Thermochemistry: Energy, Work, and Heat in Chemical Systems

Energy Basics

Thermochemistry is the study of energy changes, particularly heat, that accompany chemical reactions and physical changes. Understanding the different forms of energy and their interconversions is fundamental to analyzing chemical processes.

• Energy: The capacity to do work. In chemistry, work is a directed energy change resulting from a chemical process.

• Work: Defined as force multiplied by distance ().

• Kinetic Energy (KE): Energy of motion. Important in understanding the movement of molecules and atoms.

• Potential Energy (PE): Energy due to position or composition. For example, water behind a dam or a rock at the top of a hill.

• Radiant Energy: Energy from electromagnetic radiation, such as sunlight, which drives photosynthesis and affects climate.

• Thermal Energy: Associated with the random motion of atoms and molecules. It is related to temperature, but also depends on the amount of substance present.

• Chemical Energy: Stored within the structural units of chemical substances, determined by the arrangement of atoms in molecules.

Energy Conversion: All forms of energy can be converted from one form to another, but the total energy remains constant (law of conservation of energy).

The First Law of Thermodynamics

The First Law of Thermodynamics states that energy cannot be created or destroyed, only converted from one form to another. The total energy of the universe is constant.

• If a system gives up heat to the surroundings (exothermic, ), or does work on the surroundings (), its internal energy decreases.

• If heat is added to the system (endothermic, ), or work is done on the system (), its internal energy increases.

Work in Chemical Systems

Work is often associated with the expansion or compression of gases in chemical and biological processes.

• Pressure-Volume Work: When a gas expands or is compressed in a cylinder with a piston, the work done is given by:

• (change in volume)

• Units: 1 L·atm = 101.3 J

• If (expansion), is negative (work done by the system).

• If (compression), is positive (work done on the system).

Example: Inflating a balloon from 0.100 L to 1.85 L against 1.00 atm:

Internal Energy Change ()

The change in internal energy of a system is the sum of the heat exchanged and the work done:

• = heat exchanged (positive if absorbed by the system, negative if released)

• = work done (positive if done on the system, negative if done by the system)

Example: Burning fuel in a piston: Volume expands from 0.255 L to 1.45 L against 1.02 atm, and 875 J is emitted as heat.

This process is exothermic (energy is released).

State Functions vs. Path Functions

• State Function: A property that depends only on the current state of the system, not on the path taken (e.g., internal energy, volume, pressure).

• Path Function: A property that depends on the specific process or path taken (e.g., work, heat).

Example: Heating 100 g of water from 20°C to 30°C requires 4184 J, regardless of the heat source. Heat is not a state function because its value depends on the path.

Specific Heat and Heat Capacity

Understanding how substances absorb or release heat is essential in thermochemistry.

• Specific Heat (c): The amount of heat required to raise the temperature of 1 gram of a substance by 1°C.

• Heat Capacity (C): The amount of heat required to raise the temperature of a given quantity of a substance by 1°C.

• The relationship: , where is the mass in grams.

The amount of heat () absorbed or released can be calculated as:

• Example: How much heat is released when a 21.2 g block of iron cools from 91.6°C to 25.1°C? ( for Fe)

(The negative sign indicates an exothermic process.)

• Example: What is the final temperature when 1.2 g of iron absorbs 11.1 J, starting at 20.0°C? (Endothermic process, )

• Example: A 14.2 g ring of 18K gold requires 22.1 J to heat from 16.7°C to 28.8°C. Calculate its specific heat.

Table: Specific Heats of Selected Substances

Summary of Key Equations

• Work:

• Internal Energy Change:

• Heat (using specific heat):

• Heat Capacity:

Additional info:

• Endothermic processes absorb energy (q > 0), while exothermic processes release energy (q < 0).

• State functions are properties that depend only on the state of the system, not the path taken to reach that state.

• Path functions, such as heat and work, depend on the process taken between two states.