3.2.1 (a,b,c) Enthalpy Changes

Syllabus

a) explanation that some chemical reactions are accompanied by enthalpy changes that are exothermic (ΔH, negative) or endothermic (ΔH, positive)

(b) construction of enthalpy profile diagrams showing differences in the enthalpy of reactants and products

{Activation energy in terms of the minimum energy required for a reaction to take place.}

What does this mean?

Enthalpy is a name for the energy of a system including the energy stored in bonds, the energy that is heat etc.

You never actually measure all the enthalpy in a system - just how it changes as a reaction occurs.

Entropy is another form of energy that is studied in Year 13 but which we can ignore for now.

Generally, reactions occur that release energy and allow substances to lower their internal energy.

ΔH is the symbol given to enthalpy changes and if you studied GCSE Chemistry you'll probably be familiar with the following:

Exothermic reactions release energy (generally heat) - making their surroundings warmer.

This released energy must come from inside the reactants.

And it is accounted for by there being different strength bonds in the reactants and the products.

If the reactants are releasing energy then the energy of the products must be lower.

Which is why the products are shown below the reactants.

And also why ΔH is always negative for all exothermic reactions.

Endothermic reactions absorb energy (generally heat) - making their surroundings cooler.

The difference is still accounted for by there being different strength bonds in the reactants and the products.

If the reactants are absorbing energy then the energy of the products must be higher.

Which is why the products are shown above the reactants.

And also why ΔH is always positive for all endothermic reactions.

In reversible reactions ΔH for the forward reaction is the same but opposite of ΔH for the backward reaction.

ΔH_exo = - ΔH_endo

Matches burn exothermically, releasing energy and allowing the energy of the products to be lower than that of the reactants.

But they don't spontaneously burst into flames unless the temperature is very hot.

Or you add some energy by striking them.

Why?

It is all about making and breaking bonds.

You must break bonds in the reactants (which absorbs energy) before you can make new bonds in the products and release energy.

The minimum energy input needed for the reaction to happen is the Activation Energy.

If the reactants had very strong bonds the Activation energy will be very high.

If the reactants had very weak bonds the Activation energy will be very low.

But all reactions have a positive activation energy.

Even in an endothermic reaction you need to put in energy first, it is simply that you don't get back more as you would in an exothermic reaction.

In fact you don't even get all the input back - which is why the products end up higher than the reactants.

Unlike the Enthalpy change for reversible reactions the Activation Energies are not just the opposite of each other.

In the left-hand reaction Activation Energy (forwards) = E₁

& Activation Energy (backwards) = E₁ - E₃

In the right-hand reaction Activation Energy (forwards) = E₄

& Activation Energy (backwards) = E₄ + E₆

_{The shortest Powerpoint ever}

Google Presentation

_Videos

Back to 3.2.1?

Click here