Work, Heat, Energy — Definition

Definition

Imagine you have a chemical reaction happening inside a beaker – that beaker and its contents form your 'system' in chemistry. Everything outside the beaker, like the air around it or your hand holding it, is the 'surroundings'. Now, let's talk about what's going on with energy in this system.

Energy is the capacity to do work or produce heat. In chemistry, when we talk about the total energy stored within a system, we're usually referring to its **Internal Energy ($U$)**. Think of internal energy as the sum of all the kinetic energies (energy of motion) and potential energies (stored energy due to position or arrangement) of every single atom, molecule, and subatomic particle inside your beaker.

This includes the movement of molecules, the vibrations within molecules, the rotations of molecules, and even the energy stored in chemical bonds. What's crucial about internal energy is that it's a 'state function' – its value depends only on the current state of the system (like its temperature, pressure, and volume), not on how it got to that state.

**Heat ($q$)** is a way energy moves between your system and its surroundings because there's a temperature difference. If your reaction in the beaker gets hot, heat will flow out of the beaker into the cooler air.

If the reaction gets cold, heat will flow from the warmer air into the beaker. It's like a natural flow from a hotter region to a colder region. We use specific sign conventions: if the system absorbs heat (gets hotter), $q$ is positive ($+q$).

If the system releases heat (gets colder), $q$ is negative ($-q$). Heat is a 'path function,' meaning the amount of heat transferred depends on the specific way the process happens, not just the initial and final states.

**Work ($w$)** is another way energy moves between your system and its surroundings, but this time it's not due to a temperature difference. In chemistry, the most common type of work we encounter is 'pressure-volume work' or 'expansion work.

' Imagine your reaction produces a gas. This gas will push against the atmosphere or a piston, causing it to move. When the system (the gas) pushes outwards and expands, it's doing work *on* the surroundings.

Conversely, if the surroundings push inwards and compress the gas, work is being done *on* the system. Like heat, work is also a 'path function' – the amount of work done depends on the specific process.

The sign convention for work is: if the system does work on the surroundings (e.g., expands), $w$ is negative ($-w$). If the surroundings do work on the system (e.g., compress), $w$ is positive ($+w$).

The First Law of Thermodynamics ties these three together: it states that the change in the internal energy of a system ($Delta U$) is equal to the heat added to the system ($q$) plus the work done on the system ($w$). Mathematically, this is expressed as $Delta U = q + w$. This law is essentially a statement of the conservation of energy – energy cannot be created or destroyed, only transferred or transformed.