Enthalpy change is the amount of heat energy absorbed or released during a chemical reaction or physical process at constant pressure. Written as ΔH (delta H), it tells you whether a process gives off heat to its surroundings or pulls heat in. If you’re studying chemistry, it’s one of the most practical ways to track where energy goes when substances react.
How Enthalpy Change Works
Enthalpy (H) is defined as a system’s internal energy plus the product of its pressure and volume. You don’t need to memorize that formula to understand the key idea: when a reaction happens at constant pressure (which covers most reactions in open lab beakers or in your body), the change in enthalpy equals the heat transferred. That makes ΔH a direct measure of heat flow.
The sign of ΔH tells you the direction of that flow. A negative ΔH means the reaction releases heat into the surroundings. These are called exothermic reactions, and they feel warm or hot. Burning fuel, mixing cement with water, and neutralizing an acid with a base are all exothermic. A positive ΔH means the reaction absorbs heat from the surroundings, making them feel cooler. These are endothermic reactions. Dissolving certain salts in water or photosynthesis are common examples.
So when you see ΔH = −400 kJ, that means 400 kilojoules of energy left the reacting chemicals and entered the surroundings as heat. If ΔH = +200 kJ, the reaction pulled 200 kilojoules from its surroundings.
Where the Energy Comes From
At the molecular level, enthalpy change comes down to chemical bonds. Breaking bonds requires energy. Forming bonds releases energy. The overall ΔH for a reaction is the total energy needed to break the bonds in the reactants minus the total energy released when new bonds form in the products.
If the new bonds release more energy than the old bonds took to break, the reaction is exothermic (negative ΔH). If breaking bonds costs more energy than forming new ones gives back, the reaction is endothermic (positive ΔH). This is why you can estimate the enthalpy change of a reaction using tables of bond energies: add up the bond energies on each side and find the difference.
Common Types of Enthalpy Change
Chemists categorize enthalpy changes by the type of process involved. The most common ones you’ll encounter are:
- Enthalpy of formation: the energy change when one mole of a compound forms from its pure elements in their standard states. This serves as a reference point for comparing different substances.
- Enthalpy of combustion: the energy released when one mole of a substance burns completely in oxygen. This is how we quantify the energy content of fuels.
- Enthalpy of neutralization: the energy released when an acid reacts with a base to produce one mole of water. For strong acids reacting with strong bases, this value falls consistently between −57 and −58 kJ per mole of water formed, because the underlying reaction is always the same: hydrogen ions combining with hydroxide ions to make water. Weak acids produce smaller values because some energy is used up separating the acid molecules before they can react. Hydrocyanic acid neutralized by potassium hydroxide, for instance, releases only about −11.7 kJ per mole.
Standard Conditions
When you see the symbol ΔH° (with the degree symbol), it refers to enthalpy change measured under standard conditions. IUPAC defines these as a standard pressure applied to pure substances in a specified state of matter. Standard conditions exist so that scientists everywhere can compare enthalpy values on a level playing field. Without them, the same reaction could produce different reported values depending on the lab’s altitude or the concentration of solutions used.
Hess’s Law and Why the Path Doesn’t Matter
Enthalpy is a state function. That means the enthalpy change between a starting point and an ending point is always the same, regardless of what happens in between. If you convert reactant A into product D through intermediates B and C, the total ΔH is identical to what you’d get converting A directly into D in a single step.
This principle is called Hess’s Law, and it’s enormously useful. Many reactions are difficult or dangerous to carry out directly in a lab. Hess’s Law lets you calculate their enthalpy changes by adding up the ΔH values of simpler, measurable reactions that reach the same endpoint. If you can find a sequence of known reactions that, when combined, give you the reaction you’re interested in, you simply add their enthalpy changes together to get the answer.
Measuring Enthalpy Change With Calorimetry
In practice, enthalpy changes are measured using a technique called calorimetry. The basic idea is straightforward: carry out a reaction inside an insulated container of water and measure how much the water’s temperature changes. The heat absorbed or released by the water equals the heat released or absorbed by the reaction.
The core formula is q = mcΔT, where q is the heat transferred, m is the mass of water, c is the specific heat capacity of water, and ΔT is the temperature change. Water’s specific heat capacity is 4.18 joules per gram per degree Celsius, meaning it takes 4.18 joules of energy to raise one gram of water by one degree. Once you calculate q, you divide by the number of moles of reactant to get the molar enthalpy change in kilojoules per mole.
For example, if a reaction raises the temperature of 100 grams of water by 15 degrees Celsius, the water absorbed 6,270 joules (100 × 4.18 × 15). That heat came from the reaction, so the reaction’s enthalpy change for that quantity of reactant is −6,270 joules (negative because heat left the system). Divide by the moles of reactant used and you have the molar ΔH.
Reading an Energy Diagram
Enthalpy changes are often shown on energy diagrams where the vertical axis represents enthalpy and the horizontal axis tracks the reaction’s progress from reactants to products. In an exothermic reaction, the products sit lower on the diagram than the reactants, and the gap between the two levels equals the enthalpy change. The diagram slopes “downhill.” In an endothermic reaction, the products sit higher and the diagram goes “uphill.”
These diagrams also show a hump between reactants and products representing the activation energy, the initial energy push needed to get the reaction started. A reaction can have a large, favorable (negative) ΔH and still need a significant activation energy to begin. This is why a match needs a strike to light even though burning releases a lot of heat. The enthalpy change tells you the net energy difference between start and finish, while the activation energy tells you how hard it is to get the reaction going.