Thermochemistry

This is Unit Four of Chemistry, Part One: heat, calorimetry and enthalpy. The through-line is that four words this unit depends on — energy, heat, temperature, and enthalpy — all mean something narrower in chemistry than in ordinary speech, and most errors in thermochemistry come from carrying the loose meaning into a calculation. Heat is not something a body contains. Temperature is not how much thermal energy is present. And the heat you measure in a sealed vessel is not the enthalpy change. By the end you will be able to say what temperature measures and what it does not; tell heat capacity from specific heat capacity; say what calorimetry actually observes; say why a bomb calorimeter's heat is not an enthalpy change; and say which sign of delta H marks an exothermic reaction.

Three entries. What energy actually is as a definition, the two forms it takes, and the law that governs every conversion between them.

A capacity.

Chem 0621 Energy

Energy is the capacity of a system to supply heat or to do work on its surroundings. Inflating a bicycle tyre by hand does work, since the pump moves air against the pressure already inside. Energy names a capacity to act rather than a substance stored in one fixed location inside an object.

Position, or motion.

Chem 0624 Potential and Kinetic Energy

Potential energy comes from an object's position, composition or condition; kinetic energy comes from its motion. Water held behind a dam has potential energy that converts to kinetic energy as it falls and spins the turbines below. A substance's stored chemical energy counts as potential energy even while the substance sits perfectly still on a shelf.

Never created.

Chem 0627 Law of Conservation of Energy

During a chemical or physical change, energy changes from one form into another but is never created or destroyed overall. Gasoline's stored chemical potential energy converts into the kinetic energy of a car's moving pistons, not into additional energy. A reaction that seems to produce energy from nothing is converting some unmeasured input form; it never actually breaks this law.

Five entries, and this is where the everyday meanings do the most damage. What temperature actually measures, what heat actually is, where the observer draws the line between system and surroundings, which direction heat crosses that line, and the two units involved.

Per particle.

Chem 0630 Thermal Energy and Temperature

Thermal energy is the kinetic energy of randomly moving atoms and molecules; temperature measures how hot that motion makes a sample. Molecules in hot water move faster on average than in cold water, so the hot sample reads higher. Temperature measures motion per particle, never the total thermal energy in a sample, regardless of that sample's size.

A transfer.

Chem 0633 Heat

Heat is the transfer of thermal energy that flows between two bodies at different temperatures. A warm hand cupped around a snowball transfers heat into the ice until the snowball melts and the hand cools. Heat names a transfer already in progress, never a fixed quantity a body contains once it has reached thermal equilibrium.

Drawing the line.

Chem 0636 System and Surroundings

The system is the matter undergoing a studied change; the surroundings are everything else able to exchange heat with it. In a calorimetry experiment the reacting substances are the system, while the surrounding solution and the calorimeter body count as surroundings. That boundary is a choice the observer draws for the analysis, not a physical wall built into the apparatus.

Which way it crosses.

Chem 0637 Exothermic and Endothermic Processes

An exothermic process releases heat from the system into its surroundings; an endothermic process absorbs heat into the system. A cold pack feels cold because dissolving salt inside it absorbs heat from your hand — endothermic for the pack. The common error is labelling a process from its own temperature change rather than from the direction heat crosses the boundary.

Two units.

Chem 0640 The Joule and the Calorie

The joule is the SI unit of energy, and the calorie is a defined unit set equal to exactly 4.184 joules. A nutritional Calorie, written with a capital C, equals a thousand ordinary calories — one kilocalorie. The calorie's fixed ratio to the joule is set by definition rather than being a measured value some future experiment could revise.

Four entries. Two properties that sound alike and differ in one crucial way, then the equation that connects them to a measured temperature change, worked in both directions.

Depends on how much.

Chem 0643 Heat Capacity

Heat capacity is the heat a body absorbs or releases per degree of change, and it scales with the matter present. A swimming pool holds a far greater heat capacity than a cup of water, since it contains vastly more matter to heat. Heat capacity is extensive — comparing capacities compares amounts of substance, never what each is made of.

Depends on what.

Chem 0645 Specific Heat Capacity

Specific heat capacity is the heat needed to raise one gram of a substance by one degree, and it depends only on the substance. A small and a large iron pan yield the same specific heat for iron, though their heat capacities differ fivefold. Specific heat capacity is intensive — fixed for a substance regardless of sample size.

Three factors.

Chem 0648 Calculating Heat from Mass, Specific Heat, and Temperature Change

Calculating heat from a temperature change means multiplying a substance's mass by its specific heat and by the temperature change it undergoes. Eight hundred grams of water heated from twenty-one to eighty-five degrees, times water's specific heat of 4.184 joules per gram-degree, times the sixty-four degree change, gives about 2.1 times ten to the fifth joules absorbed. Using the initial temperature instead of the change produces an answer unconnected to how much the sample warmed.

The same equation, rearranged.

Chem 0651 Solving for an Unknown Specific Heat

Solving for an unknown substance's specific heat means dividing the total measured heat by the product of the sample's mass and its temperature change — the same equation rearranged. A 348-gram metal absorbing 6.64 kilojoules over a 21.2 degree rise gives a specific heat near 0.90, identifying aluminium against a reference table. Forgetting to convert kilojoules to joules first shifts the result by a factor of a thousand, pointing to the wrong substance entirely.

Four entries. What the technique actually does — which is less direct than it sounds — the simplest instrument for it, and two calculations that both rest on heat lost equalling heat gained.

