Archive for July, 2013

This is part 2 of a four part series in the Energy Diagram Module. Stay tuned for the other parts!

To see part 1 click here.

In order to talk about energy of the reaction, a few key concepts are needed.

• If the products have less potential energy than the reactants, the reaction will release a net amount of energy (an exothermic reaction).
• If the products are higher in energy than the reactions, the reaction will consume a net amount of energy (an endothermic reaction).

But, in order to predict how well a reaction will progress, or how spontaneous the reaction will be, enthalpy is insufficient to make this estimation.  Therefore, we rely on the thermodynamic calculation of Gibbs Free Energy (ΔG⁰) which is represented by the equation;

ΔG⁰ = ΔH⁰ – TΔS

The components of Gibbs Free Energy are:

• Enthalpy, ΔH⁰ – The heat consumed or released by the reaction.
• Temperature, T – Temperature of the system.
• Entropy, ΔS.  The change in the degree of disorder.

The sign of the reaction indicates the release (negative) or absorption (positive) of heat during the reaction. Therefore, if ΔG⁰ is negative the reaction is always spontaneous.  Similarly, if ΔG⁰ is positive the reaction is never spontaneous.  When ΔG⁰ equals zero, the reaction is at equilibrium. In accounting for these additional thermodynamic properties using Gibbs Free Energy, different terms are used.

• If a reaction has a negative  ΔG⁰ , it is therefore spontaneous and is said to be exergonic.
• Conversely, a reaction that has a  positive  ΔG⁰  is not spontaneous is considered endergonic.

Graph 2

(click on photo to enlarge)

With the addition of the temperature variable in the Gibbs Free Energy equation, it is easy to see that an endergonic reactions can be driven forward simply by increasing the temperature of the reaction so that the term TΔS is more negative than ΔH⁰, thus making ΔG⁰ negative and making the reaction spontaneous!

This is part 1 of a four part series in the Energy Diagram Module.

In order to talk about reaction mechanisms, we have to first understand what makes up a reaction. It may help to think of a reaction as a landscape.  Let us look at the components in Graph 1 below.

• the x-axis as the reaction proceeds (the reaction coordinate).
• The potential energy of the molecules throughout the reaction along the y-axis.
• The shape of the line graph represents the pathway (or mechanism) by which the reactants must take to reach the products.
• Using the landscape analogy, reactants and products that are stable can be thought of as the “valleys”, which have some potential energy (refer to Graph 1).
• As reactants collide during the course of a reaction, they must go over a “hill” in order to each the product of the reaction.  This hill represents the energy needed to break the existing bonds in the reactions (the activation energy, ΔG^‡)
• At the top of the hill, some hybrid molecule (the transition state, [ ]‡) is formed which represents the simultaneous (often referred to as concerted) bond breaking and bond forming.  Because of the hybrid nature of the molecule, transition states are NOT isolatable.

The activation energy is usually provided by sum of the surrounding system and the collision of the reactant molecules. For reactions with high activation energy barriers, a Bunsen burner or heating element can be used to add energy.  Looking at Graph 2, as transition states form, the reaction can “roll” downhill both ways, backwards to reactants or forwards to products. As products begin to form, which as illustrated is much lower in energy than the reactants, the net release of energy (the heat of reaction, ΔH⁰) is released back to the surrounding system in the form of heat (enthalpy).

Graph 1

(click on photo to enlarge)

Graph 2

(click on photo to enlarge)