Chemistry Activation Energy

Explanation: Activation energy is the energy barrier that needs to be overcome for a reaction to proceed; the higher the activation energy, the slower the reaction.

• In chemistry, activation energy is a term introduced in 1889 by the Swedish scientist Svante Arrhenius to describe the minimum energy which must be available to a.
• The Activation Energy of Chemical Reactions. Only a small fraction of the collisions between reactant molecules convert the reactants into the products of the reaction.
• This is the definition of Ea or activation energy. This is the definition of Ea or activation energy. About.com. Food; Chemistry; Chemistry Glossary and Dictionary;.
• Effects of Enzymes on Activation Energy. However, if a catalyst is added to the reaction, the activation energy is lowered because a lower-energy transition state is.
• Activation energy, in chemistry, the minimum amount of energy that is required to activate atoms or molecules to a condition in which they can undergo chemical.

Activation energy, in chemistry, minimum energy needed to cause a chemical reaction. A chemical reaction between two substances occurs only when an.

Noun, Chemistry. 1. the least amount of energy required to activate atoms or molecules to a state in which they can undergo a chemical reaction.

For additional High School Chemistry prep, visit Free Online GED and try out their course for free. Basic Atomic Structure Atoms Ions.

Activation Energy

The Activation Energy of Chemical Reactions

Only a small fraction of the collisions between reactant molecules convert

the reactants into the products of the reaction. This can be understood by turning, once

again, to the reaction between ClNO2 and NO.

ClNO2 g NO g NO2 g

ClNO g

In the course of this reaction, a chlorine atom is transferred from one

nitrogen atom to another. In order for the reaction to occur, the nitrogen atom in NO must

collide with the chlorine atom in ClNO2.

Reaction won t occur if the oxygen end of the NO molecule collides with

the chlorine atom on ClNO2.

Nor will it occur if one of the oxygen atoms on ClNO2 collides

with the nitrogen atom on NO.

Another factor that influences whether reaction will occur is the energy

the molecules carry when they collide. Not all of the molecules have the same kinetic

energy, as shown in the figure below. This is important because the kinetic energy

molecules carry when they collide is the principal source of the energy that must be

invested in a reaction to get it started.

The overall standard free energy for the reaction between ClNO2

and NO is favorable.

ClNO2 g NO g NO2 g

Go

-23.6 kJ/mol

But, before the reactants can be converted into products, the free energy

of the system must overcome the activation energy for the reaction, as shown in the

figure below. The vertical axis in this diagram represents the free energy of a pair of

molecules as a chlorine atom is transferred from one to the other. The horizontal axis

represents the the sequence of infinitesimally small changes that must occur to convert

the reactants into the products of this reaction.

To understand why reactions have an activation energy, consider what has

to happen in order for ClNO2 to react with NO. First, and foremost, these two

molecules have to collide, thereby organizing the system. Not only do they have to be

brought together, they have to be held in exactly the right orientation relative to each

other to ensure that reaction can occur. Both of these factors raise the free energy of

the system by lowering the entropy. Some energy also must be invested to begin breaking

the Cl-NO2 bond so that the Cl-NO bond can form.

NO and ClNO2 molecules that collide in the correct orientation,

with enough kinetic energy to climb the activation energy barrier, can react to form NO2

and ClNO. As the temperature of the system increases, the number of molecules that carry

enough energy to react when they collide also increases. The rate of reaction therefore

increases with temperature. As a rule, the rate of a reaction doubles for every 10oC

increase in the temperature of the system.

Purists might note that the symbol used to represent the difference

between the free energies of the products and the reactants in the above figure is Go,

not Go.

A small capital G is used to remind us that this diagram plots the free energy

of a pair of molecules as they react, not the free energy of a system that contains many

pairs of molecules undergoing collision. If we averaged the results of this calculation

over the entire array of molecules in the system, we would get the change in the free

energy of the system, Go.

