Reaction Kinetics Class Notes

Reaction Kinetics

(a.k.a. “Reaction Rates”)

Reaction kinetics is the study of rates (speeds) of chemical reactions and factors which affect the rates.

I.1 Measuring the Rate of a Chemical Reaction

Rates of reaction are usually expressed in terms of a concentration change per unit of time.

Reaction Rate = amount of product formed

time interval

= amount of reactant used

time interval

or, basically

= D amount

D time where D means “change in”

Example 1

If 16.0g of HCl are used up after 10min in a certain reaction, calculate the average rate of the reaction.

Sample

Questions:

Using Stoichiometry (Remember Chem 11?) to find reaction rates.

Example 1: H2(g) + Cl2(g) ® 2HCl(g)

0.250g of HCl are produced in 103 seconds. Find the rate of consumption of H2 and Cl2 in g/s.

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I.2 Methods of Measuring Reaction Rates

Cu(s) + 4HNO3(aq) Cu(NO3)2(aq) + 2H2O(l) + 2NO2(g) + HEAT

red-brown colourless blue colourless brown

To determine the rate of the above reaction, we could measure at least six different physical properties:

a) Color changes

The Cu(NO3)2 has a characteristic blue color. The intensity of the blue color could be measured in a spectrophotometer

Rate = D(color intensity)

D time

b) Temperature changes

Since the reaction is producing heat, we could measure and graph the temperature changes.

Rate = D(temperature changes)

D time

c) Mass changes

Since copper metal is the only solid present in the reaction, we could measure the rate at which the copper is used up by measuring its change in mass. Could we measure the change in mass of the whole beaker and its content?

Rate = D(mass)

D time

d) Concentration changes

Since the reaction uses up nitric acid, HNO3, we can measure the change in the [H+] with a pH meter.

Rate = D[H+(aq)]

D time

e) Pressure changes

Since the gas NO2 is being produced, we could measure the rate at which NO2 is produced by measuring the change in pressure.

Rate = D(pressure)

D time

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I.3 Factors That Affect Reaction Rates

The speed at which a reaction occurs is called the rate of reaction.

The rate of a reaction depends on both the nature of the reactants and on the conditions under which the reaction occurs:

a) Temperature

Increasing the temperature increases the average kinetic energy of the particles that is they jump around more. In general, each increase in temperature of 10 °C will double the reaction rate.

b) Concentration of the Reactants

The rates of both homogeneous and heterogeneous reactions are increased by an increase in concentration of reactants. Increasing concentration increases the probability of successful collisions.

c) Pressure

Pressure is just another way to define concentration. Gases only!!!!

d) The Nature of Reactants

Some reactions are naturally slow because the bonds involved are very strong and unreactive others are naturally fast because the reaction involves the breakage of weak bonds. We have no control over these big differences in the rate of a reaction.

Example: Rust vs TNT

e) The Ability for Reactants to Meet

(Surface Area and Phase Considerations)

i) Surface Area

The greater the surface area available for the reaction, the greater the rate of the reaction. This is because the greater the surface area, the greater the ability of the reactants to meet.

ii) Phase considerations

The phases (solid, liquid, gas) of reactants also has an effect on reaction rates.

Homogeneous Reactions

Reactions in which all the reactants are in the same phase is are homogenous reactions. These tend to be faster reactions because the reactant molecules are fully mixed and able to collide.

Examples:

·a reaction between two gases

·a reaction between two substances dissolved in water

·a reaction between two liquids which can easily dissolve in each other (are “miscible”)

Heterogeneous Reactions

Reactions in which the reactants are present in different phases are heterogeneous reactions. These reactions tend to be slower reactions because the reaction can only take place on the interface between the two phases. The greater the surface area, the greater the rate of reaction.

Examples:

·a reaction between a solid and a liquid

·a reaction between a liquid and a gas

·a reaction between a solid and a gas

·a reaction between two solids HOW CAN THIS BE?????????????

