Thermochemistry: Energy and Chemical Reactions

Theme 8: Energy and Chemical Reactions

Introduction to Thermochemistry

Thermochemistry is the study of energy changes, particularly heat, that accompany chemical reactions and physical changes. Understanding how energy is transferred and transformed is essential for predicting reaction behavior and for applications in biology, chemistry, and engineering.

Basic Principles of Energy

Definitions and Concepts

• Energy: The capacity to do work or transfer heat.

• Temperature: A measure of the average kinetic energy of particles in a substance, reported in °C, K, or °F.

• Heat (q): Energy transferred between systems due to a temperature difference, measured in Joules (J).

• Heat of Reaction: The energy flow to or from the surroundings as a result of a chemical reaction.

Note: Temperature (in K) is not the same as heat (in J).

First Law of Thermodynamics

The first law states that energy cannot be created or destroyed, only transferred or transformed. The total energy of the universe remains constant.

• System: The part of the universe under study (e.g., the contents of a reaction flask).

• Surroundings: Everything outside the system (e.g., air, container, lab bench).

Types of Systems

• Open System: Exchanges both energy and matter with surroundings (e.g., hot soup in an open cup).

• Closed System: Exchanges energy but not matter (e.g., soup in a cup with a lid).

• Isolated System: Exchanges neither energy nor matter (e.g., soup in a perfectly insulated thermos).

Heat Transfer and Thermal Equilibrium

Heat flows spontaneously from a hotter object to a cooler one until thermal equilibrium is reached (no net heat flow).

Exothermic and Endothermic Processes

Definitions

• Exothermic: Energy is released from the system to the surroundings (q < 0).

• Endothermic: Energy is absorbed from the surroundings into the system (q > 0).

Units of Energy

• 1 calorie (cal) = heat required to raise the temperature of 1.00 g of H2O by 1.0°C.

• 1 cal = 4.184 J (Joules)

• Food energy: 1 Calorie (Cal) = 1 kcal = 1000 cal

Heat Capacity and Specific Heat

Definitions

• Heat Capacity (C): The amount of heat required to raise the temperature of an object by 1°C (J/°C).

• Specific Heat Capacity (c): The amount of heat required to raise the temperature of 1 g of a substance by 1°C (J/g·K).

• Molar Heat Capacity: The amount of heat required to raise the temperature of 1 mol of a substance by 1°C (J/mol·K).

Factors Affecting Heat Absorption

• Mass of the substance (m)

• Magnitude of temperature change (ΔT)

• Specific heat capacity of the material (c)

The heat absorbed or released is calculated as:

$q = m \times c \times \Delta T$

Energy and Changes of State (Phase Changes)

Phase Changes

Energy is required for phase changes (e.g., melting, boiling) without a change in temperature. The amount of energy depends on the mass and the enthalpy of the phase change.

Example Calculation

To convert 500.0 g of water from liquid to gas at 100°C:

$q = 2256 \ \text{J/g} \times 500.0 \ \text{g} = 1.13 \times 10^6 \ \text{J} = 1130 \ \text{kJ}$

First Law of Thermodynamics (Internal Energy)

Internal Energy (U)

• The sum of all kinetic and potential energies of the particles in a system.

• Change in internal energy: $\Delta U = q + w$

• q = heat exchanged, w = work done

Sign Conventions

Energy Transferred As...	Sign Convention	Effect on Usystem
Heat to the system (endothermic)	q > 0 (+)	U increases
Heat from the system (exothermic)	q < 0 (−)	U decreases
Work done on system	w > 0 (+)	U increases
Work done by system	w < 0 (−)	U decreases

Enthalpy Changes for Chemical Reactions

Enthalpy (H)

• A state function representing heat content at constant pressure.

• Change in enthalpy: $\Delta H = H_{products} - H_{reactants}$

• Depends on the amount, direction, and phase of substances.

Thermochemical Equations

• Show the enthalpy change associated with a reaction.

• Example: $\text{CaO(s) + H}_2\text{O}(\ell) \rightarrow \text{Ca(OH)}_2(\text{s}) \quad \Delta_rH = -65 \ \text{kJ}$

State Functions

• Properties that depend only on the current state, not the path taken (e.g., T, P, V, H).

Calorimetry

Constant Pressure Calorimetry (Coffee-Cup Calorimeter)

• Used to measure enthalpy changes at constant pressure.

• Commonly used for reactions in solution.

Constant Volume Calorimetry (Bomb Calorimeter)

• Used to measure changes in internal energy at constant volume.

• Suitable for combustion reactions.

Hess's Law

Principle

Hess's Law states that the total enthalpy change for a reaction is the same, regardless of the number of steps or the pathway taken, because enthalpy is a state function.

• Allows calculation of enthalpy changes for reactions that cannot be measured directly.

• Sum the enthalpy changes for individual steps to find the overall change.

Standard Enthalpy of Formation

• The enthalpy change for the formation of 1 mole of a compound from its elements in their standard states.

Substance	Formula	ΔHf° (kJ/mol)
Water (liquid)	H2O(l)	-285.8
Carbon dioxide (gas)	CO2(g)	-393.5
Ammonia (gas)	NH3(g)	-45.9
Methane (gas)	CH4(g)	-74.8
Glucose (solid)	C6H12O6(s)	-1273

Summary Table: Key Equations

Equation	Description
$q = m \times c \times \Delta T$	Heat transfer for temperature change
$\Delta U = q + w$	First law of thermodynamics (internal energy change)
$\Delta H = H_{products} - H_{reactants}$	Enthalpy change for a reaction
$\Delta H_{rxn} = \sum \Delta H_f^{\circ} (products) - \sum \Delta H_f^{\circ} (reactants)$	Reaction enthalpy from standard enthalpies of formation

Additional info: Some tables and examples have been expanded for clarity and completeness. All equations are provided in LaTeX format as required for academic use.