Thermodynamics and Energy in Chemistry

Introduction to Energy

Energy is a fundamental concept in chemistry, describing the capacity to do work or transfer heat. In chemical systems, energy changes are central to understanding reactions and physical processes.

• Energy: The ability to do work (force acting over a distance, ) or to transfer heat (energy transfer due to temperature difference).

• Thermodynamics: The study of energy and its transformations.

• Thermochemistry: The branch of thermodynamics focusing on energy changes during chemical reactions, especially heat exchange.

Kinetic and Potential Energy

All forms of energy can be classified as either kinetic or potential energy.

• Kinetic Energy: Energy due to motion.

• Potential Energy: Stored energy due to position or arrangement.

• The most important form of potential energy in chemistry is electrostatic potential energy between charged particles.

Electrostatic Potential Energy

Electrostatic potential energy () describes the energy between charged particles:

• For two charges and separated by distance :

• is a proportionality constant: J·m/C2

• Unit of energy: Joule (J), where

• Like charges (repulsion):

• Opposite charges (attraction):

Example: The potential energy between two ions decreases (becomes more negative) as they approach each other if they are oppositely charged, indicating a stable, attractive interaction.

Energy Changes in Chemical Reactions

Chemical reactions involve making and breaking bonds, which are associated with energy changes:

• Bond Formation: Releases energy (exothermic, )

• Bond Breaking: Requires energy input (endothermic, )

First Law of Thermodynamics

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

• Mathematically:

• = heat exchanged, = work done

• Positive : System gains energy from surroundings

• Negative : System loses energy to surroundings

System and Surroundings

In thermodynamics, the universe is divided into the system (the part under study) and the surroundings (everything else).

• Open System: Can exchange both matter and energy with surroundings

• Closed System: Can exchange energy but not matter

• Isolated System: Cannot exchange either energy or matter

State Functions vs. Path Functions

Some thermodynamic quantities depend only on the current state of the system, not on how it got there.

• State Functions: Depend only on initial and final states (e.g., internal energy , enthalpy )

• Path Functions: Depend on the specific process or path taken (e.g., heat , work )

Internal Energy () and Its Change ()

The internal energy of a system is the sum of all kinetic and potential energies of its components. We can only measure changes in internal energy:

• If : System absorbs energy (endothermic)

• If : System releases energy (exothermic)

Heat (), Work (), and Their Signs

• q > 0: System absorbs heat (endothermic)

• q < 0: System releases heat (exothermic)

• w > 0: Work done on the system

• w < 0: Work done by the system

Enthalpy () and Enthalpy Change ()

Enthalpy is a thermodynamic quantity that is especially useful for processes at constant pressure:

At constant pressure, the change in enthalpy is:

• For most chemical reactions at constant pressure, equals the heat exchanged ():

• Endothermic process: (system absorbs heat)

• Exothermic process: (system releases heat)

Pressure-Volume Work

When a chemical reaction causes a change in volume (e.g., gas expansion), work is done:

• Work is negative when the system expands (does work on surroundings)

• Work is positive when the system is compressed (work done on system)

Summary Table: State vs. Path Functions

Examples and Applications

• Combustion reactions are typically exothermic, releasing heat to the surroundings.

• Photosynthesis is an endothermic process, requiring energy input from sunlight.

• Refrigerators move heat from inside (system) to outside (surroundings), not by creating 'cold' but by transferring energy.

Additional info: These notes expand on the provided slides by clarifying definitions, sign conventions, and including standard equations and examples for clarity and completeness.