Another important consideration regarding phases is the actual phase in which a reactant occurs:

Solid phase reactions – tend to be slow because the reacting species cannot move freely.

Liquid phase reactions – tend to be fast because of the close proximity of the particles.

Gas phase reactions – tend to fast because of the speed of gas particles.

Aqueous ion reactions – have the fastest reaction rates because of their close proximity, their ability to move throughout the solution, and their strong positive-negative attraction.

f) Catalysts and Inhibitors

A catalyst is a chemical which can be added to a reaction to INCREASE the rate of the reaction without be used up in the reaction.

An inhibitor is a chemical which REDUCES the reaction rate. Examples of inhibitors are poisons and antibiotics.

Why would we want to Control of Reaction Rates?

There are many everyday situations that require the control of reaction rates. Some examples are:

a) Body chemistry requires exact temperature control to make sure that the reactions of life continue or do not go to fast. When the body produces a fever, it destroys many forms of bacteria which cannot survive at temperatures more than a degree or so above the normal body temperature of 37.0°C.

b) Fuels burn quickly in air and extremely rapidly in pure oxygen. If the concentration of oxygen is lowered, the rate of burning eventually drops to zero. This is the reason why smothering a fire is effective.

c) Enzymes are catalysts which regulate our body chemistry.

d) The metal in car bodies will quickly rust unless we can use paint or other protection to prevent oxygen from reacting with the iron in the car.

e) The cooking of food requires an increase in temperature. Many of the desirable reactions in the cooking process will not occur except at high temperatures.

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I.5 Reaction Rates and Collision Theory

Collision theory states that all molecules (and atoms) can be thought of as small, hard spheres which can bounce off each other.

For a chemical change to occur molecules must collide with sufficient energy to cause a rearrangement of atoms (activation energy). As well, the atoms must be orient correctly (geometry).

Consider the reaction: H2 + I2 ® 2HI

Figure 1.2

HI

+ +

H2 I2 activated

complex

Collision Theory and …

a) The Effect of Concentration – if we increase [H2] or [I2] we will have more collisions possible. Therefore the rate of the reaction will increase; that is the number of collisions per second increases.

b) The Effect of Temperature – If we increase the temperature then we increase the KINETIC ENERGY (KE) of the molecules and hence the speed at which they are moving. Because the molecules are moving faster they collide more often and with more energy, and hence the reaction rate increases.

Why does collision one

happen and collision

two does not?

Sample

Questions

I.6 Enthalpy Changes in Chemical Reactions

a) Bond Energies

The amount of energy required to break a bond between two atoms is bond energy.

If we break a bond, an amount of energy equal to the bond energy must be added to the bond:

Cl2(g) + 243kJ ® 2Cl(g)

Conversely, if two atoms form a bond, an amount of energy equal to the bond energy is released by the atoms:

2Cl(g) ® Cl2(g) + 243kJ

b) Reaction Heats

When a chemical reaction occurs, new molecules are formed as chemical bonds are broken and formed. Since molecules are different after the reaction, we would expect the POTENTIAL ENERGY (PE) of the products would differ from that of the reactants.

Therefore, two things are possible:

Either - Energy has been absorbed in order to form products. (Endothermic)

Or - Energy has been released in order to form products. (Exothermic)

Enthalpy (H) The heat (energy) content of a reaction

HREACTANTS The combined enthalpies of all reactants

HPRODUCTS The combined enthalpies of all products

The Sign of DH

i) Endothermic Reactions – products have more energy than reactants.

HPROD > HREACT

Therefore DHREACT 0

EXAMPLES: N2(g) + ½O2(g) + 25kJ ® N2O (g)

N2(g) + ½O2(g) ® N2O (g) ; DH = 25kJ

ii) Exothermic Reactions – products have less energy than reactants.

HPROD < HREACT

Therefore DH°RX 0

EXAMPLES:

½ H2 + ½ Cl2 ® HCl + 50kJ

½ H2 + ½ Cl2 ® HCl ; DH = -50kJ

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I.7 Kinetic Energy Distributions

Consider the reaction: C2H2 + H2 ® C2H4

At room temperature (25°C) the reaction rate is very slow.

