Thermochemistry: Basic Concepts and Terminology

Chemical reactions are all about the transformation of matter. An integral part of that transformation is the flow and conversion of energy. The study of energy changes that accompany chemical reactions is called thermochemistry. It is a practical application of the broader field of thermodynamics, and it allows us to answer one of the most fundamental questions about a reaction: what is the quantity of energy transferred as heat when the reaction takes place?

Before we can answer that question, we must first establish a clear framework and a precise vocabulary for discussing energy.

Defining the Universe: System and Surroundings

To track the flow of energy, we must first define exactly what we are observing. In thermodynamics, we partition the universe into two distinct parts: the system and the surroundings.

The system is the specific part of the universe we are interested in studying. In chemistry, this is usually the collection of reactants and products in a chemical reaction.
The surroundings are everything else in the universe outside of the system. This could be the beaker the reaction is in, the lab bench, the air in the room, and everything beyond.

The universe is simply the sum of the system and the surroundings. \[ \text{Universe} = \text{System} + \text{Surroundings} \] The relationship between the system and its surroundings is the key to all of thermochemistry. We are interested in the energy of the system, but we measure it by observing the effects of energy exchange on the surroundings.

A Fundamental Convention: The System’s Point of View

To build a consistent and unambiguous framework, thermodynamics adopts a universal convention: All energy changes are described from the point of view of the system.

Think of the system as the protagonist of our story. When we discuss energy transfers like heat (q) or work (w), we are always asking, “Did the system’s energy go up or down?”

If energy flows into the system, the change is positive.
If energy flows out of the system, the change is negative.

This system-centric viewpoint is the bedrock of all thermodynamic equations. When you see a variable like q or w without a subscript, it is always referring to the change for the system (q_sys or w_sys). This is not just a common practice; it is the essential rule that allows for a universal language of energy accounting.

Occasionally, we may need to discuss the energy change of the surroundings. Only in these specific cases will a variable be given an explicit subscript, such as q_surr. This makes our accounting clear: q always means the system, and any other perspective must be explicitly labeled.

Defining the System: A Question of Perspective

In thermodynamics, the system is the specific part of the universe we are interested in studying. The crucial idea is that we, the observers, get to define the system’s boundary. The way we define this boundary depends entirely on the question we are trying to answer.

Let’s consider a familiar process: a car’s internal combustion engine.

The Chemist’s Perspective

The Goal: To understand the fundamental energy released by the fuel.
The System: The mixture of gasoline and air molecules inside the cylinder.
The Analysis: A chemist wants to know the enthalpy of combustion (Δ_cH). They would measure the heat that flows out of the reacting molecules (the system) into the engine block (the surroundings). From this perspective, the combustion is a highly exothermic process, releasing a specific amount of energy per mole of gasoline. This is a fundamental chemical property of the fuel.

The Mechanical Engineer’s Perspective

The Goal: To determine the engine’s efficiency and power output.
The System: The entire engine block, including the pistons and cylinders.
The Analysis: An engineer cares about how much of the chemical energy is converted into useful work (pushing the pistons) versus how much is “wasted” as heat that must be carried away by the cooling system. They define the whole engine as the system and measure the work done by the system on the crankshaft and the heat transferred out of the system into the radiator and exhaust. Their goal is to maximize work and minimize waste heat.

The Environmental Scientist’s Perspective

The Goal: To assess the car’s impact on the atmosphere.
The System: The entire car, including the engine, fuel tank, and exhaust pipe.
The Analysis: An environmental scientist defines the entire car as the system and studies what crosses the boundary into the surroundings (the atmosphere). They are not just interested in heat, but in the mass transfer of products like CO₂, H₂O, and pollutants like NO_x. Their analysis focuses on the chemical composition of what the system outputs into the environment.

The way you draw the boundary—around the molecules, the engine, or the entire car—changes what you measure and the questions you can answer. This flexibility allows thermodynamics to be applied to systems ranging from a single chemical bond to the entire global climate.

Classifying Systems by Boundary

To properly account for the flow of energy and matter, it is useful to classify a system based on the nature of its boundary.

An open system can exchange both energy and matter with its surroundings. (Example: An open beaker of water).
A closed system can exchange energy but not matter with its surroundings. (Example: A sealed flask containing a reaction mixture. This is the most common type of system we will study).
An isolated system cannot exchange either energy or matter with its surroundings. (Example: An ideal, perfectly insulated thermos).

The Energy of a System: Internal Energy

Every chemical system contains a certain amount of energy. The sum of all the kinetic and potential energies of all the particles (atoms, molecules, ions) within that system is called its internal energy (U).

Kinetic Energy: This includes the movement of molecules from place to place (translation), the rotation of molecules, and the vibration of atoms within molecules.
Potential Energy: This is the energy stored in chemical bonds and the intermolecular forces between particles.

We can think of the internal energy as the system’s “energy bank account.” It is impossible to know the absolute total amount of energy in this account at any given moment. However, we can precisely measure the change in internal energy (ΔU) that occurs during a chemical or physical process. A change is always calculated as the final state minus the initial state.

\[ \Delta U = U_{\mathrm{final}} - U_{\mathrm{initial}} \]

How Energy is Transferred: Heat and Work

Energy can be transferred between the system and the surroundings, but it cannot be created or destroyed. This is the Law of Conservation of Energy, which is the foundational principle of thermochemistry.

There are only two ways that a system can exchange energy with its surroundings: heat and work. These two concepts represent the two fundamental modes of energy transfer at the molecular level.

Heat (q) is the transfer of thermal energy between the system and the surroundings. Thermal energy is the energy associated with the random, disorganized motion of atoms and molecules.
Work (w) is the transfer of energy that results in organized, uniform motion. A common example in chemistry is the work done by a gas as it expands, moving the molecules of its surroundings in a uniform direction.

Heat and work are not forms of energy themselves, but rather processes of energy transfer. They are the only two mechanisms by which a system can exchange energy with its surroundings, changing its internal energy balance. The relationship between a system’s internal energy and the heat and work transferred across its boundary is elegantly summarized by one of the most important laws in all of science: the First Law of Thermodynamics.

A Point of Precision: The Scientific Meaning of ‘Heat’

In everyday language, we use the word “heat” loosely, often as a synonym for “hotness” or “thermal energy”. We might say, “This coffee has a lot of heat”.

In thermodynamics, however, the word heat (q) has a very precise and different meaning. It is not a substance or a property that an object has or contains. Instead, heat is a process: it is the transfer of thermal energy from one object to another.

A perfect analogy is the distinction between rain and water.

A cloud does not contain rain; it contains water vapor.
Rain is the process of that water falling.
We can measure the amount of water transferred during that process (e.g., “one inch of rainfall”).

The same is true for heat:

An object does not contain heat; it contains thermal energy.
Heat is the process of that thermal energy being transferred.
We can measure the quantity of energy transferred during that process (the value of q).

Why This Distinction Matters: This is not just a matter of semantics; it is fundamental to understanding the First Law of Thermodynamics (ΔU = q + w). This equation describes the transfer of energy across a system’s boundary. The terms q and w represent quantities of energy that have been “deposited” or “withdrawn,” changing the system’s internal energy, U. If you think of heat as something a system already possesses, the equation becomes conceptually nonsensical.

In this course, we will be precise. A substance contains thermal energy. Heat is the transfer of that thermal energy.