Inferred, not seen.

Chem 0654 Calorimetry

Calorimetry measures the heat transferred during a process by observing the temperature change of a calibrated calorimeter. A reaction run inside an insulated calorimeter changes its temperature by an amount revealing how much heat it released or absorbed. Calorimetry measures a temperature change and infers the system's heat from it — it never observes the heat flow directly.

A polystyrene cup.

Chem 0657 The Coffee-Cup Calorimeter

A coffee-cup calorimeter is a simple, constant-pressure device measuring heat exchange in a solution from its temperature change. A polystyrene cup holding a reacting solution lets a thermometer track the rise or fall as the reaction proceeds. Its imperfect insulation lets some heat escape to the room, giving less accurate values than a sealed instrument.

Lost equals gained.

Chem 0660 Calculating Heat Transfer Between Two Substances

Calculating heat transfer between two substances at different temperatures means setting the heat lost by the hotter equal in magnitude to the heat gained by the cooler, since ideally none escapes. A 360-gram steel rod dropped into 425 millilitres of water at twenty-four degrees settles at 42.7, and solving back gives a rod starting near 248 degrees. Treating their final temperatures as different values breaks the equation the whole calculation rests on.

Equal and opposite.

Chem 0663 Calculating Heat from Solution Calorimetry

Calculating a reaction's heat from solution calorimetry means treating that heat as equal in magnitude and opposite in sign to the heat gained by the surrounding solution. Mixing fifty millilitres each of one-molar acid and base at twenty-two degrees raises the combined solution to 28.9, a gain of about 2.9 kilojoules — so the reaction's own heat is minus 2.9. Reporting the solution's positive gain without flipping the sign mislabels an exothermic reaction as endothermic.

Two entries. A sealed instrument that holds volume constant instead of pressure, and the calculation it requires — which has one term the coffee cup did not.

Sealed steel.

Chem 0666 The Bomb Calorimeter

A bomb calorimeter measures heat at constant volume, inside a sealed steel vessel built for combustion reactions. A weighed fuel sample is ignited by a spark inside the bomb, and the surrounding water warms by an amount revealing the heat released. The steel body absorbs a share of that heat too, so its known heat capacity enters the calculation.

Two contributions.

Chem 0669 Calculating Heat from Bomb Calorimetry

Calculating heat from bomb calorimetry means adding the heat absorbed by the surrounding water to the heat absorbed by the bomb itself, since both make up the calorimeter's full surroundings. Burning 3.12 grams of glucose in a bomb holding 775 grams of water, with the steel's own capacity at 893 joules per degree, gives a total release near 48.7 kilojoules. Computing only the water's gain understates the heat the reaction actually released.

Four entries to close the part. The law relating heat and work to a system's internal energy, the property that makes a quantity path-independent, the state function chemists actually use — and the sign convention that decides whether your answer means the reaction gave heat out or took it in.

Heat plus work.

Chem 0672 Internal Energy and the First Law of Thermodynamics

A system's internal energy changes by the heat it absorbs or releases plus the work done on or by it. A metal wire warms both when dropped in hot water, which adds heat, and when bent back and forth, which adds work. Heat and work are two separate routes to the same internal-energy change, not rival or mutually exclusive explanations.

Path-independent.

Chem 0675 State Function

A state function is a property whose value depends only on a system's current state, not on the path taken to reach it. A hiker's elevation on a mountain is identical whether the trail was direct or roundabout, unlike the distance walked. Heat and work each depend on the path taken, so neither alone is a state function — though their combined internal-energy change is.

Heat at constant pressure.

Chem 0678 Enthalpy

Enthalpy is a state function equal to internal energy plus pressure times volume, and its change equals heat flow at constant pressure. The heat released by a Bunsen burner's flame equals its combustion reaction's enthalpy change, since it burns at roughly constant pressure. A reaction's heat measured in a sealed bomb calorimeter is not its enthalpy change, since pressure there can shift freely.

Which sign means which.

Chem 0681 Sign Convention for Enthalpy Change

A negative enthalpy change marks an exothermic reaction and a positive one marks endothermic, and reversing a reaction flips the sign. Hydrogen and oxygen forming liquid water release 286 kilojoules per mole, so delta H is minus 286; decomposing water back carries plus 286. Reasoning that a reaction gets cold so it must release heat inverts the convention — a falling temperature marks endothermic, a positive delta H.

Five things to carry out of this part. Temperature measures the kinetic energy of motion per particle, and what it does not measure is the total thermal energy in a sample — which is why a bathtub of warm water holds far more thermal energy than a spark that is very much hotter. Heat capacity is extensive: it scales with how much matter is present, so a pool has a larger one than a cup. Specific heat capacity is intensive: it is fixed for a substance regardless of sample size, which is why a small and a large iron pan share the same value. Calorimetry never observes a reaction's heat directly — it observes a temperature change in a calibrated calorimeter and infers the heat from it, which is why every calorimetry calculation is really an inference from the surroundings back to the system. A bomb calorimeter's heat is not an enthalpy change because a bomb holds volume constant while pressure shifts freely, and enthalpy change equals heat flow only at constant pressure — which is what the coffee-cup calorimeter provides. And a negative delta H marks an exothermic reaction: heat leaves the system, so the system's enthalpy falls. A reaction that turns cold is endothermic and carries a positive delta H. Part Two takes up thermochemical equations and Hess's law.