Exhaustive University Review Notes: Senior High School Chemistry Year 2 Chemistry (Ghana Ministry of Education)

Fundamental Concepts of Thermochemistry

• Systems and Surroundings: In thermochemistry, the System is the specific portion of the universe being studied, while the Surroundings comprise everything outside the system.

• Classification of Systems:

  • Open System: Capable of exchanging both matter and heat with the surroundings (e.g., water boiling in an open beaker).
  • Closed System: Capable of exchanging heat but not matter with the surroundings (e.g., water in a sealed beaker).
  • Isolated System: Incapable of exchanging either matter or heat with the surroundings. A perfectly isolated system is theoretical; real-world approximations include highly insulated calorimeters or thermos flasks.

• Enthalpy ($H$): A thermodynamic property describing the total energy of a system, defined as the sum of its internal energy ($U$) and the product of its pressure ($P$) and volume ($V$):

  • $H = U + PV$
  • Isobaric Reactions: Reactions occurring at constant pressure, where the enthalpy change equals the heat added or removed.

• Enthalpy Change ($\Delta H$): The difference between the heat content of the products and the reactants:

  • $\Delta H = H_{\text{products}} - H_{\text{reactants}}$

• Reaction Types:

  • Exothermic: Release heat to the surroundings. The products have lower energy than the reactants ($\Delta H < 0$). Example: Combustion of methane ($\text{CH}_4(g) + 2\text{O}_2(g) \rightarrow \text{CO}_2(g) + 2\text{H}_2\text{O}(l) \quad \Delta H = -890.4\,kJ\,mol^{-1}$).
  • Endothermic: Absorb heat from the surroundings. The products have higher energy than the reactants ($\Delta H > 0$). Example: Decomposition of calcium carbonate ($\text{CaCO}_3(s) \rightarrow \text{CaO}(s) + \text{CO}_2(g) \quad \Delta H = +177.8\,kJ\,mol^{-1}$).

Standard Enthalpy Definitions and Varieties

• Standard Conditions:

  • Pressure: $1\,atm$ ($101.3\,kPa$).
  • Temperature: $298\,K$ ($25^\text{o}\text{C}$).
  • Concentration: $1.0\,mol\,dm^{-3}$.
  • Symbol: $\Delta H^\circ$

• Specific Standard Enthalpies:

  • Standard Enthalpy of Formation ($\Delta H_f^\circ$): Heat change when 1 mole of a compound is formed from its constituent elements in their standard states under standard conditions. The $\Delta H_f^\circ$ for all elements in their standard states is zero ($H_2(g)$, $O_2(g)$, $C(\text{graphite})$).
  • Standard Enthalpy of Combustion ($\Delta H_c^\circ$): Heat change when 1 mole of a substance is completely burned in excess oxygen. Always negative (exothermic).
  • Standard Enthalpy of Neutralisation ($\Delta H_n^\circ$): Heat change when an acid and alkali react to form 1 mole of water. For strong acids/bases, this is approximately $-57.3\,kJ\,mol^{-1}$ because both react via the same net ionic equation: $\text{H}^+(aq) + \text{OH}^-(aq) \rightarrow \text{H}_2\text{O}(l)$.
  • Standard Enthalpy of Solution ($\Delta H_{\text{soln}}^\circ$): Heat change when 1 mole of an ionic substance dissolves in excess water to infinite dilution.
  • Standard Enthalpy of Hydration ($\Delta H_{\text{hyd}}^\circ$): Heat change when 1 mole of gaseous ions dissolves in water to form aqueous ions.
  • Atomisation Energy: Energy required to break all bonds in 1 mole of a substance to form individual gaseous atoms.
  • Lattice Energy: Energy required to separate 1 mole of an ionic solid into its gaseous constituent ions.

Quantitative Thermochemical Laws and Cycles

• Hess's Law of Constant Heat Summation: The total enthalpy change of a chemical reaction is independent of the pathway taken, provided the initial and final states are the same.

