Thermochemistry: Energy Changes in Chemical Reactions

Thermochemistry: Energy Changes in Chemical Reactions

Introduction to Thermochemistry

Thermochemistry is a branch of thermodynamics that focuses on the energy changes, particularly heat, involved in chemical reactions. Understanding these energy changes is essential for predicting reaction behavior and for practical applications such as energy production and biological processes.

• Energy is defined as the capacity to do work or transfer heat.

• Thermodynamics is the study of energy and its transformations.

• Thermochemistry specifically studies the energy changes (mainly heat) that accompany chemical reactions.

Chemical Reactions and Energy

Chemical reactions are a primary source of energy in daily life and biological systems.

• Examples include the combustion of gasoline, production of electricity from coal, heating homes with natural gas, and powering devices with batteries.

• Plants use solar energy for photosynthesis, storing energy in chemical bonds. Animals obtain energy by digesting these molecules.

5.1. The Nature of Chemical Energy

Energy in chemical systems can be classified as kinetic or potential energy. The energy associated with chemical reactions is mainly due to changes in potential energy, especially electrostatic potential energy between charged particles.

• Electrostatic potential energy ($E_{el}$) is given by:

$E_{el} = \frac{k Q_1 Q_2}{d}$

• Where $Q_1$ and $Q_2$ are the charges, $d$ is the distance between them, and $k$ is a proportionality constant ($8.99 \times 10^9$ J·m/C2).

• Energy is released when chemical bonds form and must be supplied to break bonds.

• Ionic bonds (e.g., in NaCl) and covalent bonds (e.g., in H2O, CH4) are both fundamentally electrostatic in nature.

5.2. The First Law of Thermodynamics

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

• Energy conversions are central to chemical reactions and practical applications (e.g., converting chemical energy to heat or electricity).

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

• Surroundings: Everything else outside the system.

Types of Systems

• Open system: Exchanges heat and mass with surroundings.

• Closed system: Exchanges heat but not mass.

• Isolated system: Exchanges neither heat nor mass.

Internal Energy ($E$)

The internal energy of a system is the sum of all kinetic and potential energies of its components. We usually measure changes in internal energy ($\Delta E$), not absolute values.

$\Delta E = E_{final} - E_{initial}$

• If $\Delta E < 0$, the system releases energy (exothermic).

• If $\Delta E > 0$, the system absorbs energy (endothermic).

• Energy can be exchanged as heat ($q$) or work ($w$):

$\Delta E = q + w$

• Sign conventions: $q > 0$ (heat absorbed), $w > 0$ (work done on system).

Quantity	Sign (+)	Sign (-)
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 ($q > 0$).

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

State Functions

A state function depends only on the current state of the system, not on the path taken to reach that state. Internal energy ($E$) is a state function, but heat ($q$) and work ($w$) are not.

5.3. Enthalpy ($H$)

Enthalpy is a thermodynamic quantity that reflects the heat content of a system at constant pressure.

• $H = E + PV$

• Change in enthalpy at constant pressure:

$\Delta H = \Delta E + P\Delta V$

• At constant pressure, $\Delta H = q_p$ (heat at constant pressure).

5.4. Enthalpy of Reaction ($\Delta H_{rxn}$)

The enthalpy change for a reaction is the difference between the enthalpy of products and reactants:

$\Delta H_{rxn} = H_{products} - H_{reactants}$

• Enthalpy is an extensive property (depends on amount of substance).

• Reversing a reaction changes the sign of $\Delta H$.

• Enthalpy change depends on the physical states of reactants and products.

5.5. Calorimetry

Calorimetry is the measurement of heat flow in a chemical reaction. The device used is called a calorimeter.

Heat Capacity and Specific Heat

• Heat capacity: Energy required to raise temperature of a substance by 1 K.

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

• Molar heat capacity: Energy required to raise 1 mol of a substance by 1 K.

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

• Where $q$ is heat, $C_s$ is specific heat, $m$ is mass, and $\Delta T$ is temperature change.

Constant Pressure Calorimetry

• Often performed in a "coffee-cup" calorimeter at atmospheric pressure.

• Heat change for the reaction is measured by the temperature change of water.

Constant Volume Calorimetry (Bomb Calorimeter)

• Used for reactions involving gases or combustion.

• Measures change in internal energy ($\Delta E$) at constant volume.

5.6. Hess's Law

Hess's Law states that the total enthalpy change for a reaction is the same, no matter how many steps the reaction is carried out in, because enthalpy is a state function.

• Allows calculation of $\Delta H$ for complex reactions by combining known reactions.

5.7. Enthalpies of Formation

The standard enthalpy of formation ($\Delta H_f^\circ$) is the enthalpy change when one mole of a compound is formed from its elements in their standard states (usually 1 atm, 25°C).

• Standard enthalpy change for a reaction:

$\Delta H^\circ_{rxn} = \sum n \Delta H_f^\circ (products) - \sum m \Delta H_f^\circ (reactants)$

• Where $n$ and $m$ are stoichiometric coefficients.

5.8. Bond Enthalpy

Bond enthalpy is the energy required to break one mole of a specific bond in a gaseous molecule. It is always positive (endothermic process).

• Energy is released when bonds form (exothermic).

• Reaction enthalpy can be estimated using bond enthalpies:

$\Delta H_{rxn} = \sum D_{bonds\,broken} - \sum D_{bonds\,formed}$

• If more energy is required to break bonds than is released in forming new ones, the reaction is endothermic; otherwise, it is exothermic.

5.9. Energy in Foods and Fuels

The energy released when one gram of food or fuel is combusted is called its fuel value. Most energy in foods comes from carbohydrates, fats, and proteins.

• Carbohydrates: ~17 kJ/g

• Fats: ~38 kJ/g

• Proteins: ~17 kJ/g

Alternative energy sources include solar, wind, geothermal, hydroelectric, and biomass.