Purists might also note that the symbol used to represent the activation

energy is written with a capital E. This is unfortunate, because it

leads students to believe the activation energy is the change in the internal energy of

the system, which is not quite true. Ea measures the change in the

potential energy of a pair of molecules that is required to begin the process of

converting a pair of reactant molecules into a pair of product molecules.

Catalysts and the Rates of Chemical Reactions

Aqueous solutions of hydrogen peroxide are stable until we add a small

quantity of the I- ion, a piece of platinum metal, a few drops of blood, or a

freshly cut slice of turnip, at which point the hydrogen peroxide rapidly decomposes.

2 H2O2 aq 2 H2O aq

O2 g

This reaction therefore provides the basis for understanding the effect of

a catalyst on the rate of a chemical reaction. Four criteria must be satisfied in order

for something to be classified as catalyst.

Catalysts increase the rate of reaction.

Catalysts are not consumed by the reaction.

A small quantity of catalyst should be able to affect the rate of

reaction for a large amount of reactant.

Catalysts do not change the equilibrium constant for the reaction.

The first criterion provides the basis for defining a catalyst as

something that increases the rate of a reaction. The second reflects the fact that

anything consumed in the reaction is a reactant, not a catalyst. The third criterion is a

consequence of the second; because catalysts are not consumed in the reaction, they can

catalyze the reaction over and over again. The fourth criterion results from the fact that

catalysts speed up the rates of the forward and reverse reactions equally, so the

equilibrium constant for the reaction remains the same.

Catalysts increase the rates of reactions by providing a new mechanism

that has a smaller activation energy, as shown in the figure below. A larger proportion of

the collisions that occur between reactants now have enough energy to overcome the

activation energy for the reaction. As a result, the rate of reaction increases.

To illustrate how a catalyst can decrease the activation energy for a

reaction by providing another pathway for the reaction, let s look at the mechanism for

the decomposition of hydrogen peroxide catalyzed by the I- ion. In the presence

of this ion, the decomposition of H2O2 doesn t have to occur in a

single step. It can occur in two steps, both of which are easier and therefore faster. In

the first step, the I- ion is oxidized by H2O2 to form

the hypoiodite ion, OI-.

H2O2 aq I- aq

H2O aq

OI- aq

In the second step, the OI- ion is reduced to I- by

H2O2.

OI- aq H2O2 aq

O2 g I- aq

Because there is no net change in the concentration of the I-

ion as a result of these reactions, the I- ion satisfies the criteria for a

catalyst. Because H2O2 and I- are both involved in the

first step in this reaction, and the first step in this reaction is the rate-limiting

step, the overall rate of reaction is first-order in both reagents.

Determining the Activation Energy of a Reaction

The rate of a reaction depends on the temperature at which it is run. As

the temperature increases, the molecules move faster and therefore collide more

frequently. The molecules also carry more kinetic energy. Thus, the proportion of

collisions that can overcome the activation energy for the reaction increases with

temperature.

The only way to explain the relationship between temperature and the rate

of a reaction is to assume that the rate constant depends on the temperature at which the

reaction is run. In 1889, Svante Arrhenius showed that the relationship between

temperature and the rate constant for a reaction obeyed the following equation.

In this equation, k is the rate constant for the reaction, Z

is a proportionality constant that varies from one reaction to another, Ea

is the activation energy for the reaction, R is the ideal gas constant in joules

per mole kelvin, and T is the temperature in kelvin.

The Arrhenius equation can be used to determine the activation

energy for a reaction. We start by taking the natural logarithm of both sides of the

equation.

We then rearrange this equation to fit the equation for a straight line.

y mx b

According to this equation, a plot of ln k versus 1/T should

give a straight line with a slope of - Ea/R, as shown in the

figure below.

By paying careful attention to the mathematics of logarithms, it is

possible to derive another form of the Arrhenius equation that can be

used to predict the effect of a change in temperature on the rate constant for a reaction.

The Arrhenius equation can also be used to calculate what happens to the

rate of a reaction when a catalyst lowers the activation energy.