A-Level H2 Chemistry students in junior colleges (JC) in Singapore always complain about this topic called Chemical Energetics (also known as Thermochemistry) and how difficult the questions are set in their schools. This topic is usually covered in term 1 or term 2 in JC1.
There are 3 sections in Chemical Energetics, namely:
- Enthalpy Changes, ΔH
- Entropy Changes, ΔS
- Gibbs Free Energy, ΔG
After teaching batches of students in the last 16 years, i realised those students who are weak in this topic tends to take short-cuts. They are too “lazy” to understand and remember the key definitions of each of the different Standard Enthalpy Changes of Reactions. No wonder they are struggling with this topic. If you want to master this topic, you cannot take short-cuts and expect yourself to do well.
Students need to have a good grasp of all these fundamentals before they can answer application questions on Hess’ Law, Born Haber Cycle, Energy Level Diagram, etc.
Enthalpy Change of Reaction, ΔH
To start, let’s look at Enthalpy Change of Reaction, ΔH, which is defined as the heat change (heat energy absorbed or evolved) when the reaction takes place between the reagents as indicated by the stoichiometric equation for the reaction.
If the reaction is carried out at standard conditions, it will then be known as Standard Enthalpy Change of Reaction, ΔHθ.
Let’s now take a look at the definitions of the key Standard Enthalpy Changes of Reactions (with examples):
Standard Enthalpy Change of Formation, ΔHfθ
Energy change when 1 mole of the compound is formed from its constituent elements under standard conditions
e.g. H2 (g) + 1/2 O2 (g) → H2O (l) ΔHfθ(H2O)
Standard Enthalpy Change of Combustion, ΔHcθ
Energy released when 1 mole of a substance is completely burned in oxygen under standard conditions
e.g. C2H4 (g) + 3 O2 (g) → 2 CO2 (g) + 2 H2O (l) ΔHcθ(C2H4)
Standard Enthalpy Change of Neutralisation, ΔHneuθ
Energy released when 1 mole of water is formed in the neutralisation between an acid and an alkali under standard conditions
e.g. HCl (aq) + NaOH (aq) → NaCl (aq) + H2O (l) ΔHneuθ
Standard Enthalpy Change of Atomisation, ΔHatomθ
Energy absorbed when an element is converted into 1 mole of free gaseous atoms under standard conditions
e.g. 1/2 O2 (g) → O (g) ΔHatomθ(O)
Standard Enthalpy Change of Hydration, ΔHhydθ
Energy released when 1 mole of the gaseous ion is dissolved in large amount of water under standard conditions
e.g. Na+ (g) + water → Na+ (aq) ΔHhydθ(Na+)
Standard Enthalpy Change of Solution, ΔHsolnθ
Energy change when 1 mole of a substance dissolves in such a large volume of solvent that addition of more solvent produces no further heat change under standard conditions
e.g. NaCl (s) + water → Na+ (aq) + Cl– (aq) ΔHsolnθ(NaCl)
Lattice Energy, ΔHlattθ
Energy released when 1 mole of an ionic crystalline solid is formed from its separate gaseous ions under standard conditions
e.g. Na+ (g) + Cl– (g) → NaCl (s) ΔHlattθ(NaCl)
Besides the definitions above, it is important that you know three more basic definitions also in order to solve all the application type questions, and they are:
Ionisation Energy, I.E.
First Ionisation Energy, 1st I.E.
Energy absorbed to remove 1 mole of electrons from 1 mole of gaseous atoms to form 1 mole of singly charged gaseous cations.
Second Ionisation Energy, 2nd I.E.
Energy absorbed to remove 1 mole of electrons from 1 mole of singly charged gaseous cations to form 1 mole of doubly charged gaseous cations.
e.g. Al (g) → Al+ (g) + e– 1st I.E. = +577 kJ/mol
e.g. Al+ (g) → Al2+ (g) + e– 2nd I.E. = +1820 kJ/mol
- Outermost electrons get removed first
- Ionisation is endothermic process, i.e. ∆H = +ve
- Further removal of electrons becomes more difficult as they are being removed from an increasingly positively charged species i.e. 1st I.E. < 2nd I.E. < 3rd I.E.
Electron Affinity, E.A.
Enthalpy change when 1 mole of electrons is added to 1 mole of atoms or ions in the gaseous state
First Electron Affinity, 1st E.A.
Energy released when 1 mole of electrons is added to 1 mole of gaseous atoms to form 1 mole of singly charged gaseous anions.
Second Electron Affinity, 2nd E.A.
Energy change when 1 mole of electrons is added to 1 mole of singly charged gaseous anions to form 1 mole of doubly charged gaseous anions.
e.g. O (g) + e– → O– (g) 1st E.A. (O)
e.g. O– (g) + e– → O2– (g) 2nd E.A. (O)
- 1st E.A. is usually negative and is the measure of affinity of the atom for the incoming electron. The stronger the affinity for the electron, the more energy is given off in the formation of the anion and the more negative the E.A.
- 2nd E.A. is always positive. Energy is required to overcome the repulsion between the incoming electron and the already negatively charged ion.
Bond Energy, B.E.
Energy absorbed to break 1 mole of a covalent bond between 2 atoms in the gaseous state.
e.g. A-A (g) → 2 A (g) B.E. (A-A)
e.g. A-B (g) → A (g) + B (g) B.E. (A-B)
Standard condition (under the new H2 Chemistry syllabus code 9729) refers to:
- Pressure of 1 bar
- Temperature of 298K
- Substance in its most stable physical form e.g. H2 (g), H2O (l), Br2 (l), NaCl (s), etc
I hope you find the content easy for your understanding and if you have any questions, leave me a comment below. Feel free to share this blog post with your friends.
Do stay tuned to the upcoming posts as i will be discussing on the applications of Enthalpy Changes using Hess’ Law. I will also be sharing the concepts and applications involved in Entropy Changes and Gibbs Free Energy.
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- Chemical Energetics: Experimental Method to Determine Enthalpy Change of Combustion
- Atomic Structure: Ionisation Energy Video Tutorial
- Chemical Energetics: Application of Hess’ Law & Energy Cycle Diagram
- Concentrations of Solutions in Atoms, Molecules & Stoichiometry
- Chemical Energetics: Gibbs Free Energy in Thermodynamics