Thermodynamics Concepts in Chemistry 1412

Thermodynamics

The study of energy and its transformations, especially during physical and chemical changes.

System

The part of the universe we're studying (e.g., chemical reaction).

Surroundings

Everything else (e.g., container, air).

Open System

Exchange of mass and energy.

Closed System

Exchange of energy only.

Isolated System

No exchange of mass or energy.

First Law of Thermodynamics

Energy cannot be created or destroyed, only transformed.

Equation for Internal Energy Change

ΔE = q + w

ΔE

Change in internal energy.

q

Heat (positive if gained by system).

w

Work (positive if done on the system).

Endothermic

q > 0 → absorbs heat.

Exothermic

q < 0 → releases heat.

Work Done on System

w > 0 → work done on system (compression).

Work Done by System

w < 0 → work done by system (expansion).

Internal Energy (E)

Sum of kinetic and potential energy of all particles in the system; a state function.

Work Equation

w = -PΔV

P

External pressure.

ΔV

Change in volume.

Enthalpy (ΔH)

Heat change at constant pressure.

Standard Enthalpy Change (ΔH°)

Change when reactants and products are in their standard states (1 atm, 25°C).

Enthalpy of Formation (ΔHf°)

Heat change when 1 mole of a compound is formed from elements in standard states.

Enthalpy of Reaction (ΔHrxn)

ΔHrxn = ΣnΔHf°(products) - ΣnΔHf°(reactants).

Hess's Law

If a reaction occurs in multiple steps, total ΔH is the sum of ΔH of the steps.

Calorimetry

Used to measure heat flow.

Calorimetry Equation

q = mcΔT

m

Mass (g).

c

Specific heat capacity (J/g·°C).

ΔT

Change in temperature.

Entropy (S)

Measure of disorder or randomness.

Units of Entropy

J/mol·K.

Second Law of Thermodynamics

For any spontaneous process, the total entropy of the universe increases: ΔSuniverse = ΔSsystem + ΔSsurroundings > 0.

Qualitative Entropy Trends

Solid → liquid → gas → ΔS increases.

Gibbs Free Energy (ΔG)

The energy available to do work.

Gibbs Free Energy Equation

ΔG = ΔH - TΔS.

Temperature in Gibbs Equation

T = temperature in Kelvin.

Spontaneity Conditions

ΔG < 0 → spontaneous; ΔG = 0 → equilibrium; ΔG > 0 → non-spontaneous.

Standard Free Energy (ΔG°)

Standard Conditions: 1 atm, 25°C, 1 M concentrations.

Relationship with Equilibrium Constant (K)

ΔG° = -RT lnK.

R

8.314 J/mol·K.

Implications of ΔG°

ΔG° < 0 → K > 1 → products favored; ΔG° > 0 → K < 1 → reactants favored.