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Chemistry Honors/AP

Chemistry 14: 5-6

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WebLecture

Scientific Method

Outline

Activation energy

Reactions occur when reactant molecules collide under exactly the right conditions to break the bonds that hold them together, and allow the individual atoms to reassemble in a different, more energy-efficient pattern. These conditions include the velocity of the colliding molecules and the angle and impact site of the collision. For example, two molecules might collide with enough energy but not hit the right site for the collision to be effective in separating the atoms in either molecule.

The Swedish chemist Svante Arrhenius defined amount of energy required for an effective collision as the activation energy of the reaction. As with enthalpy, this energy is measured in kJ/mole. Some of the energy can come from heat, which increases the impact velocity of the reactant molecules. In general, increasing the temperature of a solution will increase the rate of the reaction so that the rate constant is dependent on the minimum energy for the reaction at a specific temperature:

k = Ae-Ea/RT [Arrhenius Equation]

Here, A is the frequency at which effective collisions occur. T is the absolute temperature (in Kelvin), R is the gas constant. Ea is the activation energy. The fraction E/RT is greater than 1, so e-E/RT will always be less than one. It gives us the portion of the total sample that has enough activation energy to participate in the reaction.

It is possible under some circumstances for Ea to be negative. Such reactions are "barrierless": they will be spontaneous when any collisions occur between molecules of the reacting species.

If a catalyst is present, it will reduce the amount of activation energy required, so that a larger number of molecules will experience effective collisions.

The way in which concentration levels affect the rate of a reaction can indicate something about the reaction mechanics, that is, the rate can tell us how individual molecule interact during the reaction.

Reaction Mechanisms

Many important chemical reactions are actually multiple reactions, in which the products of each step become the reactants for the next step. This is spectacularly true in two of the most critical biological reactions, cellular respiration and photosynthesis, where each step of the reaction "cycle" requires its own unique enzyme to catalyze the particular reaction at temperature a cell can survive.

Consider the simple two-step reaction between nitrogen oxide and carbon

NO2 + CO → NO + CO2

From the equation, we might conclude that we have a reaction between one CO molecule and one NO2 molecule, which would give us a first order reaction.

Experimentally, though, it turns out that the reaction rate law is rate = k[NO2]2: that is, the reaction is second order and dependent only on the concentration of NO2 available. This leads chemists to realize that something else is going on. The reaction doesn't occur in a simeple one step manner: there is an intermediate step involving collisions between the NO2 molecules that must occur for the CO collision to be effective:

Step 1NO2 + NO2 → NO3 + NOSlow
Step 2NO3 + CO → NO2 + CO2Fast
Overall ReactionNO2 + CO → NO + CO2Slow

The reaction rate is actually determined by this "hidden" reaction step, called the rate-determining step since in this case, it controls the rate of the overall reaction. Experimental measurement of the rate as a function of concentration of reactants is fundamental in determining the individual reaction steps that make up an overal "reaction mechanism".