Difficulty: Introductory to Intermediate | Prerequisites: Basic chemistry concepts, familiarity with energy and matter.
Big Picture
Thermochemistry sits at the intersection of chemistry and physics, asking one central question: where does the energy go during a chemical or physical change? This section lays the groundwork by defining what counts as your "system" versus everything else (the surroundings), introducing the first law of thermodynamics, and distinguishing the different types of processes a system can undergo. You also learn the difference between state functions (properties that depend only on the current state) and process functions (which depend on the path taken). All of thermochemistry builds on these foundations.
Thermochemistry tracks energy changes during chemical and physical processes. Every problem starts by defining a system (the matter you are watching) and its surroundings (everything else), then applying the first law of thermodynamics: energy is conserved, so any change in a system's internal energy equals the heat added minus the work it does. State functions like enthalpy and entropy depend only on where you start and finish, not on the route you take.
System
The matter being observed in a process, such as all reactants and products in a reaction, or the gas inside a balloon. Think of it as the part of the universe you have drawn a boundary around.
Surroundings (environment)
Everything outside the system. If the system is the gas in a piston, the surroundings include the piston walls, the room, and the rest of the universe.
Isolated system
A system that cannot exchange energy or matter with its surroundings. In simple terms, nothing gets in or out. Example: an insulated bomb calorimeter.
Closed system
A system that can exchange energy but not matter with its surroundings. Energy crosses the boundary, but atoms and molecules stay put. Example: a steam radiator.
Open system
A system that can exchange both energy and matter with its surroundings. Think of it as the least restrictive type. Example: a pot of boiling water on a stove.
Internal energy (U)
The total energy contained within a system, including kinetic and potential energy of all particles. It is a state function.
Isothermal process
A process that occurs at constant temperature, meaning the internal energy does not change (ΔU = 0). In simple terms, any heat added is entirely converted to work.
Adiabatic process
A process where no heat is exchanged between the system and surroundings (Q = 0). The change in internal energy equals the negative of the work done.
Isobaric process
A process that occurs at constant pressure. Think of it as what happens inside an open container exposed to the atmosphere.
Isovolumetric (isochoric) process
A process where the volume does not change, so no work is performed (W = 0). All energy transfer appears as heat.
Spontaneous process
A process that can occur without continuous external energy input. This does not mean it happens quickly or goes to completion.
Nonspontaneous process
A process that requires a continuous supply of energy to proceed. It can sometimes be driven by coupling it to a spontaneous reaction.
State function
A property that depends only on the current state of the system, not on the path taken to get there. The mnemonic PTV HUGS D covers: Pressure, Temperature, Volume, entHalpy (H), internal energy (U), Gibbs free energy (G), entropy (S), and Density.
Process function
A quantity that depends on the pathway taken between two states. Heat (Q) and work (W) are the main examples.
Standard conditions
25 °C (298 K), 1 atm pressure, 1 M concentration for solutions. Used for kinetics, equilibrium, and thermodynamics calculations. In simple terms, the "default settings" for chemistry problems.
Standard temperature and pressure (STP)
0 °C (273 K) and 1 atm. Used specifically for ideal gas calculations. Students often confuse STP with standard conditions; they are different.
The system is whatever matter you are studying: all reactants and products, the solute and solvent in a solution, or the gas inside a container.
The surroundings are everything else, extending in principle to the rest of the universe.
Every thermochemistry problem begins by drawing this boundary clearly.
Isolated system: no exchange of energy or matter. Example: insulated bomb calorimeter.
Closed system: energy can cross the boundary, matter cannot. Example: steam radiator.
Open system: both energy and matter can cross the boundary. Example: pot of boiling water.
Energy cannot be created or destroyed, only transferred or converted.
Expressed as ΔU = Q − W, where U is internal energy, Q is heat added to the system, and W is work done by the system.
This equation is the starting point for every process type below.
Isothermal (constant temperature):
ΔU = 0, so Q = W.
Appears as a hyperbolic curve on a P-V diagram. Work equals the area under the curve.
Adiabatic (no heat exchange):
Q = 0, so ΔU = −W.
Also a hyperbolic curve on P-V, steeper than the isothermal curve.
Isobaric (constant pressure):
The first law still applies without modification.
Appears as a horizontal (flat) line on a P-V diagram.
Isovolumetric / isochoric (constant volume):
W = 0 because no expansion or compression occurs.
ΔU = Q: all energy transfer is heat.
Appears as a vertical line on a P-V diagram, with no area under the curve.
A spontaneous process proceeds without continuous external energy, but that says nothing about speed. It may be slow, partial, or temperature-dependent.
Most spontaneous processes are rare in practice because they require overcoming a high activation energy.
A nonspontaneous process needs energy input. Coupling it with a spontaneous reaction can supply the necessary energy.
State functions depend only on the current state, not the path. The mnemonic PTV HUGS D covers: Pressure, Temperature, Volume, entHalpy (H), internal energy (U), Gibbs free energy (G), entropy (S), Density.
