Chemical Thermodynamics

1. Introduction

Thermodynamics is the study of the energy changes that accompany physical and chemical processes. Chemical thermodynamics applies these ideas to chemical reactions: how much heat is absorbed or released, and whether a reaction is spontaneous (feasible on its own). It answers "will it happen and how far", though not "how fast" (that is kinetics).

When hot water is left on a table it cools down: energy always flows from a hotter body to a colder one until they reach the same temperature. Thermodynamics is built to describe such energy flows quantitatively.

1.1 Some basic terms

• System: the part of the universe under study. Surroundings: everything else. Together they make the universe.
• Types of system: open (exchanges both matter and energy), closed (energy only), isolated (neither), e.g. a sealed vacuum flask.
• State function: a property that depends only on the present state, not the path taken to reach it, e.g. internal energy $U$, enthalpy $H$, entropy $S$, temperature, pressure, volume. Path functions depend on the route, e.g. heat $q$ and work $w$.
• Thermodynamic processes: isothermal (constant $T$), adiabatic (no heat exchange, $q = 0$), isobaric (constant $P$), isochoric (constant $V$).

2. Internal Energy ($U$) and Work

Internal energy $U$ is the total energy (kinetic + potential) stored in a system. Its absolute value cannot be measured, but the change $\Delta U$ can.

Work done by a gas expanding against external pressure $P$ (pressure-volume work):

$w = -P\,\Delta V$

(The minus sign follows the convention that work done by the system is negative because the system loses energy.)

3. First Law of Thermodynamics

Statement: Energy can neither be created nor destroyed, only converted from one form to another; the total energy of an isolated system is constant (law of conservation of energy).

Mathematically, for a closed system:

$\Delta U = q + w$

where $q$ is heat added to the system and $w$ is work done on the system. Substituting $w = -P\Delta V$:

$\Delta U = q - P\,\Delta V$

At constant volume ($\Delta V = 0$): $\Delta U = q_V$, so the heat absorbed at constant volume equals the change in internal energy.