Thermochemistry tracks the energy that flows in and out of chemical reactions and physical changes as heat and work. By the end of this unit you should be able to tell exothermic from endothermic processes by sign, calculate ΔH using Hess's law or formation enthalpies, and turn calorimeter temperature data into a heat of reaction.
Read the overview first, then work through each topic note and try the checkpoint question before moving on.
This unit connects the particle picture of matter to energy you can actually measure. You'll define the system and surroundings, follow heat and work across that boundary, and use enthalpy to describe how much energy a reaction releases or absorbs.
Honors Chemistry Unit 8 covers thermochemistry, the study of energy changes during chemical reactions and physical changes. It starts with the forms of energy and the difference between heat, work, and temperature. Next comes the first law of thermodynamics (ΔU = q + w) and enthalpy change (ΔH), including how to calculate ΔH with Hess's law, standard enthalpies of formation, and bond enthalpies. It ends with calorimetry, where you use q = mcΔT and temperature data from a coffee cup or bomb calorimeter to find the heat of a reaction.
The system is the reaction you're studying, and the surroundings are everything else, including the water and the beaker. A positive q means heat flows into the system. A negative q means the system releases heat and the surroundings warm up.
Most lab reactions open to the air occur at approximately constant pressure. When pressure-volume work is the only work involved, ΔH equals the heat exchanged at constant pressure. Exothermic reactions have ΔH < 0, and endothermic reactions have ΔH > 0. Because enthalpy is a state function, you can find ΔH by adding reaction steps with Hess's law or by using tabulated formation values.
You measure the temperature change of the surroundings rather than the reaction's heat itself. Use q = mcΔT to calculate the heat gained or lost by the solution or water. If the calorimeter's heat capacity and heat loss are negligible, the reaction's heat has the same magnitude with the opposite sign.
The first law says the energy lost by the system equals the energy gained by the surroundings. That single idea explains the sign conventions, lets Hess's law work, and makes calorimetry possible. If you can track where the energy goes, you can solve almost every problem in this unit.
Energy is the ability to do work or transfer heat, and it's measured in joules (J).
The first law of thermodynamics says energy can't be created or destroyed.
Calorimetry determines a reaction's heat from the temperature change of its surroundings, such as a solution, water bath, and calorimeter.
Energy is the ability to do work or transfer heat, and it's measured in joules (J). Kinetic energy is energy of motion, and potential energy comes from position or arrangement. Chemical energy is potential energy associated with the arrangement and electrostatic interactions of atoms and electrons. Breaking bonds requires energy, while forming bonds releases energy.
Heat and temperature are not the same thing. Temperature measures the average kinetic energy of the particles in a sample. Heat (q) is energy that moves from a hotter object to a cooler one because of a temperature difference. A full bathtub at 40°C holds far more thermal energy than a cup of tea at 90°C, even though the tea has the higher temperature.
Work (w) is energy transferred when a force moves something. In chemistry this is usually pressure-volume work. At constant external pressure, . When a gas expands against the atmosphere, is positive, so w is negative and the system transfers energy to the surroundings as work.
In an exothermic process the system releases heat, so the surroundings get warmer. In an endothermic process the system absorbs heat, so the surroundings get cooler. An instant cold pack feels cold because dissolving ammonium nitrate pulls heat out of your hand.
For units, 1 cal = 4.184 J, and the Calorie on a food label is a kilocalorie, or 4184 J.
Check yourself. You dissolve a salt in water and the beaker gets colder. Is the dissolving process exothermic or endothermic, and what is the sign of q for the system? (Answer: The process is endothermic. The system absorbed heat from the water, so q is positive.)
| Feature | Exothermic | Endothermic |
|---|---|---|
| Direction of heat | System to surroundings | Surroundings to system |
| Sign of q and ΔH | Negative | Positive |
| Surroundings temperature | Rises | Falls |
| Example | Burning methane, ΔH = -890 kJ/mol | Dissolving NH₄NO₃ in a cold pack |
The first law of thermodynamics says energy can't be created or destroyed. For a system, , where is the change in internal energy. If a system absorbs 250 J of heat and does 100 J of work on its surroundings, J.
Enthalpy (H) is defined as . At constant pressure, when pressure-volume work is the only work involved, the enthalpy change equals the heat exchanged, so . ΔH depends on the physical states involved and scales with the coefficients. Doubling a reaction doubles ΔH, and reversing it flips the sign.
