Difficulty: Introductory to Intermediate
Prerequisites: Basic algebra, familiarity with the mole concept and balancing chemical equations. If you are shaky on unit conversions (especially joules, kilojoules, litres, atmospheres), review those first.
Thermochemistry sits at the intersection of chemistry and physics: it asks where the energy goes when bonds break and form. Everything that follows in chemical kinetics, equilibrium and electrochemistry leans on the language you learn here, so the payoff for getting it right is large. You need to be comfortable with the idea that energy is conserved, that heat flows from hot to cold, and that the sign of a number tells you the direction of that flow. If you missed the lectures on chemical bonding, go back and review them; this chapter assumes you know what a bond is.
Energy is the capacity to do work or produce heat, and it is always conserved. Chemical reactions either release energy to the surroundings (exothermic, feels hot) or absorb energy from the surroundings (endothermic, feels cold). The internal energy change of a system equals the heat added plus the work done on it: Delta E = q + w.
Energy
The capacity to do work (w) or to produce heat (q). Energy in the universe is conserved: it can change form but cannot be created or destroyed.
In simple terms, energy is what makes things happen, whether that is moving an object or warming it up.
Potential energy
Stored energy associated with the position or composition of an object.
Think of it as energy waiting to be released. Water behind a dam, petrol in a parked car, a ball at the top of a hill and the chemical bonds in a reactant are all examples.
Kinetic energy (KE)
The energy of motion, dependent on mass and velocity: KE = 1/2 mv squared.
In simple terms, anything that is moving has kinetic energy. The heavier it is or the faster it moves, the more it has.
Temperature
A measure of the random motion of the particles in a substance. It tells you how hot or cold something is, but it is not the same as the total energy stored in that substance.
Heat (q)
The transfer of energy between two objects due to a temperature difference. Heat flows spontaneously from a hotter object to a cooler one. Heat is not a substance contained by an object; it is a process of energy transfer.
Think of it as energy in transit. Once it arrives, it becomes part of the internal energy of the receiving object.
System
The part of the universe you are focusing on. In chemistry this is typically the reaction itself: the reactants and products.
Surroundings
Everything else in the universe that is not the system. In a lab setting, this includes the reaction vessel, the bench, the air, and you.
Exothermic reaction
A reaction that releases energy (as heat) to the surroundings. The system loses energy, the surroundings gain it. Exothermic reactions feel hot to the touch. The sign of q is negative (q < 0) from the system's point of view.
Endothermic reaction
A reaction that absorbs energy (as heat) from the surroundings. The system gains energy, the surroundings lose it. Endothermic reactions feel cold to the touch. The sign of q is positive (q > 0) from the system's point of view.
Thermodynamics
The study of energy and its interconversions between different forms.
First Law of Thermodynamics
The total energy of the universe is constant. Energy can be transferred or converted, but the total never changes. This is another way of stating the law of conservation of energy.
Internal energy (E)
The sum of all kinetic and potential energies of every particle in the system. You cannot measure E directly, but you can measure the change in internal energy, Delta E.
Work (w)
In chemical systems, work is usually the energy associated with gas expansion or compression against an external pressure. Calculated as w = -P times Delta V.
Energy is defined as the capacity to do work (w) or to produce heat (q).
The concept is familiar but surprisingly hard to pin down precisely.
The key principle: energy in the universe is conserved. It can be converted from one form to another, but it can never be created or destroyed (this is the First Law of Thermodynamics).
Potential energy: stored energy, linked to position or chemical composition.
Water behind a dam (gravitational potential energy)
Petrol in a parked car (chemical potential energy)
A ball at the top of a hill
Chemical bonds in the reactants of a reaction
Kinetic energy: the energy of motion, determined by mass and velocity.
KE = 1/2 mv squared
A heavier or faster-moving object has more kinetic energy.
Temperature measures the average random motion of particles in a substance. It is an intensive property (does not depend on how much substance you have).
Heat (q) is the transfer of energy between objects caused by a temperature difference.
Heat flows spontaneously from hot to cold.
