Thermochemistry, Calorimetry, and Phase Transitions

Thermochemistry and Calorimetry

Heat of Reaction and Calorimetry

Thermochemistry studies the heat changes that occur during chemical reactions. Calorimetry is the experimental technique used to measure these heat changes, often by observing temperature changes in a controlled environment.

• Calorimeter: A device used to measure the heat absorbed or released during a chemical or physical process.

• q = CΔT: The heat exchanged (q) is equal to the heat capacity (C) times the change in temperature (ΔT).

• Isobaric Calorimetry (Constant Pressure): Typically uses coffee cup calorimeters; the heat released by the system is absorbed by the surroundings (usually water).

• Bomb Calorimetry (Constant Volume): Uses a sealed container; measures heat at constant volume, often for combustion reactions.

Key Equations:

• $q = C\Delta T$

• $\Delta U = q + w$

• At constant volume: $\Delta U = q_v$

• At constant pressure: $\Delta H = \Delta U + P\Delta V$

• For reactions involving gases: $\Delta H = \Delta U + \Delta n_{gas}RT$

Example: In a coffee cup calorimeter, the heat released by a reaction is absorbed by water, and the temperature change is measured to calculate the enthalpy change.

Energy from Food: Calorimetry Applications

Calorimetry is used to determine the energy content of food by measuring the heat released during combustion.

• Food provides energy for growth, repair, regulation, and action.

• Energy is released mainly through combustion reactions.

Combustion Reactions:

• For glucose (a carbohydrate): $\mathrm{C_6H_{12}O_6(s) + 6O_2(g) \rightarrow 6CO_2(g) + 6H_2O(l)}$ $\Delta H_f^\circ = -2803\ \mathrm{kJ}$

• For glycerol trimyristate (a fat): $\mathrm{C_{45}H_{86}O_6(s) + 127/2\ O_2(g) \rightarrow 45CO_2(g) + 43H_2O(l)}$ $\Delta H_f^\circ = -27,820\ \mathrm{kJ}$

Average Energy Values:

• Carbohydrates: 4.0 kcal/g

• Fats: 9.0 kcal/g

Example: The energy content of food is determined by burning a sample in a bomb calorimeter and measuring the temperature change.

Fossil Fuels and Energy Production

Combustion of fossil fuels is a major source of heat, work, and electricity. Fossil fuels are formed from ancient organic matter subjected to pressure and bacterial decay.

• Types of Fossil Fuels: Petroleum (oil), natural gas, coal.

• Natural Gas: Primarily methane (CH4), with small amounts of ethane (C2H6), propane (C3H8), and butane (C4H10).

• Petroleum: A mixture of hydrocarbons; gasoline is a refined product.

Source	Percentage of Total Generation
Hydro	60%
Natural Gas	14%
Coal & Coke	12%
Uranium	6%
Wind	5%
Biomass	1%
Solar	1%
Petroleum	0.3%

Additional info: Canada is noted for its green electricity production, with hydroelectric power being the largest contributor.

Phase Transitions and Thermodynamics

Phase Transitions

A phase transition is the change of a substance from one physical state (solid, liquid, gas) to another. Each transition has a specific name and is associated with an enthalpy change.

• Melting (Fusion): Solid to liquid

• Freezing: Liquid to solid

• Vaporization: Liquid to gas

• Condensation: Gas to liquid

• Sublimation: Solid to gas

• Deposition: Gas to solid

Example: Water undergoes melting, vaporization, and condensation at specific temperatures and pressures.

Vapour Pressure and Boiling Point

The boiling point of a liquid depends on the external pressure. Increasing pressure raises the boiling point, while decreasing pressure lowers it. This is explained by the concept of vapour pressure.

• Vapour Pressure: The partial pressure of the vapour above the liquid at equilibrium.

• At equilibrium, the rate of evaporation equals the rate of condensation (dynamic equilibrium).

• Vapour pressure increases with temperature.

Example: At high altitudes, the lower atmospheric pressure causes water to boil at a lower temperature.

Enthalpy of Phase Changes

Energy must be added or removed for a substance to change phase. During phase changes, temperature remains constant while heat is absorbed or released.

• Heat of Fusion ($\Delta H_{fus}$): Energy required to convert a solid to a liquid at its melting point.

• Heat of Vaporization ($\Delta H_{vap}$): Energy required to convert a liquid to a gas at its boiling point.

Key Equations:

• $\mathrm{H_2O(l) \rightarrow H_2O(g)}$; $\Delta H_{vap} = 40.7\ \mathrm{kJ/mol}$

• $\mathrm{H_2O(s) \rightarrow H_2O(l)}$; $\Delta H_{fus} = 6.01\ \mathrm{kJ/mol}$

Example: When ice melts, it absorbs heat without a temperature change until all ice is converted to liquid water.

Allotropes and Reference Forms

Some elements exist in more than one form in the same physical state, called allotropes. The reference form is the most stable allotrope under standard conditions.

• Carbon: Exists as graphite (reference form) and diamond.

• Oxygen: Exists as O2 (reference form) and O3 (ozone).

Example: The enthalpy of formation for an element is defined for its reference form.

Summary of Key Concepts

• The first law of thermodynamics

• System and surroundings

• Thermochemistry and calorimetry

• Enthalpy of reactions and phase transformations

• Significant figures and sources of uncertainty

• Unit conversion

• Handling numbers representing physical quantities, including scientific notation and SI prefixes

Additional info: These concepts are foundational for understanding energy changes in chemical reactions and physical processes.