Unit 1 – Energetics
Enthalpy Changes
Chemical reactions usually have enthalpy changes.
When a chemical reaction happens, there is usually a change in
energy.
Enthalpy change ΔH = The heat energy transferred in a reaction at
constant pressure – The units are kj mol -1
The δ – delta symbol by an enthalpy change is telling you that
substances were in their standard states and the measurement was
made under standard conditions. Standard conditions are 100kPa
pressure and a stated temperature.
Reactions can be either exothermic or endothermic
Exothermic reactions give out energy.
ΔH is negative.
Oxidation reactions are usually exothermic.
- The combustion of fuel like methane.
- The oxidation of carbohydrates, such as glucose in respiration.
Endothermic reactions absorb energy.
ΔH is positive.
- Thermal decomposition of calcium carbonate
- The main reactions of photosynthesis are also endothermic.
Reactions are all about breaking and making bonds.
When reactions happen, reactant bonds are broken and product
bonds are formed.
1. You need energy to break bonds, so bond breaking is
endothermic.
Stronger bonds take more energy to break.
, 2. Energy is released when bonds are formed, so bond making is
exothermic.
Stronger bonds release more energy when they form.
3. The enthalpy change for a reaction is the overall effect of
these 2 changes – If you need more energy to break bonds
than is released when bonds are made, ΔH is positive. If its less
ΔH is negative.
Mean Bond Enthalpies are no exact
1. Bond enthalpy is the energy required to break bonds.
2. The energy needed to break a bond depends on the environment its
in.
3. In calculations you use mean bond enthalpy – the average energy
needed to break a certain type of compound over a range of
compounds.
E.g Water has 2 O-H bonds.
For the first bond H-OH +492 kj mol -1
For the second bond H-O +428 kj mol -1
Mean bond enthalpy (492+428) / 2 = 460 kj mol -1
4. Breaking bonds is always endothermic, so the mean bond
enthalpies are always positive.
Enthalpy Changes can be calculated using mean bond enthalpies.
In any chemical reaction, energy is absorbed to break bonds and given
out during bond formation. The difference between the energy absorbed
to break bonds and released in making bonds is the overall enthalpy
change of reaction.
Enthalpy Change of reaction = Total energy absorbed – Total energy
released.
As they use average values, enthalpy changes calculated using mean
bond enthalpies aren’t exact – They are slightly less accurate than
enthalpy change value calculated using Hess’s Law.
There are different types of ΔH:
Standard Enthalpy of Formation = The enthalpy change when 1 mole of a
compound is formed from its elements in their standard states under
standard conditions.
Standard Enthalpy of Combustion = The enthalpy change when 1 mole of
a substance is completely burned in oxygen under standard conditions.
Calorimetry
Enthalpy Changes
Chemical reactions usually have enthalpy changes.
When a chemical reaction happens, there is usually a change in
energy.
Enthalpy change ΔH = The heat energy transferred in a reaction at
constant pressure – The units are kj mol -1
The δ – delta symbol by an enthalpy change is telling you that
substances were in their standard states and the measurement was
made under standard conditions. Standard conditions are 100kPa
pressure and a stated temperature.
Reactions can be either exothermic or endothermic
Exothermic reactions give out energy.
ΔH is negative.
Oxidation reactions are usually exothermic.
- The combustion of fuel like methane.
- The oxidation of carbohydrates, such as glucose in respiration.
Endothermic reactions absorb energy.
ΔH is positive.
- Thermal decomposition of calcium carbonate
- The main reactions of photosynthesis are also endothermic.
Reactions are all about breaking and making bonds.
When reactions happen, reactant bonds are broken and product
bonds are formed.
1. You need energy to break bonds, so bond breaking is
endothermic.
Stronger bonds take more energy to break.
, 2. Energy is released when bonds are formed, so bond making is
exothermic.
Stronger bonds release more energy when they form.
3. The enthalpy change for a reaction is the overall effect of
these 2 changes – If you need more energy to break bonds
than is released when bonds are made, ΔH is positive. If its less
ΔH is negative.
Mean Bond Enthalpies are no exact
1. Bond enthalpy is the energy required to break bonds.
2. The energy needed to break a bond depends on the environment its
in.
3. In calculations you use mean bond enthalpy – the average energy
needed to break a certain type of compound over a range of
compounds.
E.g Water has 2 O-H bonds.
For the first bond H-OH +492 kj mol -1
For the second bond H-O +428 kj mol -1
Mean bond enthalpy (492+428) / 2 = 460 kj mol -1
4. Breaking bonds is always endothermic, so the mean bond
enthalpies are always positive.
Enthalpy Changes can be calculated using mean bond enthalpies.
In any chemical reaction, energy is absorbed to break bonds and given
out during bond formation. The difference between the energy absorbed
to break bonds and released in making bonds is the overall enthalpy
change of reaction.
Enthalpy Change of reaction = Total energy absorbed – Total energy
released.
As they use average values, enthalpy changes calculated using mean
bond enthalpies aren’t exact – They are slightly less accurate than
enthalpy change value calculated using Hess’s Law.
There are different types of ΔH:
Standard Enthalpy of Formation = The enthalpy change when 1 mole of a
compound is formed from its elements in their standard states under
standard conditions.
Standard Enthalpy of Combustion = The enthalpy change when 1 mole of
a substance is completely burned in oxygen under standard conditions.
Calorimetry
