Hello everyone! In this article I am going to group the main types of enthalpy change reactions to the best of my abilities, as well as summarise what they do and how to use them. Enthalpy change to put in layman terms, is the energy change to a system after any reaction. Hence if a system releases heat it is losing energy (hence exothermic and negative enthalpy change), and if it absorbs heat it is gaining energy(hence endothermic and positive enthalpy change). This is based on topic 5 of IB HL chemistry syllabus for reference.
Section 1: Changing of compounds into “common compounds/elements”
The first two enthalpy changes are general types of reactions which can be done for (almost) all compounds. These reactions are commonly used in enthalpy change calculations as part of Hess’s Law which I will cover in the following article. They are standard enthalpy change of formation and standard enthalpy change of combustion.
1. Standard Enthalpy Change of Formation
Definition: Enthalpy change when one mole of a compound is formed from its elements in their standard states under standard conditions.
As can be seen from the name, this is the formation of any compound from its base elements. Generally, it can be seen as elements $\rightarrow$ compound, an example can be the formation of carbon dioxide.
$C(s) + O_2(g) \rightarrow CO_2$(g)
2. Standard Enthalpy Change of Combustion
Definition: Enthalpy change when one mole of substance in its standard state is completely burned in excess oxygen under standard conditions.
This reaction is most commonly used for organic compounds, as most of them are combustible. This reaction follows the general trend of compound $\rightarrow$ combustion products. In the case of common organic compounds such as alkanes, alkenes or alcohols, the combustion products usually comprise of water and carbon dioxide, as most of them only contain such elements. One example would be the combustion of methane.
$CH_4 (g) + O_2(g) \rightarrow CO_2(g) + H_2O$(l)
What is common between these types of reactions is that they provide a link between compounds and their substituent elements/commonly found compounds. This would be effective as these reactions bring the products and reactants to common elements/compounds, acting as “common ground” in any reaction in Hess’s law. What is important also to take note here is that this reaction does NOT specifically result in a change of state, instead every compound is in their standard state. Standard state is the state of the compound or element found under standard conditions, which is 298K and 100kPa.
Section 2: Changes of state and conversion into individual atoms
The second section is mainly focused on reactions that overall changes the state and/or breaks down the product into its substituent atoms, mainly bond enthalpies and standard enthalpy change of atomisation. These reactions are either used to serve a similar purpose: to convert an element/compound in its standard state to gaseous atoms. This is useful for the third section of enthalpy changes.
3. Bond Enthalpies
Definition: Energy required to break one mole of bonds in gaseous molecules, under standard conditions.
This is only used where the compound in question in gaseous state and converts gaseous molecules into their substituent atoms. The “quantifying factor” here is the bond, where it is the energy associated with breaking/forming one mole of a bond, and in order to fully break down a compound, we need to sum up the bond enthalpies of the individual bonds in the compound. This is useful as a conversion of a gaseous compound (like chlorine gas) into individual elements (2 chlorine atoms).
$Cl_2(g) \rightarrow 2Cl (g)$
4. Standard Enthalpy Change of atomisation
Definition: Enthalpy change when one mole of gaseous atoms is formed from the element in its standard state.
This reaction is the main, all-encompassing reaction which converts all compounds/elements into their gaseous, atomic form. The important “quantifying factor” here is the formation of one mole of gaseous atoms. Hence the main use of this reaction is to convert elements in their standard state to their gaseous, atomic form. An example includes halogens such as chlorine, where the reaction is as shown.
$\frac{1}{2}Cl_2(g) \rightarrow Cl (g)$
Hopefully from the examples above the differences between the “quantifying factors” of both reactions are evident!
Section 3: Adding or removing electrons from atoms
This section is very straightforward once you are familiar with the reactions in section 2. Two important “conditions”, gaseous, and also one mole of singular atoms. I hope that you can see why typically every compound has to be converted by a reaction in section 2 before using either of these reactions. The definitions are as shown
5. $1^{st}$ Electron Affinity
Definition: Enthalpy change when one mole of gaseous electrons is added to one mole of gaseous atoms to form one mole of singly negatively charged gaseous ions.
$Cl(g) + e^- \rightarrow Cl^-(g)$
6. $1^{st}$ Ionisation Energy
Definition: Minimum energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of singly positive charged gaseous ions.
$Na(g) \rightarrow Na^+(g) + e^-$
If the question states $2^{nd}, 3^{rd}…n^{th}$ Electron affinity or ionisation energy, it just represents the number of electrons added ore removed respectively. This is can be linked to the next section involving ions and ionic compounds.
Section 4: Reactions involving Ionic Compounds.
Now, once the atoms have been converted into their gaseous and ionic form, they are the components of a ionic compound, and hence can be used in the following 3 enthalpy changes.
7. Enthalpy Change of hydration
Definition: Enthalpy change when one mole of gaseous ions forms one mole of hydrated ions in water under standard conditions
$Na^+(g)$ $\rightarrow$ $Na^+(aq)$
8. Enthalpy Change of Solution
Definition: Enthalpy change when one mole of solute is dissolved in a solvent to infinite dilution under standard conditions
$NaCl(s)$ $\rightarrow$ $Na^+(aq)$ + $Cl^-(aq)$
9. Lattice energy
Definition: Enthalpy change when one mole of a solid ionic compound is separated into gaseous ions under standard conditions.
$NaCl(s) \rightarrow Na^+(g) + Cl^-(g)$
Personally, I find this section rather difficult to memorise, because of the similar names and states. I found that specifically remembering that enthalpy change of solution as representing dissolution, the act dissolving any ionic compound in water. This means that there can’t just be one ion involved (like in hydration), or any gaseous molecules involved, as dissolving implies a change from solid state to aqueous state.
Then from there I work backwards to derive the other two reactions, Lattice energy which must involve an ionic lattice and hence is the breaking down of the solid ionic lattice into gaseous ions, and enthalpy change of hydration to be adding water to gaseous ions to form hydrated ions. Just a disclaimer this is just how I remember the reactions, what works for me may not work for everyone, but I hope that everyone has learnt something from this article :)
If you have any questions feel free to msg the chemistry tag in our discord channel or dm me @cowcowchem !!
Author: Cowcowchemistry
