Thermochemistry Fundamentals
Enthalpy and Heat Changes · Standard Enthalpy Changes and Hess’s Law Preparation
IB Chemistry topic guide
Reactivity 1: What Drives Reactions?
Reactivity 1: What Drives Reactions? is a core part of IB Chemistry. This guide connects the syllabus ideas behind Thermochemistry Fundamentals, Calorimetry, Energy Cycles, Bonding and Enthalpy, Entropy and Spontaneity, shows how they appear in worked problems, and points you to the formulas and full lessons needed for exam revision.
What you will learn
Reactivity 1: What Drives Reactions? syllabus outline
The units below follow the structure used in the full Study to Learn course. Use the outline to identify exactly which idea needs attention, then work through the public example before continuing to the complete lesson path.
Calorimetry
The Calorimetry Calculation · Sources of Error and Improving Accuracy
Energy Cycles
Hess’s Law · Born–Haber Cycles
Bonding and Enthalpy
Bond Enthalpies · Bond Length, Bond Order, and Why Averages Vary
Entropy and Spontaneity
Entropy and Gibbs Free Energy · Predicting Entropy Changes and Finding the Switching Temperature
Free worked preview
Enthalpy and Heat Changes
This complete preview comes from the Thermochemistry Fundamentals unit. It introduces the core language, shows the method in context, and gives you a real example of the lesson quality before you create an account.
The Energy of Chemical Systems
Every chemical reaction involves energy change — some release heat (exothermic), others absorb it (endothermic). This lesson introduces enthalpy, the key thermodynamic quantity for measuring these energy changes, and the sign conventions and standard conditions you will use throughout the energetics section. After this lesson, you will understand the difference between exothermic and endothermic reactions, what standard enthalpy changes of combustion and formation mean, and how to interpret enthalpy level diagrams correctly.
Enthalpy (H) is a measure of the total energy of a system. We cannot measure H directly, but we can measure the change in enthalpy (ΔH) during a reaction.
Key Definitions
- Exothermic: ΔH < 0 (negative) — energy is released to the surroundings; the surroundings’ temperature increases.
- Endothermic: ΔH > 0 (positive) — energy is absorbed from the surroundings; the surroundings’ temperature decreases.
- Standard enthalpy change (ΔH⁰): measured at 100 kPa and a specified temperature (usually 298 K).
- Standard enthalpy of combustion (ΔHᴄ⁰): energy released when 1 mol of a substance is completely burned in excess oxygen under standard conditions.
- Standard enthalpy of formation (ΔHₐ⁰): energy change when 1 mol of a compound is formed from its constituent elements in their standard states.
Enthalpy Level Diagrams
Exothermic reaction: products are drawn lower (at lower enthalpy) than reactants; ΔH arrow points downward (negative). Endothermic reaction: products are drawn higher than reactants; ΔH arrow points upward (positive). The activation energy (Eₐ) is always drawn as an energy barrier (hump) between reactants and products, regardless of whether the reaction is endothermic or exothermic.
Common ErrorThe sign convention is always from the system’s perspective. ΔH negative = the system lost energy (exothermic). ΔH positive = the system gained energy (endothermic). Don’t confuse this with what the thermometer shows (which is the surroundings, not the system).
Reviewed by the Study to Learn editorial team · Updated 2026-08-04