  • Mathematically: $\Delta H_{\text{rxn}} = \sum n\Delta H_f^\circ(\text{products}) - \sum m\Delta H_f^\circ(\text{reactants})$

• Born-Haber Cycles: Energy cycles used to calculate lattice energies of ionic compounds by relating formation, sublimation, ionization, dissociation, and electron affinity. For Lithium Fluoride ($\text{LiF}$):

  1. Sublimation of Li: $\text{Li}(s) \rightarrow \text{Li}(g) \quad \Delta H = +155.2\,kJ\,mol^{-1}$
  2. Dissociation of F: $\frac{1}{2}\text{F}_2(g) \rightarrow \text{F}(g) \quad \Delta H = +75.3\,kJ\,mol^{-1}$
  3. Ionization of Li: $\text{Li}(g) \rightarrow \text{Li}^+(g) + e^- \quad \Delta H = +520\,kJ\,mol^{-1}$
  4. Electron affinity of F: $\text{F}(g) + e^- \rightarrow \text{F}^-(g) \quad \Delta H = -328\,kJ\,mol^{-1}$
  5. Formation of LiF lattice: $\Delta H_{\text{formation}} = \Delta H_{\text{sub}} + \Delta H_{\text{diss}} + \Delta H_{\text{IE}} + \Delta H_{\text{EA}} + \Delta H_{\text{lattice}}$

• Bond Enthalpy: The average energy required to break 1 mole of a specific covalent bond in a gaseous molecule.

  • Breaking bonds: Endothermic (energy absorbed).
  • Forming bonds: Exothermic (energy released).
  • $\Delta H_{\text{rxn}} = \sum \Delta H_{\text{bonds broken}} - \sum \Delta H_{\text{bonds formed}}$

Chemical Kinetics: Rates and Mechanisms

• Reaction Rate: The change in concentration of a reactant or product per unit time ($mol\,dm^{-3}\,s^{-1}$).

  • Formula for $a\text{A} + b\text{B} \rightarrow c\text{C} + d\text{D}$:
  • $\text{Rate} = -\frac{1}{a}\frac{\Delta[A]}{\Delta t} = -\frac{1}{b}\frac{\Delta[B]}{\Delta t} = \frac{1}{c}\frac{\Delta[C]}{\Delta t} = \frac{1}{d}\frac{\Delta[D]}{\Delta t}$

• Rate Classifications:

  • Average Rate: Change over a specific time interval ($\Delta t$).
  • Instantaneous Rate: The rate at a specific moment, calculated by the tangent slope of a concentration-time graph at that point.
  • Initial Rate: The rate at $t = 0$.

• Collision Theory: For a reaction to occur, particles must collide with:

  1. Energy greater than or equal to the Activation Energy ($E_a$).
  2. Correct spatial orientation.

• Factors Affecting Rate:

  • Temperature: Increases kinetic energy; a greater fraction of molecules exceed $E_a$. Doubling temperature usually significantly accelerates reaction frequency.
  • Concentration/Pressure: Increases the frequency of collisions per unit volume.
  • Surface Area: Finer particles increase the collision sites available in solid-phase reactions.
  • Catalyst: Provides an alternative reaction pathway with a lower $E_a$. It is not consumed in the reaction.

• Rate Law and Expression:

  • $\text{Rate} = k[A]^x[B]^y$
  • $k$: Specific rate constant. For 0-order: $M\,s^{-1}$; 1st-order: $s^{-1}$; 2nd-order: $M^{-1}\,s^{-1}$.
  • Order of Reaction: The exponent to which a reactant concentration is raised. Summed exponents give the Overall Order.
  • Half-life ($t_{1/2}$): Time for concentration to reduce by 50%.
    • 1st Order: $t_{1/2} = \frac{0.693}{k}$

Dynamic Equilibrium

• Characteristics: Occurs in a closed system; forward and reverse reactions proceed at the same rate; concentrations of reactants and products remain constant over time.

• Equilibrium Constants:

  • For $a\text{A} + b\text{B} \rightleftharpoons c\text{C} + d\text{D}$
  • $K_c = \frac{[C]^c[D]^d}{[A]^a[B]^b}$
  • $K_p = \frac{P_C^c \times P_D^d}{P_A^a \times P_B^b}$
  • Relationship: $K_p = K_c(RT)^{\Delta n}$ (where $\Delta n = (c+d) - (a+b)$).

• Heterogeneous Equilibria: Concentrations of pure solids and pure liquids are omitted from the expression as they remain constant.

• Le Chatelier’s Principle: If a system at equilibrium is disturbed, the system shifts to counteract the disturbance.

  • Concentration: Adding a reactant shifts equilibrium to the product side.
  • Pressure: Increasing pressure shifts the system toward the side with fewer gas molecules.
  • Temperature: Increasing temperature shifts equilibrium toward the endothermic direction. Only temperature changes the actual value of $K$.
  • Catalysts: Increase rate to reach equilibrium faster but do not change the equilibrium constant or position.