At 100°C the reaction rate is faster, but still quite slow.

At 200°C the reaction rate is relatively rapid.

Molecules at room temperature and pressure undergo about 1010 collisions/second, so we can conclude that the lack of reactivity at room temperature is not due to a lack of collisions. Molecules need to have a sufficient amount of energy to react. The higher the temperature, the greater the average energy of the molecules, the more molecules with a sufficient amount of energy to react.

Figure 1.3 Graph of Kinetic Energy vs. Number of Molecules showing the difference between molecules at 25°C, 100°C, and 200°C.

Note: 10°C rule-of-thumb

I.8 Activation Energies

Activation Energy is the minimum energy requirement before a molecule can react.

Consider the reaction: H2 + I2 ® 2HI

(H2) and (I2) must collide with sufficient energy for a reaction to occur. If they do not collide with sufficient force they will bounce off of each other and no product (HI) will be formed.

Figure 1.4 Potential Energy Diagram

As molecules approach each other they repel the other and thus slowing down the molecules and converting their KE into PE. If the electrons can gain enough PE, the reactant bonds can be broken and the product bonds can be formed. At this critical moment a unstable, high energy molecule is formed – the ACTIVATED COMPLEX.

Activated Complex: the very short lived molecule that forms at the moment of collision when the reactant molecules are in the process of rearranging to form products.

Activation Energy: (Ea), the minimum potential energy required to change the reactants into the activated complex.

We can now explain why a reaction proceeds at a fast rate, a slow rate, or even an effectively zero rate. High Ea means slow rate and low Ea means high rate.

For a chemical reaction to occur, the reactant molecules must have two things

1. Sufficient KE

2. Correct alignment (a.k.a. reaction geometry)

If the reactants are not correctly aligned, the reaction will need more energy to be completed.

Consider the example of hydrogen gas reacting with chlorine gas to form hydrochloric acid:

H2 + Cl2 ® 2HCl

We now introduce an extremely important concept; so important is this concept, that it will have a major impact on the rest of your lives. *You can’t uncook a steak*

The following Potential Energy diagram is the same as those we have seen before, however it shows that for all reversible reactions there is a Forward Activation Energy(Ea(f)) and a Reverse Activation Energy(Ea(r)) where:

Ea(f) = the activation energy for the FORWARD reaction.

Ea(r) = the activation energy for the REVERSE reaction.

Note: Activation energy is always endothermic, that is energy must be added to get to the top of the hill.

Label on the graph, the Activation Energies of forward and reverse reactions and the DH.

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I.9 Reaction Mechanisms

Most reactions occur in more than one step.

Reaction Mechanism:The actual sequence of steps which make up an overall reaction is a reaction mechanism.

Consider the following reaction:

4HBr(g) + O2(g) ® 2H2O(g) + 2Br2(g)

It is unlikely (or the probability is extremely low) that 4 HBr molecules and 1 O2 molecule will collide simultaneously to undergo a chemical change. Complex reactions cannot go in a single step, but must consist of more than one step in getting from reactants to products. For example:

Step HBr(g) + O2(g) ® HOOBr(g) slow

Step HBr(g) + HOOBr(g) ® 2HOBr(g) fast

Step HBr(g) + HOBr(g) ® H2O(g) + Br2(g) fast

Sample

Problem

The slowest step in the mechanism determines the overall rate of the reaction and is called the rate determining step.

Step is the rate determining step in the above example.

To determine the overall reaction from the reaction mechanism, we simply add up all the steps in the reaction equations and cancel any species which occur on both sides of the final equation.

Sample

Problem

Reaction intermediates:

In the earlier example HOOBr and HOBr are “reaction intermediates” or simply “intermediates”. They are formed in one step are are subsequently consumed in another step and not found written in the overall reaction. Unlike an activated complex, intermediates can exist on their own.

NOTE: YOU WILL NEVER HAVE TO PREDICT A REACTION MECHANISM.