thermal Stability: Carbonates and nitrates become more stable down Group 1/2 as cation size increases and polarizing power decreases.

THE HALOGENS (GROUP 17)

• Physical and Chemical Trends     - Physical State: $F_2, Cl_2$ (Gas); $Br_2$ (Liquid); $I_2, At$ (Solid). Colour darkens down the group.     - Oxidising Strength: Decreases down (F_2 > Cl_2 > Br_2 > I_2). Standard reduction potential for $F_2/2F^- = +2.87\,V$.     - Reducing Power of Halides: Increases down (I^- > Br^- > Cl^- > F^-).

• Reactions with $H_2SO_4(conc)$     - Chlorides: Produce $HCl(g)$ only.     - Bromides: Produce $HBr$, which further reduces $H_2SO_4$ to $SO_2$ and $Br_2(g)$.     - Iodides: Strongest reducer; reduces $H_2SO_4$ to $SO_2, S, H_2S$, and $I_2(g)$.

• Hydrogen Halides (HX)     - Acid Strength: HF \ll HCl < HBr < HI. $HF$ is weak due to high bond strength ($565\,kJ/mol$). $HI$ is the strongest acid ($K_a \approx 5.0 \times 10^{10}$).     - Thermal Stability: Decreases down as bond length increases and bond energy decreases.

MOLECULAR STRUCTURE AND BONDING

• Electronegativity and Polarity     - Pauling Scale: $F = 4.0$, $Cs = 0.7$. Difference \Delta EN > 1.7 is ionic; $0.5 - 1.7$ is polar covalent.     - Dipole Moment ($D$): Symmetrical molecules (e.g., $CCl_4, CO_2$) have no net dipole even if bonds are polar. Asymmetrical molecules (e.g., $H_2O, NH_3$) are polar.

• VSEPR Theory     - Linear: $CO_2$, $180^\circ$.     - Trigonal Planar: $BF_3$, $120^\circ$.     - Tetrahedral: $CH_4$, $109.5^\circ$.     - Trigonal Pyramidal: $NH_3$, $107^\circ$.     - Bent: $H_2O$, $104.5^\circ$.

• $Sigma (\sigma)$ and $Pi (\pi)$ Bonds     - Sigma (\sigma): Head-on orbital overlap ($s-s, s-p, p-p$). Allows free rotation. Found in single bonds.     - Pi (\pi): Lateral/sideways p-orbital overlap. Restricted rotation. Found in double (1 \sigma, 1 \pi) and triple (1 \sigma, 2 \pi) bonds.

• Hybridization in Carbon     - $sp^3$: 4 equivalent orbitals, tetrahedral ($CH_4, C_2H_6$).     - $sp^2$: 3 equivalent orbitals + 1 unhybridized p-orbital, trigonal planar ($C_2H_4$, benzene).     - $sp$: 2 equivalent orbitals + 2 unhybridized p-orbitals, linear ($C_2H_2$).

ORGANIC FUNCTIONAL GROUPS

• Alkanes ($C_n H_{2n+2}$)     - Reactivity: Low due to strong, non-polar \sigma bonds. Undergo free-radical halogenation via initiation, propagation, and termination steps.     - Cracking/Reforming: Thermal or catalytic processes to produce shorter fuels or branched isomers (higher octane).

• Alkenes ($C_n H_{2n}$)     - Reactions: Electrophilic addition.         - Hydrogenation: Uses $Ni$ catalyst, $130\,^\circ C$.         - Markovnikov’s Rule: In unsymmetrical additions (e.g., $HX$ to propene), $H$ adds to the carbon with more hydrogens.     - Bromine Water Test: Alkenes decolorize reddish-brown bromine water; alkanes do not.

• Benzene ($C_6H_6$)     - Bonding: $sp^2$ hybridized with delocalized \pi-electrons above and below the ring. All $C-C$ bond lengths are equal.     - Reactions: Prefers Electrophilic Substitution (Nitration, Halogenation, Friedel-Crafts) to preserve aromatic stability.

• Alkanols ($C_n H_{2n+1}OH$)     - Classification: Primary (1\circ), Secondary (2\circ), Tertiary (3\circ).     - Oxidation: 1\circ \rightarrow Aldehyde \rightarrow Carboxylic Acid; 2\circ \rightarrow Ketone; 3\circ \rightarrow No reaction.     - Lucas Test: Tertiary alcohols turn cloudy instantly; secondary in 5-10 mins; primary only on heating.

• Alkanoic Acids ($R-COOH$)     - Acidity: Weak acids that form dimers via hydrogen bonding (high boiling points).     - Esterification: Carboxylic Acid + Alkanol $\rightleftharpoons$ Ester + $H_2O$ ($H_2SO_4$ catalyst).     - Sodium Bicarbonate Test: Carboxylic acids release $CO_2(g)$ (effervescence); alkanols do not.

1. Question 1: Consider the following reversible reaction: $N<em>2(g) + 3H</em>2(g) \rightleftharpoons 2NH_3(g)$.

   • Describe what happens to the equilibrium position if the concentration of $H_2(g)$ is increased.

2. Question 2: For the exothermic reaction: $A(g) + B(g) \rightleftharpoons C(g) + D(g)$, explain how increasing the temperature will affect the yield of products at equilibrium.

3. Question 3: Given the equilibrium constant expression $K_c = \frac{[C][D]}{[A][B]}$, what does it signify about the relationship between the concentrations of products and reactants at equilibrium?

4. Question 4: In the Contact Process: $2SO<em>2(g) + O</em>2(g) \rightleftharpoons 2SO<em>3(g)$, it is known that an increase in pressure shifts the equilibrium towards the production of $SO</em>3$. Describe the reasoning behind this shift in equilibrium position.

5. Question 5: How does the addition of a catalyst affect the position of equilibrium in a chemical reaction?