You can calculate ΔH three ways. Hess's law adds the ΔH values of steps that sum to the overall reaction. Standard enthalpies of formation use , where any element in its standard state has . Bond enthalpies give an estimate, .
Worked example. Find ΔH° for . Use values of -74.8 kJ/mol for CH₄, -393.5 for CO₂, -285.8 for H₂O(l), and 0 for O₂. Products give kJ. Reactants give kJ. So kJ, which is strongly exothermic.
Check yourself. The reaction A → 2B has ΔH = +40 kJ. What is ΔH for 4B → 2A? (Answer: -80 kJ. Reversing the reaction changes the sign to -40 kJ, and doubling the coefficients doubles ΔH.)
Calorimetry determines a reaction's heat from the temperature change of its surroundings, such as a solution, water bath, and calorimeter. Because energy is conserved, . You may use when the water is treated as the only part of the surroundings that absorbs or releases heat.
The heat gained or lost by a substance is , where m is mass in grams, c is specific heat in J/g·°C, and . Water's specific heat is a high 4.18 J/g·°C, so water absorbs a lot of heat with a small temperature change.
A coffee cup calorimeter usually consists of nested foam cups with a lid and thermometer. It operates at approximately constant pressure, so you use the temperature change to infer the reaction's ΔH, assuming pressure-volume work is the only work involved. Heat leaking to the room and heat absorbed by the cup are common sources of error. A bomb calorimeter keeps volume constant, so its calibrated temperature change gives and therefore ΔU for reactions such as combustion.
Worked example. You mix 50.0 mL of 1.0 M HCl with 50.0 mL of 1.0 M NaOH, and the temperature rises from 22.0°C to 28.7°C. Treat the solution as 100.0 g of water and ignore heat absorbed by the cup. J, so kJ. The reaction formed 0.0500 mol of water, so kJ/mol.
Check yourself. How much heat does 200.0 g of water absorb when it warms from 20.0°C to 25.0°C? (Answer: q = (200.0)(4.18)(5.0) = 4.2 × 10³ J, or 4.2 kJ.)
| Feature | Coffee cup calorimeter | Bomb calorimeter |
|---|---|---|
| Condition held constant | Pressure (approximately) | Volume |
| Quantity determined | ΔH, inferred from q at constant pressure | ΔU, from q at constant volume |
| Typical use | Dissolving, neutralization, aqueous reactions | Combustion of fuels and foods |
| Precision | Lower, some heat escapes | Higher, well insulated and sealed |
| Term | Definition |
|---|---|
| System | The system is the part of the universe you are studying, usually the reacting chemicals. |
| Surroundings | The surroundings are everything outside the system, such as the water, container, and air around a reaction. |
| Exothermic | An exothermic process releases heat to the surroundings and has a negative ΔH. |
| Endothermic | An endothermic process absorbs heat from the surroundings and has a positive ΔH. |
| First law of thermodynamics | Energy is conserved, so a system's energy change equals the heat plus work it exchanges, ΔU = q + w. |
| Enthalpy change (ΔH) | When pressure-volume work is the only work involved, enthalpy change is the heat absorbed or released by a reaction at constant pressure. |
| Hess's law | The total ΔH for a reaction is the same no matter which set of steps you use to get from reactants to products. |
| Bond enthalpy | Bond enthalpy is the energy needed to break one mole of a particular bond in the gas phase, and it can be used to estimate ΔH. |
| Specific heat capacity | Specific heat capacity is the heat required to raise the temperature of 1 gram of a substance by 1°C. |
| Calorimeter | A calorimeter is a device that measures the heat of a reaction by tracking the temperature change of its surroundings. |
When the water warms up because of a reaction, q for the water is positive, so it's tempting to call the reaction positive too. The water gained heat because the reaction lost it, so q_rxn = -q_surroundings. In a constant-pressure calorimeter, a surroundings temperature rise caused by the reaction indicates an exothermic reaction with a negative ΔH.
Temperature is an average particle energy, while heat is energy in transit. A large mass at a lower temperature can hold more thermal energy than a small hot one. Keep them separate by remembering that q depends on mass and specific heat, not just ΔT.
ΔH°f values are per mole, so each one has to be multiplied by its coefficient in the balanced equation. Also remember that elements in their standard states, like O₂(g) or Fe(s), have ΔH°f = 0, not a missing value.
If you reverse a step to match the target reaction, change the sign of its ΔH. If you multiply a step by 2 or 3, multiply its ΔH by the same number. Check your work by confirming the intermediates cancel and the steps add up to exactly the target equation.