Chemical Energetics and Kinetics

Why reactions do not all happen at the same rate

Some reactions are extremely rapid, whereas others proceed slowly. Reacting particles must first meet, so collisions are necessary. However, not every collision causes a reaction: the particles must collide with at least a minimum amount of energy.

This minimum energy is the activation energy, often written $E_a$. It is an energy barrier between reactants and products. If colliding particles have less energy than $E_a$, they separate without reacting. A collision with sufficient energy can lead to bonds being rearranged and products being formed; this is a successful collision.

The rate of a reaction depends on how many successful collisions occur in a given time. Reactions may therefore have different rates because their energy barriers differ, although collision frequency and the energies of the colliding particles also matter. A reaction with a large activation energy can be slow because fewer collisions have enough energy to overcome its barrier.

Reaction profiles show the barrier

A reaction profile plots energy vertically against progress through a reaction horizontally. The reactants begin at one energy level and the products finish at another. The curve rises to a maximum because energy must initially be supplied to overcome the activation-energy barrier—even when the reaction eventually transfers energy to its surroundings.

The activation energy is the vertical difference between the reactants’ level and the top of the curve. It is not the difference between reactants and products; that second difference represents the overall energy transfer.

A catalyst makes a reaction faster by providing a different pathway with a lower activation energy. At the same conditions, more collisions can then overcome the smaller barrier. The catalyst is not used up overall, and it does not change the starting and finishing energy levels.

Energy changes as bonds rearrange

Chemical reactions involve breaking bonds in the reactants and forming bonds in the products. Breaking bonds requires energy, whereas forming bonds releases energy. The balance between these two processes determines the direction of the overall energy transfer.

In an exothermic reaction, forming the product bonds releases more energy than is required to break the reactant bonds. Energy is transferred from the reaction system to the surroundings, so the products have a lower energy than the reactants. Combustion and acid–alkali neutralisation are examples; the surroundings often become warmer.

In an endothermic reaction, breaking the reactant bonds requires more energy than is released when product bonds form. Energy is transferred from the surroundings to the reaction system, so the products have a higher energy than the reactants. Thermal decomposition reactions provide an example and require a continuing energy supply.

Energy is conserved

The rise or fall in the chemical energy of the reacting substances does not mean that energy has appeared or vanished. Energy is conserved: it can be transferred, but it can be neither created nor destroyed.

It is important to define the boundary being considered. The reaction system contains the reacting substances; the surroundings include such things as the container, solution and nearby air. In an exothermic reaction, a decrease in the system’s energy is matched by energy transferred to the surroundings. In an endothermic reaction, an increase in the system’s energy is matched by energy transferred from the surroundings. The total energy of the system and surroundings remains constant.

Activation energy and overall energy transfer answer different questions. Activation energy explains the barrier that affects whether collisions produce a reaction and therefore helps explain rate. The difference between reactant and product energy shows whether the completed reaction transfers energy to or from its surroundings.

The energy barrier

• Particles must collide to react.
• A reaction occurs only if a collision has at least the activation energy, $E_a$.
• More successful collisions per second means a faster reaction.
• Different barriers help explain why reactions have different rates.

Reading a reaction profile

• Vertical axis: energy; horizontal axis: progress of reaction.
• $E_a$ runs from the reactants’ level to the peak.
• A catalyst provides another pathway with a lower $E_a$ and is not used up overall.
• A catalyst does not change the reactant or product energy levels.

Energy transfer

• Breaking bonds requires energy.
• Forming bonds releases energy.
• Exothermic: energy transfers from system to surroundings; products are lower in energy.
• Endothermic: energy transfers from surroundings to system; products are higher in energy.

Conservation

Energy is neither created nor destroyed. Any energy decrease in the reaction system is matched by a transfer to the surroundings, and any increase is supplied by the surroundings.