Systematic Chemistry: Period 3 and Halogens

• Period 3 Trends (Na to Ar):

  • Atomic Radius: Decreases due to increasing nuclear charge pulling electrons closer.
  • Electronegativity: Increases.
  • Melting Points: High for metals ($Na, Mg, Al$), highest for Si (giant covalent), low for non-metals ($P_4, S_8, Cl_2$ based on Van der Waals forces ($S_8 > P_4 > Cl_2 > Ar$)).
  • Oxides: Basic ($Na_2O, MgO$) $\rightarrow$ Amphoteric ($Al_2O_3$) $\rightarrow$ Acidic ($SiO_2, P_4O_{10}, SO_2, Cl_2O_7$).

• Group 17 (Halogens):

  • Physical States: Fluorine/Chlorine (gas), Bromine (liquid), Iodine (solid).
  • Oxidising Power: Decreases down the group ($F_2 > Cl_2 > Br_2 > I_2$).
  • Displacement: A more reactive halogen displaces a less reactive halide from its salt ($\text{Cl}_2 + 2\text{KBr} \rightarrow 2\text{KCl} + \text{Br}_2$).
  • Halides + Sulfuric Acid:
    • $Cl^-$ + $H_2SO_4$: Produces $HCl$ (Steamy fumes).
    • $Br^-$ + $H_2SO_4$: Produces $HBr, Br_2$ (brown fumes), and $SO_2$.
    • $I^-$ + $H_2SO_4$: Produces $HI, I_2$ (purple fumes), $SO_2, S$ (yellow solid), and $H_2S$ (rotten egg smell).

Acids, Bases, and Salts

• Theories:

  • Arrhenius: Acids produce $H^+$/$H_3O^+$ in water; Bases produce $OH^-$ in water.
  • Brønsted-Lowry: Acids are proton ($H^+$) donors; Bases are proton acceptors.
  • Lewis: Acids are electron pair acceptors; Bases are electron pair donors.

• Terminology:

  • Conjugate Acid-Base Pair: Two species differing by a single $H^+$.
  • Amphiprotic: Species acting as both acid and base ($H_2O, HCO_3^-$).

• Titrations:

  • Standard Solution: Known concentration ($mol\,dm^{-3}$).
  • Endpoint: Visual signal (color change) that the reaction is complete.
  • Back Titration: Reacting an analyte with an excess of a standard reagent, then titrating the remaining excess.

Bonding and Molecular Geometry

• VSEPR Theory: Molecular shape is determined by the number of bonding and lone pairs of electrons around the central atom, minimizing repulsion.

  • Repulsion Order: LP-LP > LP-BP > BP-BP.
  • Shapes:
    • 2 charge centers: Linear ($180^\text{o}$, e.g., $CO_2$).
    • 3 charge centers: Trigonal Planar ($120^\text{o}$, e.g., $BF_3$).
    • 4 charge centers: Tetrahedral ($109.5^\text{o}$, e.g., $CH_4$); Pyramidal ($107^\text{o}$, e.g., $NH_3$); Bent ($104.5^\text{o}$, e.g., $H_2O$).

• Hybridisation:

  • $sp^3$: 4 sigma bonds, tetrahedral (Alkanes).
  • $sp^2$: 3 sigma bonds, 1 pi bond, trigonal planar (Alkenes/Benzene).
  • $sp$: 2 sigma bonds, 2 pi bonds, linear (Alkynes).

Organic Chemistry: Carbon Compounds

• Alkanes ($C_n H_{2n+2}$): Saturated hydrocarbons. Primary reaction is free-radical substitution (chlorination).

• Alkenes ($C_n H_{2n}$): Unsaturated (double bond). Reactions: Electrophilic addition (Hydrogenation with Ni catalyst, hydration to form alcohols).

  • Markovnikov’s Rule: In the addition of $HX$ to an unsymmetrical alkene, the hydrogen attaches to the carbon with more hydrogens.

• Alkynes ($C_n H_{2n-2}$): Triple bond. ethyne produced from calcium carbide: $\text{CaC}_2 + 2\text{H}_2\text{O} \rightarrow \text{C}_2\text{H}_2 + \text{Ca}(\text{OH})_2$.

• Benzene ($C_6H_6$): Aromatic ring with delocalized pi electrons. Undergoes Electrophilic Substitution (Nitration with $HNO_3/H_2SO_4$; Friedel-Crafts alkylation with $AlCl_3$ catalyst).

• Alkanols ($R-OH$): Primary, Secondary, Tertiary.

  • Lucas Test: Tertiary alcohols turn cloudy immediately; Secondary take 5-10 mins; Primary no reaction.
  • Iodoform Test: Positive for Ethanol and secondary alcohols with a methyl group next to the $-OH$.

• Alkanoic Acids ($R-COOH$): Weak acids. Form esters with alcohols via concentrated $H_2SO_4$ catalyst: $\text{Acid} + \text{Alcohol} \rightleftharpoons \text{Ester} + \text{Water}$.