Heat is not a substance; it is energy in transit.
Once transferred, the energy becomes part of the receiving object's internal energy.
System vs surroundings: the system is the reaction you are studying. The surroundings are everything else.
Example: for CH4(g) + 2O2(g) -> CO2(g) + 2H2O(g) + energy, the system is the reactants and products, and the surroundings are the container, the lab, the air.
Exothermic reactions release energy to the surroundings.
Some of the potential energy stored in the chemical bonds is converted to thermal energy.
Energy flows out of the system, so q is negative (q < 0).
The surroundings warm up; the reaction vessel feels hot.
Example: combustion of methane.
Endothermic reactions absorb energy from the surroundings.
Energy flows into the system, so q is positive (q > 0).
The surroundings cool down; the reaction vessel feels cold to the touch.
Example: N2(g) + O2(g) + energy (heat) -> 2 NO(g).
Two key definitions before moving on:
Thermodynamics = the study of energy and its interconversions.
First Law of Thermodynamics = the total energy of the universe is constant.
Internal energy (E) = the sum of all kinetic and potential energies of every particle in the system.
You cannot measure E itself, but you can measure the change: Delta E.
Internal energy can be changed by a flow of work (w), heat (q), or both.
The central equation:
Delta E = q + w
All signs are from the system's point of view:
(a) Endothermic reaction: heat flows into the system, so q is positive (q > 0).
(b) Exothermic reaction: heat flows out of the system, so q is negative (q < 0).
(c) System does work on surroundings: energy flows out, so w is negative (w < 0). This happens when a gas expands.
(d) Surroundings do work on system: energy flows in, so w is positive (w > 0). This happens when a gas is compressed.
Worked example: Calculate Delta E for a system undergoing an endothermic process in which 15.6 kJ of heat flows and where 1400 J of work is done on the system.
q = +15.6 kJ (endothermic, heat flows in)
w = +1400 J = +1.400 kJ (work done on the system, energy flows in)
Delta E = q + w = +15.6 kJ + 1.400 kJ = +17.0 kJ
The system is gaining 17.0 kJ of energy.
In chemical processes, work is usually done by gases expanding against an external pressure or by gases being compressed.
The relationship: w = -P times Delta V, where P is external pressure (usually constant) and Delta V = V(final) - V(initial).
Case 1: Gas expands
Volume increases, so Delta V is positive.
P is always positive, so w = -P(Delta V) is negative.
Energy flows out of the system (the system is pushing the surroundings back).
Consistent with sign convention (c) above: w < 0.
Case 2: Gas is compressed
Volume decreases, so Delta V is negative.
w = -P(Delta V) is positive (negative times negative).
Energy flows into the system (the surroundings are squeezing the gas).
Consistent with sign convention (d) above: w > 0.
Worked example: Calculate the work for a gas expanding from 46 L to 64 L at a constant external pressure of 15 atm.
w = -P(Delta V) = -(15 atm)(64 L - 46 L) = -(15 atm)(+18 L) = -270 L atm
The negative sign means energy is flowing out of the system as work.
Tip: do not think of the minus sign as "negative number." Think of it as a direction indicator.
Worked example: A balloon is inflated by heating the air inside it. The volume increases from 4.00 x 10^6 L to 4.50 x 10^6 L by the addition of 1.3 x 10^8 J of energy as heat. The balloon expands against a constant pressure of 1.0 atm. Calculate Delta E. (Note: 1 L atm = 101.3 J)
q = +1.3 x 10^8 J (heat added to the system)
Delta V = 4.50 x 10^6 - 4.00 x 10^6 = 5.0 x 10^5 L
w = -(1.0 atm)(5.0 x 10^5 L) = -5.0 x 10^5 L atm
Convert: -5.0 x 10^5 L atm x (101.3 J / 1 L atm) = -5.1 x 10^7 J
Delta E = q + w = (+1.3 x 10^8 J) + (-5.1 x 10^7 J) = 8.0 x 10^7 J
Formula | Meaning | Units |
|---|---|---|
KE = 1/2 mv^2 | Kinetic energy | J (kg m^2/s^2) |
Delta E = q + w | Change in internal energy | J or kJ |
w = -P Delta V | Work from gas expansion or compression | L atm (convert to J with 1 L atm = 101.3 J) |
Unit conversion to remember: 1 L atm = 101.3 J
The combustion of methane (CH4 + 2O2 -> CO2 + 2H2O + energy) is the reaction running every natural-gas boiler and gas hob. Understanding the energy balance is how engineers size heating systems.
