BIO Lecture 2: Kinetics

In our last lecture we talked about how DeltaG says nothing about the rate of a reaction, only if it is spontaneous or not.  The study of reaction rates is called kinetics.  

All reactions proceed through a transient intermediate that is unstable and takes a great deal of energy to produce (called activation energy).  More specifically, activation energy is the barrier that prevents many reactions from taking place even though they may have a negative DeltaG.  Therefore, lowering the activation energy increases the rate of the reaction (lowering of the activation energy is done through catalysts--more on these below).  This is because we have essentially reduced the amount of energy that is required to reach the transition state.  Also, the lower the energy barrier, the more likely the reaction is to occur because more reactant molecules have enough energy to achieve the transition state.  

For clarification, a transition state exists for a very, very short time.  Because of this, we really don't know what they look like because we are unable to isolate them via analytical techniques.

A catalyst lowers the activation energy without changing the DeltaG.  It does this by stabilizing the transition state.  It is also not consumed in the reaction and is therefore regenerated with each cycle.  Note that enzymes are catalysts and they increase the rate of reactions that have a negative DeltaG.

Thermodynamically unfavorable reactions (positive DeltaG) are driven forward by reaction coupling.  This means that a very favorable reaction is used to drive an unfavorable one (i.e. ATP hydrolysis).  This is possible because free energy changes are additive.