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Summary of equations

This article briefly summarises the energy and enthalpy equations and concepts covered so far.

The First Law of Thermodynamics

The first law of thermodynamics means that energy is conserved. Any change in internal energy (U) is accounted for by an exchange of heat (q) and work done by or to the system (w).

[Delta U = q + w]

What is Heat?

Heat is the energy exchanged by surroundings through microscopic collisions. It is path dependent – the heat exchanged between a system and surroundings depends on the path the system takes.

What is Work?

Energy is the ability to do work. Work, w, is defined as either a change in volume (V) multiplied by pressure (p) resisting it, or a change in distance (h) multiplied by the force (F) resisting it. It is negative because expansion lowers the energy of the system:

[w = -p times Delta V = -F times Delta h]

What is Enthalpy?

Enthalpy is defined as the heat energy (q) released or absorbed by a system at constant pressure. This is convenient for chemistry, and is the main definition of heat energy we use. Enthalpy (H) is the sum of internal energy (U) and the pressure-volume work (pV) needed for the system to establish itself.

[Delta H = Delta U + p Delta V]

How do we Measure Enthalpy?

Enthalpy (H) and/or heat transferred (q) are not the same as temperature (T). They are proportional, and the proportionality constant is known as the heat capacity (Cp). We can determine heat exchanged between a system and surroundings by multiplying the heat capacity, temperature change, and amount of material – as heat capacity is usually given “per mass”.

[q = -C_p times m times Delta T]

These are the key concepts and equations from this section.