Chapter 5: Thermochemistry – Study Notes

Thermochemistry

Introduction to Energy and Thermochemistry

Thermochemistry is a branch of thermodynamics focused on the study of energy changes, particularly heat, that accompany chemical reactions. Understanding energy transformations is essential for predicting reaction behavior and energy flow in chemical systems.

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

• Thermodynamics: The study of energy and its transformations.

• Thermochemistry: The study of chemical reactions and the energy changes involving heat.

5.1 The Nature of Chemical Energy

Chemical energy is a form of potential energy stored in the arrangement of atoms and molecules. The most important type in charged particles is electrostatic potential energy.

• Electrostatic Potential Energy ($E_{el}$): Energy due to the attraction or repulsion between charged particles. $E_{el} = \frac{kQ_1Q_2}{d}$ where $k$ is a constant, $Q_1$ and $Q_2$ are charges, and $d$ is the distance between them.

• Unit of Energy: The Joule (J) $1\,J = 1\,kg\,m^2/s^2$

Attraction Between Ions

Electrostatic attraction occurs between oppositely charged ions, forming ionic bonds.

• Energy is released when chemical bonds are formed ($E_{el} < 0$).

• Energy is consumed when chemical bonds are broken ($E_{el} > 0$).

5.2 First Law of Thermodynamics

The first law states that energy can be converted from one form to another, but it cannot be created or destroyed.

• Examples: Chemical energy converted to heat (heating homes), sunlight to chemical energy (photosynthesis).

Definitions: System and Surroundings

In thermochemistry, it is crucial to define the system and its surroundings.

• System: The part of the universe singled out for study (e.g., reactants in a chemical reaction).

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

Types of Systems

• Open System: Can exchange both heat and mass with surroundings.

• Closed System: Can exchange heat but not mass.

• Isolated System: Cannot exchange heat or mass.

Internal Energy ($E$)

The internal energy of a system is the sum of all kinetic and potential energies of its components.

• Symbol: $E$

• Change in internal energy: $\Delta E$

• We usually know only the change, not the absolute value.

Change in Internal Energy

• If $\Delta E < 0$, the system releases energy to the surroundings.

• If $\Delta E > 0$, the system absorbs energy from the surroundings.

• Formula: $\Delta E = E_{final} - E_{initial}$

• $\Delta E$ is a state function: depends only on initial and final states, not the path.

Thermodynamic Quantities

• Each quantity has three parts: a number, a unit, and a sign.

• Sign conventions:

  • Positive $\Delta E$: system gains energy.

  • Negative $\Delta E$: system loses energy.

Relating $\Delta E$ to Heat and Work

• Energy exchanged as heat ($q$) or work ($w$).

• Formula: $\Delta E = q + w$

Sign Conventions Table

Quantity	Positive (+)	Negative (-)
q (heat)	System gains heat	System loses heat
w (work)	Work done on system	Work done by system
$\Delta E$	Net gain of energy	Net loss of energy

Endothermic and Exothermic Processes

• Endothermic: System absorbs heat ($q > 0$), temperature drops.

• Exothermic: System releases heat ($q < 0$), temperature rises.

State Functions

• State function: Property dependent only on the current state, not the path (e.g., $E$, $\Delta E$).

• Path functions: Properties dependent on the path (e.g., $q$, $w$).

5.3 Enthalpy ($H$)

Enthalpy is a thermodynamic quantity that accounts for heat flow at constant pressure.

• Definition: $H = E + PV$

• Change in enthalpy: $\Delta H = \Delta E + P\Delta V$

• At constant pressure: $\Delta H = q$

Pressure-Volume Work

• Work done by a gas: $w = -P\Delta V$

• Negative sign indicates work done by the system.

Enthalpy Change and Heat

• Endothermic process: $\Delta H$ is positive.

• Exothermic process: $\Delta H$ is negative.

Heat of Reaction

• Enthalpy of reaction ($\Delta H_{rxn}$): The heat change for a chemical reaction at constant pressure.

• Example: $\Delta H = -483.6\,kJ$ for the formation of water from hydrogen and oxygen.

5.4 Enthalpies of Reaction

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

• Reverse reaction: Same magnitude, opposite sign.

Enthalpy Guidelines

1. Enthalpy is an extensive property: depends on the amount of substance.

2. Enthalpy change for a reaction is equal in magnitude but opposite in sign for the reverse reaction.

3. Enthalpy change depends on the states (phases) of reactants and products.

5.5 Calorimetry

Calorimetry is the measurement of heat flow in chemical reactions.

• Calorimeter: Instrument used to measure heat flow.

Heat Capacity and Specific Heat

• Heat capacity (C): Energy required to raise temperature by 1 K (1°C).

• Specific heat (C_s): Energy required to raise temperature of 1 g by 1 K.

• Molar heat capacity (C_m): Energy required to raise temperature of 1 mol by 1 K.

• Formula: $q = C \times m \times \Delta T$

Substance	Specific Heat (J/g·K)
H2O (l)	4.18
Al (s)	0.90
Fe (s)	0.45
CaCO3 (s)	0.82

Constant-Pressure Calorimetry

• Uses a coffee-cup calorimeter.

• Specific heat of water: 4.184 J/g·K.

• Formula: $q_{soln} = C_s \times m_{soln} \times \Delta T = -q_{rxn}$

Bomb Calorimetry (Constant Volume)

• Uses a sealed bomb calorimeter.

• Measures change in internal energy ($\Delta E$), not enthalpy ($\Delta H$).

• Formula: $q_{rxn} = -C_{cal} \times \Delta T$

Sample Exercises

• Calculate $\Delta E$ for a process where a system absorbs 140 J of heat and does 85 J of work: $\Delta E = q + w$

• Calculate heat absorbed by rocks: $q = C_s \times m \times \Delta T$

Additional info:

• These notes cover the core principles of thermochemistry, including energy, enthalpy, calorimetry, and the first law of thermodynamics, as relevant to a General Chemistry college course.