Revision Notes - Blog Posts
Enthalpy - a thermodynamic property
When I first learned about enthalpy, I was shocked - it felt more like a physics lesson than a chemistry lesson. The thought of learning more about thermodynamics than my basic understanding from my many science lessons in lower school made me bored out of my mind. But enthalpy is actually pretty interesting, once you get your head around it…
Reactions which release heat to their surroundings are described to be exothermic. These are reactions like combustion reactions, oxidation reactions and neutralisation reactions. Endothermic reactions take in heat from their surroundings, such as in thermal decomposition. Reversible reactions are endothermic in one direction and exothermic in the other.
These facts are important when you start to look at enthalpy. Enthalpy is basically a thermodynamic property linked to internal energy, represented by a capital H. This is pretty much the energy released in bond breaking and made in bond making. We usually measure a change in enthalpy, represented by ∆H. ∆H = enthalpy of the products (H1) - enthalpy of the reactants (H2). This is because we cannot measure enthalpy directly.
In exothermic reactions, ∆H is negative whereas in endothermic reactions, ∆H is positive.
∆H is always measured under standard conditions of 298K and 100kPa.
In reversible reactions, the ∆H value is the same numerical value forwards and backwards but the sign is reversed. For example, in a forward exothermic reaction, the ∆H value would be -ve but in the backwards reaction (endothermic) the ∆H would be +ve.
Reaction profiles are diagrams of enthalpy levels of reactants and products in a chemical reaction. X axis is enthalpy rather than ∆H and the Y axis is the progress of reaction, reaction coordinate or extent of reaction. Two horizontal lines show the enthalpy of reactants and products with the reactants on the left and the products on the right. These should be labelled with their names or formulae.
In an endothermic reaction, product lines are higher enthalpy values than reactants. In an exothermic reaction, product lines are lower enthalpy values than reactants. The difference between product and reactant lines is labelled as ∆H. Values are measured in kJ mol-1.
Reaction pathways are shown with lines from the reactants to the products on enthalpy level diagrams. This shows the “journey” that the enthalpy takes during a reaction. They require an input of energy to break bonds before new bonds can form the products. The activation energy is the peak of the pathway above the enthalpy of reactants. It is the minimum amount of energy that reactants must have to react.
Standard enthalpy values are the ∆H values for enthalpy changes of specific reactions measured under standard conditions, represented by ⊖. There are three of these:
1. Standard enthalpy of reaction ( ΔHr⊖ )
The enthalpy change when substances react under standard conditions in quantities given by the equation for the reaction.
2. Standard enthalpy of formation ( ΔfH⊖ )
The enthalpy change when 1 mole of a compound is formed from its constitutent elements with all reactants and products in standard states under standard conditions.
The enthalpy of formation for an element is zero is it is in it’s standard state for example, O2 enthalpy is zero.
3. Standard enthalpy of combustion ( ΔcH⊖ )
The enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in their standard states under standard conditions.
Values for standard enthalpy of formation and combustion must be kept to per mole of what they refer.
Summary
Reactions which release heat to their surroundings are described to be exothermic. Endothermic reactions take in heat from their surroundings, such as in thermal decomposition.
Reversible reactions are endothermic in one direction and exothermic in the other.
Enthalpy is a thermodynamic property linked to internal energy, represented by a capital H. We usually measure a change in enthalpy, represented by ∆H.
∆H = enthalpy of the products (H1) - enthalpy of the reactants (H2). We cannot measure enthalpy directly.
In exothermic reactions, ∆H is negative whereas in endothermic reactions, ∆H is positive.
∆H is always measured under standard conditions of 298K and 100kPa.
In reversible reactions, the ∆H value is the same numerical value forwards and backwards but the sign is reversed.
Reaction profiles are diagrams of enthalpy levels of reactants and products in a chemical reaction. They
In an endothermic reaction, product lines are higher enthalpy values than reactants. In an exothermic reaction, product lines are lower enthalpy values than reactants.
The difference between product and reactant lines is labelled as ∆H.
Values are measured in kJ mol-1.
Reaction pathways are shown with lines from the reactants to the products on enthalpy level diagrams. They plot enthalpy against reaction progress.
Reactions require an input of energy to break bonds before new bonds can form the products. The activation energy is the peak of the pathway above the enthalpy of reactants. It is the minimum amount of energy that reactants must have to react.
Standard enthalpy values are the ∆H values for enthalpy changes of specific reactions measured under standard conditions, represented by ⊖.
Standard enthalpy of reaction ( ΔHr⊖ ) is the enthalpy change when substances react under standard conditions in quantities given by the equation for the reaction.
Standard enthalpy of formation ( ΔfH⊖ ) is the enthalpy change when 1 mole of a compound is formed from its constitutent elements with all reactants and products in standard states under standard conditions.
The enthalpy of formation for an element is zero is it is in it’s standard state.
Standard enthalpy of combustion ( ΔcH⊖ ) is the enthalpy change when 1 mole of a substance is burned completely in excess oxygen with all reactants and products in their standard states under standard conditions.
Values for standard enthalpy of formation and combustion must be kept to per mole of what they refer.
Happy studying!