Thermodynamics is all about energy changes in chemical reactions. Understanding...
Chemistry199 views·Updated 9 Jul 2026·2 pagesUnderstanding Thermodynamics and Enthalpy: Key Definitions in Physical Chemistry
Aasiyah Rahman@aasiyahrahman
1
of 2
Enthalpy Definitions
Enthalpy change (ΔH) represents the heat energy transferred during a reaction that occurs at constant pressure. When this change is measured under standard conditions (298K and 100kPa), we call it standard enthalpy change (ΔH°). Reactions are classified as either exothermic (negative ΔH, releasing energy) or endothermic (positive ΔH, absorbing energy).
There are several specific types of enthalpy changes you'll need to know. Enthalpy of formation measures the energy change when one mole of a compound forms from its elements in their standard states. Enthalpy of combustion measures energy released when one mole of substance burns completely in oxygen.
The enthalpy of neutralisation describes energy released when an acid and alkali react to form one mole of water. Bond dissociation enthalpy is the energy needed to break all bonds of the same type in one mole of gaseous molecules.
⚡ Think of enthalpy changes like bank transactions: exothermic reactions are like deposits (giving energy to surroundings), while endothermic reactions are withdrawals (taking energy from surroundings).
Other important enthalpy terms include ionisation energy (forming positive ions), electron affinity (forming negative ions), enthalpy of atomisation (forming gaseous atoms), hydration enthalpy (when ions dissolve in water), and solution enthalpy (when a solute completely dissolves).
2
of 2
Types of Enthalpy Changes
Every enthalpy change measures a specific type of chemical or physical process. For instance, formation enthalpy (ΔHf) tells us about making compounds from elements, while combustion enthalpy (ΔHc) measures energy released during burning. These values help chemists predict whether reactions will be feasible and how much energy they'll release.
Most formation reactions are exothermic (negative ΔH), meaning they release energy when compounds form. Combustion reactions are almost always exothermic too, which is why burning fuels generates heat. Neutralisation reactions between acids and bases also release energy, with a consistent enthalpy change when water forms.
Breaking bonds always requires energy, which is why bond dissociation and lattice dissociation enthalpies are endothermic (positive ΔH). Similarly, changing states from solid to liquid (fusion) or liquid to gas (vaporisation) requires energy input. Knowing these patterns helps you predict reaction outcomes.
🔍 Remember that exothermic reactions feel hot to touch because they release energy to surroundings, while endothermic reactions feel cold because they absorb energy from surroundings.
Ionisation enthalpies are always endothermic because removing electrons requires energy. The first electron affinity is usually exothermic for non-metals (they want electrons), but the second electron affinity is endothermic (they resist taking another electron due to repulsion).
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Chemistry199 views·Updated 9 Jul 2026·2 pagesUnderstanding Thermodynamics and Enthalpy: Key Definitions in Physical Chemistry
Aasiyah Rahman@aasiyahrahman
Thermodynamics is all about energy changes in chemical reactions. Understanding enthalpy changes helps us predict whether reactions release or absorb heat energy, and by how much. This topic is crucial for explaining why some reactions feel hot while others cool...
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of 2
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Enthalpy Definitions
Enthalpy change (ΔH) represents the heat energy transferred during a reaction that occurs at constant pressure. When this change is measured under standard conditions (298K and 100kPa), we call it standard enthalpy change (ΔH°). Reactions are classified as either exothermic (negative ΔH, releasing energy) or endothermic (positive ΔH, absorbing energy).
There are several specific types of enthalpy changes you'll need to know. Enthalpy of formation measures the energy change when one mole of a compound forms from its elements in their standard states. Enthalpy of combustion measures energy released when one mole of substance burns completely in oxygen.
The enthalpy of neutralisation describes energy released when an acid and alkali react to form one mole of water. Bond dissociation enthalpy is the energy needed to break all bonds of the same type in one mole of gaseous molecules.
⚡ Think of enthalpy changes like bank transactions: exothermic reactions are like deposits (giving energy to surroundings), while endothermic reactions are withdrawals (taking energy from surroundings).
Other important enthalpy terms include ionisation energy (forming positive ions), electron affinity (forming negative ions), enthalpy of atomisation (forming gaseous atoms), hydration enthalpy (when ions dissolve in water), and solution enthalpy (when a solute completely dissolves).
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Types of Enthalpy Changes
Every enthalpy change measures a specific type of chemical or physical process. For instance, formation enthalpy (ΔHf) tells us about making compounds from elements, while combustion enthalpy (ΔHc) measures energy released during burning. These values help chemists predict whether reactions will be feasible and how much energy they'll release.
Most formation reactions are exothermic (negative ΔH), meaning they release energy when compounds form. Combustion reactions are almost always exothermic too, which is why burning fuels generates heat. Neutralisation reactions between acids and bases also release energy, with a consistent enthalpy change when water forms.
Breaking bonds always requires energy, which is why bond dissociation and lattice dissociation enthalpies are endothermic (positive ΔH). Similarly, changing states from solid to liquid (fusion) or liquid to gas (vaporisation) requires energy input. Knowing these patterns helps you predict reaction outcomes.
🔍 Remember that exothermic reactions feel hot to touch because they release energy to surroundings, while endothermic reactions feel cold because they absorb energy from surroundings.
Ionisation enthalpies are always endothermic because removing electrons requires energy. The first electron affinity is usually exothermic for non-metals (they want electrons), but the second electron affinity is endothermic (they resist taking another electron due to repulsion).
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You can download the app from Google Play Store and Apple App Store.
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