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Summary Complete Energetics I Revision Notes (A Level Edexcel) £2.99
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  3. PEARSON (PEARSON)
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  5. Chemistry 2015
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  7. Unit 8 - Energetics I

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Summary Complete Energetics I Revision Notes (A Level Edexcel)
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Comprehensive study guide for Chemistry A Level, made by an Oxford Biochemistry student with all 9s at GCSE and 3 A*s at A Level! Information arranged by spec point. Notes written using past papers, textbooks and more.

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  • March 20, 2021
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ENERGETICS I
1. know that standard conditions are 100 kPa and a specific temperature, usually 298
K
Chemical bonds are the forces of attraction that bind atoms together
- Chemical energy lies within these chemical bonds.
- It is a form of potential energy.
- In chemical reactions, energy is changed from one form to another.
o E.g. chemical energy may change to thermal energy.
- No energy is lost – it is converted from one form to another.
Thermochemistry is the study of energy changes
- We can use thermochemistry to decide whether or not reactions are likely to occur and
explain the stability of compounds.
- The ‘system’ describes the material or mixture of chemicals being studied. Everything
else around the system is called the ‘surroundings.’
o This includes the apparatus, the air in the laboratory etc.
o The system includes the atoms and bonds involved in the chemical reaction.
- We often use ‘standard conditions’ in chemistry so that we can talk about the changes in
chemical reactions under set conditions.
o These standard conditions are: 100 kPa and 298 K (= 25oC).
- Elements should be written in their most stable state for the standard enthalpy change of
formation.
o E.g. carbon should be graphite (not diamond, fullerene etc.)
o Any aqueous solutions should be at a concentration of 1 mol dm-3.



2. know that enthalpy change is the heat energy change measured at constant
pressure
Enthalpy, H, is the thermal energy that is stored in a system.
- We can’t measure the direct enthalpy of products and reactants.
- Instead, we can measure the amount of energy that is absorbed or released to the
surroundings.
- We use the enthalpy change to do this.
Enthalpy change is the energy transferred
between a system and its surroundings when a
change happens at constant pressure.
- It is defined as: ‘the heat energy change
measured at constant pressure.’
- The symbol for enthalpy change is ΔH.
- Enthalpy change is measured in kJ mol-1.
- You must express these values under standard conditions, using the symbol ΔH⦵.


3. be able to construct and interpret enthalpy level diagrams showing an enthalpy
change, including appropriate signs for exothermic and endothermic reactions

,Exothermic changes
- Exothermic changes give out energy that often heats up
the surroundings.
- The enthalpy of the products is smaller than the
enthalpy of the reactants.
- Making bonds is an exothermic reaction (MEXican).
- An example of an exothermic reaction is the combustion
of hydrocarbons.
o Hydrocarbons are fuels; they release energy when burnt.
o CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (l), ΔH = – 890 kJ mol-1
- Another example is the reaction between solid
calcium oxide and water. This forms calcium
hydroxide solution.
o CaO (s) + H2O (l)  Ca(OH)2 (aq), ΔH = –
1067 kJ mol-1
o 1067 kJ energy is given out in the reaction,
assuming that one mole of CaO reacts with
one mole H2O.
o The system loses energy by heating the
surroundings.
- Anaerobic respiration is the release of energy from
glucose in the presence of oxygen.
o C6H12O6 (aq) + 6 O2 (g)  6 CO2 (g) + 2 H2O (l), ΔH = – 2800 kJ mol-1
- This loss of energy from the system means that ΔH is negative.

Endothermic changes
- Endothermic changes take in energy from the
surroundings.
- The enthalpy of the products is greater than the
enthalpy of the reactants.
- They are the opposite of exothermic changes.
- Breaking bonds is endothermic (BENDy).
o Melting and vaporisation are endothermic
changes of state as they involve the breaking of intermolecular forces between
the molecules.
- An example of an endothermic chemical change is
photosynthesis.
o Plants take in energy from the Sun in order to
convert CO2 and H2O to glucose.
o This is the reverse of aerobic respiration.
o 6 CO2 (g) + 2 H2O (l)  C6H12O6 (aq) + 6 O2 (g),
ΔH = + 2800 kJ mol-1
- Thermal decomposition reactions also take in heat
energy from their surroundings.
o CaCO3 (s)  CaO (s) + CO2 (g), ΔH = + 178 kJ
mol-1
- The system takes in energy from the surroundings.
o This means that the energy change, ΔH, is positive.
o The products are at a higher energy level than the reactants.

, Enthalpy level diagrams
- These have no x axis. They show ‘enthalpy’ on the y axis.
- They also do not show the activation energy, unlike enthalpy profile diagrams.
o Energy profile diagrams are ‘high profile.’ They have a high hump.
o Enthalpy level diagrams are ‘level.’ They do not have a hump.
- They show the enthalpy change, ΔH, between the reactants and the products.
Enthalpy profile diagrams
- These show the ‘progress of reaction’ on the x axis and the ‘enthalpy’ on the y axis.
- These show the activation energy.
- The activation energy is the minimum energy that the reacting particles must possess
before the can react. Its units are kJ mol-1.
Activation energy
- For an exothermic reaction, the sign of the enthalpy change is negative and the products
are at lower enthalpy than the reactants.
- For this type of reaction, the products are said to be thermodynamically stable with
respect to the reactants.
- Many exothermic reactions occur spontaneously and it is often assumed for simplicity
that the more exothermic the ΔHreaction, the more likely the reaction is to occur in the
forward direction.
- However, some exothermic reactions do not occur spontaneously.
o For example, sucrose is found in cane sugar. Despite the highly negative
standard enthalpy change of the combustion of sucrose, cane sugar is not
observed to oxidise in air at ordinary temperatures.
o C12H22O11 (s) + 12 O2 (g)  12 CO2 (g) + 11 H2O (l), ΔH⦵c = 6000 kJ mol-1.
- The direction of change may depend on the conditions of temperature and pressure.
- One example is the condensation of a vapour such as steam.
o Steam condenses to water below 100oC and energy is given out. This is an
exothermic change.
o H2O (g) → H2O (l), ΔH = −44 kJ mol−1
o At temperatures above 100 °C, the change goes in the opposite direction and this
process is endothermic.
- There are some examples of endothermic reactions which occur readily under normal
conditions. So some reactions for which ΔH is positive can happen.
o One example of this is the reaction of citric acid solution with sodium
hydrogencarbonate. The mixture fizzes vigorously and cools rapidly.
o This suggests that there are other factors that determine the direction of change.
- Some exothermic reactions never occur because the rate of reaction is so slow and the
mixture of reactants is effectively inert.
o For example, the change from diamond to graphite is exothermic, but diamonds
do not suddenly turn into black flakes.
- This can also be explained when looking at the energy profile diagram for this reaction.
- The minimum energy that the sucrose and oxygen molecules must possess before they
can collide successfully (to form carbon dioxide and water) is called the activation energy
(EA) for the reaction.
o At ordinary temperatures, the molecules of sucrose and oxygen do not have
sufficient energy to overcome the activation energy barrier and so the reaction
does not occur.

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