These are independent of the process taken to reach that state, but they are not independent of each other.
They describe the system at equilibrium.
Process functions depend on the pathway: heat (Q) and work (W).
Standard conditions: 25 °C (298 K), 1 atm, 1 M solutions. Used for thermodynamics, kinetics, and equilibrium.
STP (standard temperature and pressure): 0 °C (273 K), 1 atm. Used for ideal gas calculations.
Under standard conditions, the most stable form of a substance is its standard state:
H₂ is a gas, H₂O is a liquid, NaCl is a solid, O₂ is a gas, C is a solid.
Standard enthalpy (ΔH°), entropy (ΔS°), and free energy (ΔG°) are all measured under standard conditions.
\Delta U = Q - WFirst law of thermodynamics. ΔU is change in internal energy, Q is heat added to the system, W is work done by the system.
Isothermal: ΔU = 0, so W = Q.
Adiabatic: Q = 0, so ΔU = −W.
Isovolumetric: W = 0, so ΔU = Q.
The first law of thermodynamics is the principle behind every engine, refrigerator, and power plant: energy in minus work out equals the change in the system's stored energy. Car engines run roughly isobaric combustion (constant atmospheric pressure pushes the piston), and a thermos flask is an everyday approximation of an isolated system, keeping your coffee hot by minimising energy exchange with the surroundings.
Students often think "spontaneous" means "fast." It does not. A spontaneous reaction can be imperceptibly slow if its activation energy is high. Diamond converting to graphite is spontaneous, but you will never watch it happen.
Confusing standard conditions (25 °C, 298 K) with STP (0 °C, 273 K) is one of the most common errors on exams. They are used in different contexts.
Students sometimes treat heat and work as state functions. They are not. Both depend on the path the process takes, which is why they are called process functions.
Assuming an isolated system means "nothing happens inside it" is wrong. Reactions can and do proceed inside isolated systems; the constraint is that nothing crosses the boundary.
⚠️ Be able to identify a system as isolated, closed, or open from a description. Exam questions love giving a scenario and asking which type it is.
⚠️ Know the first law cold: ΔU = Q − W. You will rearrange it for every process type.
⚠️ Know which variable is zero for each process type: isothermal (ΔU = 0), adiabatic (Q = 0), isovolumetric (W = 0). This is a near-guaranteed exam question.
⚠️ Distinguish state functions from process functions. If an exam asks "which of the following depends on the path," the answer is heat or work.
⚠️ Standard conditions vs STP: know which is which and when each applies.
True or False: In an adiabatic process, heat is exchanged between the system and surroundings.
Answer: False. Q = 0 in an adiabatic process.
Fill in the blank: In an isovolumetric process, W = ___, so ΔU = ___.
Answer: W = 0, ΔU = Q.
True or False: "Spontaneous" means the reaction happens quickly.
Answer: False. Spontaneous refers to thermodynamic favourability, not speed.
Fill in the blank: Standard conditions are ___ K and ___ atm, while STP is ___ K and ___ atm.
Answer: 298 K and 1 atm; 273 K and 1 atm.
True or False: Enthalpy is a process function.
Answer: False. Enthalpy is a state function.
Q: A gas expands inside a perfectly insulated cylinder. No heat enters or leaves the system. What type of process is this, and what happens to the internal energy if the gas does 150 J of work?
A: This is an adiabatic process (Q = 0). From ΔU = Q − W: ΔU = 0 − 150 = −150 J. The internal energy decreases by 150 J.
Q: A sealed, rigid steel container holds a reacting gas mixture. Classify the system and the process by volume, then state which variable in the first law equals zero.
A: The system is closed (energy can transfer through the walls, but matter cannot leave). The process is isovolumetric (rigid container, no volume change), so W = 0.
Q: Why is heat considered a process function rather than a state function?
A: Heat depends on the path taken between two states. Two different processes connecting the same initial and final states can transfer different amounts of heat, even though the change in state functions (like ΔU) is the same.
Q: A student claims that STP and standard conditions are the same thing. Explain the error.
A: STP is 0 °C (273 K) and 1 atm, used for ideal gas calculations. Standard conditions are 25 °C (298 K), 1 atm, and 1 M solutions, used for thermodynamics, kinetics, and equilibrium. The temperatures differ by 25 degrees.
Q: List the state functions covered by the mnemonic PTV HUGS D.
A: Pressure, Temperature, Volume, entHalpy (H), internal energy (U), Gibbs free energy (G), entropy (S), and Density.
This material connects directly to enthalpy, entropy, and Gibbs free energy (covered in Part 2 and Part 3 of these notes), because all three are state functions introduced here. The first law of thermodynamics reappears in calorimetry problems, where you measure Q to find ΔH. Phase changes (also Part 2) are classified by process type: boiling at atmospheric pressure is isobaric, and bomb calorimetry is isovolumetric. Understanding system boundaries is also essential for general physics (engines, refrigeration cycles) and biology (metabolic pathways as open systems).
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