Thermodynamics

1. Thermodynamics

1.1 First Law of Thermodynamics

  • The Law of Conservation of Energy states that energy can be exchanged between a system and its surroundings, but the total energy remains constant.

    • System: The part of the environment concerned (e.g., a chemical reaction).

    • Surroundings: Everything outside the system.

  • Energy exchanges:

    • Endothermic: The system absorbs energy/heat from surroundings.

    • Exothermic: The system loses energy/heat to surroundings.

1.2 Key Terms

  • Enthalpy (H): A thermodynamic function to describe the energy/heat content of a system.

    • ΔH = Hfinal – Hinitial.

  • For chemical reactions: ΔH = Hproducts – Hreactants.

  • Heat of Reaction (ΔHrxn): The enthalpy change during a reaction.

  • If ΔH > 0: Endothermic reaction (energy supplied).

  • If ΔH < 0: Exothermic reaction (energy released).

1.3 Spontaneity of Reactions

  • Exothermic reactions are generally more favorable as they do not require additional energy, rendering them thermodynamically spontaneous.

2. Calorimetry

2.1 Heat Flow and Measurement

  • Calorimetry: A method to measure heat flow during a reaction.

  • Calorimeter: Device for measuring heat flow (e.g., Styrofoam cup).

2.2 Example Calculation

  • Heating Water: Determine heat for warming 250g of water from 22°C to 98°C.

    • Specific heat capacity of water: 4.18 J g-1 K-1.

    • Formula: q = mcΔT where

      • q = heat (J),

      • m = mass (g),

      • c = specific heat capacity,

      • ΔT = change in temperature.

  • Calculation:

    • ΔT = 98 – 22 = 76.

    • q = (250 g)(4.18 J g-1 K-1)(76 K) = 7.9 x 10^4 J.

3. Bond Enthalpy and Energetics

3.1 Bond Properties

  • Bond Length: Single bonds > Double bonds > Triple bonds in length.

  • Bond Strength: C–C < C=C < C≡C (as bond length increases, bond strength decreases).

  • Bond Enthalpy: Energy required to break a bond.

    • Example: ΔH of Cl2 = 242 kJ when 1 mole dissociates into Cl atoms.

3.2 Reaction Example

  • For the reaction: CH4(g) + Cl2(g) → CH3Cl(g) + HCl(g).

    • Calculate ΔHrxn using bond enthalpy values.

    • Total energy needed = sum of bonds broken - sum of energy released by bonds formed.

3.3 Hess’ Law

  • States ΔH for an overall reaction equals the sum of ΔH for individual steps.

  • Useful for calculating energy changes not directly measurable.

  • Example: Use thermochemical equations to find ΔH for the formation of ethene from ethane.

4. Born-Haber Cycle

4.1 Ionic Compounds Formation

  • Involves energy changes when forming ionic compounds.

  • Born-Haber Cycle: Graphical representation of energy changes in ionic compound formation.

4.2 Energy Changes Defined

  • Enthalpy of Formation (ΔHf): Energy change when forming one mole of a compound from its elements.

  • Enthalpy of Vaporization (ΔHvap): Energy change when converting a solid/liquid to gas.

  • Enthalpy of Atomization (ΔHatom): Change when a gaseous molecule is converted into its atoms.

  • Ionization Energy (ΔHi.e.): Energy to remove electrons from gaseous atoms to form cations.

  • Electron Affinity (ΔHe.a.): Energy change when atoms gain electrons to form anions.

  • Lattice Energy (ΔHlatt): Energy released when an ionic compound forms from gaseous ions.

5. Coulomb’s Law and Lattice Energy

5.1 Electrostatic Forces

  • Coulomb’s Law relates force of attraction or repulsion between charged particles to their charge and distance.

  • Formula: F = k (Q1Q2 / d^2)

  • Energy = Force x Distance, where E represents lattice energy.

5.2 Example Calculations

  • Use Coulomb’s Law to determine energy for specific ionic interactions.

  • Calculate lattice energy per mole of ions and total reaction energy using Hess’ Law.

5.3 Summary of Steps

  • Calculate lattice energy, write all transformations, sum ΔH values to find overall energy change.

  • Consider the proportions and relationships through each aspect of the Born-Haber cycle.