U3AOS1 Topic 2: Thermodynamics

In progress (updating)

The golden rule of thermochemistry - Law of energy conservation: 

Energy cannot be created or destroyed, only transformed between different forms of energy

Remember the terms endothermic and exothermic?

In the topic of thermochemistry, the principles of these terms are expanded further

Now, what is thermochemistry? 

Decomposing the word, thermo - meaning heat, chemistry - interactions between particles

A really important term to get familiarised with is enthalpy

Think of enthalpy as stored energy - in the form of chemical energy. Enthalpy is given the symbol H

How does a chemical reaction take place?

Particles collide with sufficient energy and correct orientation, intramolecular bonds are rearranged, involving energy changes (chemical energy) into new particles.

Reminder: Intramolecular bonds include metallic, ionic and covalent

This topic involves more quantification therefore its best to become familiar with unit conversions involving energy

J	10-3 kJ, 10-6 MJ, 10-9 GJ
kJ	103 J, 10-3 MJ, 10-6 GJ
MJ	106 J,103 kJ, 10-3 GJ
GJ*	109 J, 106 kJ, 103 MJ

Note: the units are case sensitive eg. make sure to write kJ instead of KJ

The difference between each unit incrementation is 1000

*rarely used

Study design dot point: comparison of exothermic and endothermic reactions, with reference to bond making and bond breaking, including enthalpy changes (∆H) measured in kJ, molar enthalpy changes measured in kJ mol-1 and enthalpy changes for mixtures measured in kJ g-1, and their representations in energy profile diagrams

Endothermic:

Reaction absorbs energy from external environment, thereby total enthalpy of products is greater than total enthalpy of reactants as energy is gained in the system

Analogy:

Kind of think of a vacuum cleaner, where the mass of dust sucked into the vacuum cleaner is energy absorbed. As more dust is sucked in, the vacuum cleaner becomes more filled with dust - representing the total enthalpy. The end vacuum cleaner is heavier than the vacuum cleaner at the beginning prior to the cleaning session, hence total enthalpy of products (after cleaning) is greater than reactants (prior to cleaning)

Represented as: 

Hproducts<Hreactants or H>0

Examples:

• Boiling water

• Photosynthesis

• Cooling pack

• Frying eggs

Exothermic reaction:

Reaction releases energy into the external environment, thereby total enthalpy of reactants is greater than total enthalpy of products as energy is reduced in the system

Hproducts>Hreactants or H<0

Examples:

• Acid base reaction

• Freezing water into ice cubes

• Campfire 

• Combustion

Thermochemical equations:

Study design dot point: combustion (complete and incomplete) reactions of fuels as exothermic reactions: the writing of balanced thermochemical equations, including states, for the complete and incomplete combustion of organic molecules using experimental data and data tables

• Balanced equations with states however with added H component

Eg. CH4(g) + O2(g) →CO2(g) + H2O(l) H=-890kJ/mol 

Note: Make sure to double check and ensure -ve sign is added and always keep water as liquid state as databook states combustion operates at SLC

Author's normal tip: When balancing combustion equations/thermochemical equations ensure to follow the order CHO, balance carbon first, then hydrogen, finally oxygen.

Note: State of fuels is never aqueous

Thermochemical equations could be written as integer coefficients or fraction coefficients, I recommend keep it as fraction as there are more chances for error eg. must multiply H if changed to integer coefficient from fraction.

Example:
a) C4H10(g) + 132O2(g) →4CO2(g) + 5H2O(l) H=-2880kJ/mol

b) 2C4H10(g) + 13O2(g) →8CO2(g) + 10H2O(l) H=-5760kJ/mol

Both are valid however I would recommend stick with thermochemical equation a)

Effects on H value:

Reversing chemical equation changes the sign of H

Eg.

C6H12O6(aq) + 6O2(g) →6CO2(g) + 6H2O(l) H=-2803kJ/mol

6CO2(g) + 6H2O(l) →C6H12O6(aq) + 6O2(g) H=+2803kJ/mol

Energy profile diagrams:

Below is an example of an exothermic reaction

Below is an example of an endothermic reaction

See any difference in shape?

The exothermic reaction has products enthalpy lower than reactants enthalpy whereas the endothermic reaction has products enthalpy higher than reactants enthalpy.

Reasoning for exothermic profile diagram shape:

Activation energy, which is the minimum quantity of energy required to initiate the breaking of intramolecular bonds is smaller than the energy released during the formation of bonds. Thereby there is a net decrease in enthalpy, resulting in a negative change in enthalpy value.

Reasoning for endothermic profile diagram shape:

Activation energy, which is the minimum quantity of energy required to initiate the breaking of intramolecular bonds is larger than the energy released during the formation of bonds. Thereby there is a net increase in enthalpy, resulting in a positive change in enthalpy value.

Different types of combustion:
Incomplete combustion	Complete combustion
O2 limiting reagent	O2 excess reagent
CO produced*	CO2 produced
Produces soot and black particles	No black particles produced
C2H6(g) + 52O2(g) →2CO(g) + 3H2O(l)	C2H6(g) + 72O2(g) →2CO2(g) + 3H2O(l)

*Also produces carbon (C) however rarely tested

Extended knowledge:

Chemical energy is the stored energy in chemical bonds between atoms and molecules (visited in U3AOS1 Topic 1)

Chemical energy source

Attraction between electrons + protons

Movement of electrons

Rotations and vibrations from bonds

Repulsion between nuclei

Repulsion between electrons

Formation of intramolecular bonds actually decreases chemical energy

A good way to think of this is to note that chemical energy is primarily from repulsion between atoms, however if bonds are formed this is minimised by releasing energy as shown after the transition state of the energy profile diagram

You might’ve looked at the energy profile diagram and wondered what the transition state is…

Transition state:

• State at which all intramolecular bonds are broken and the constituent atoms which make up the reactants are separated

• At this state, potential energy is at its greatest

Hess’ law:

The total change in enthalpy of a chemical process is determined only by the reactants and products and their respective quantities; it is independent of any intermediate steps in the process and so the enthalpies of intermediate steps may be summed to produce the enthalpy change of an overall reaction.

In simpler terms, we only care about the beginning and end of the reaction, not the reactions that take place in the middle.