Hot packs and cold packs in first-aid kits rely on exothermic and endothermic dissolution reactions, respectively. The chemistry in this chapter explains exactly why one feels warm and the other feels cold.
"Heat and temperature are the same thing." They are not. Temperature measures average particle motion; heat is the transfer of energy caused by a temperature difference. A large lake and a small cup of tea can be at the same temperature, but the lake stores vastly more thermal energy.
"Exothermic means the system is hot." The system is actually losing energy. It is the surroundings that warm up. The system's energy goes down.
"A negative sign on w means you did something wrong." Not at all. The negative sign in w = -P Delta V is a direction indicator. When a gas expands, work is negative because energy leaves the system.
"Internal energy E can be measured directly." You can only measure the change, Delta E. The absolute internal energy of a system is not experimentally accessible.
Exams frequently test whether you can assign the correct sign to q and w for a described process. Memorise the sign conventions from the system's point of view.
A classic exam question gives you q and w (sometimes in different units) and asks you to calculate Delta E. Watch your unit conversions, especially between J and kJ, and between L atm and J.
Expect a question that describes a gas expanding or being compressed and asks you to calculate w using w = -P Delta V. Make sure you compute Delta V as V(final) minus V(initial), not the other way round.
Know the difference between temperature and heat. Short-answer and multiple-choice questions routinely test this distinction.
True or false: Heat flows spontaneously from a cold object to a hot object.
False. Heat flows from hot to cold.
Fill in the blank: In an exothermic reaction, q is ______ (positive/negative).
Negative.
True or false: If a gas is compressed, w is positive.
True. The surroundings are doing work on the system, so energy flows in.
Fill in the blank: Delta E = ______ + ______.
q + w.
True or false: Temperature and heat are the same quantity measured in different units.
False. Temperature measures average particle motion; heat is the transfer of energy due to a temperature difference.
Q: A system absorbs 22.0 kJ of heat and does 8.5 kJ of work on the surroundings. What is Delta E?
A: q = +22.0 kJ (absorbed). w = -8.5 kJ (system does work, energy leaves). Delta E = 22.0 + (-8.5) = +13.5 kJ.
Q: A gas expands from 12.0 L to 25.0 L against a constant external pressure of 3.0 atm. Calculate w in L atm and in joules.
A: Delta V = 25.0 - 12.0 = 13.0 L. w = -P Delta V = -(3.0 atm)(13.0 L) = -39.0 L atm. In joules: -39.0 x 101.3 = -3950.7 J, or approximately -3.95 kJ.
Q: Is the following reaction exothermic or endothermic? N2(g) + O2(g) + energy -> 2 NO(g)
A: Endothermic. Energy appears on the reactant side, meaning the system must absorb energy for the reaction to proceed.
Q: A reaction releases 50 kJ of heat at constant volume, and no work is done. What is Delta E?
A: q = -50 kJ (released). w = 0 (constant volume, no expansion). Delta E = -50 + 0 = -50 kJ.
Q: Explain why an endothermic reaction feels cold to the touch.
A: In an endothermic reaction, heat flows from the surroundings (including your hand) into the system. Your hand loses thermal energy, so it registers a drop in temperature.
This material leads directly into calorimetry (measuring heat experimentally) and enthalpy, which are covered in Part 2 of these notes.
The First Law of Thermodynamics reappears in discussions of entropy and the Second Law later in the course. The sign conventions for q and w established here carry through.
The concept of potential energy stored in chemical bonds connects back to the chapter on chemical bonding and will return when you study bond energies.
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