Chemistry 122 Chapter 18 Reaction Rate and Equilibrium 1
Our study of chemical kinetics will focus on two things:
1) the rate at which a chemical reaction takes place.
2) the factors which affect the speed of the reaction.
Reaction rate deals with how fast a reactant is used up during the chemical reaction. The reaction rate is the number of particles that react in a given time to form product. As the reaction proceeds, the total concentration (Molarity) of the reactant particles will decrease. The chemist uses the Collision Theory to explain chemical reactions.
Collision Theory
The Collision Theory uses the Kinetic theory to help explain how a chemical reaction occurs. The kinetic theory sates:
1. All matter is made up of particles (atoms, molecules or ions).
2. The tiny particles are in constant motion. These particles are in rapid, random, constant straight-line motion.
Particles in solids have very strong attractive forces, which hold the particles in place. These strong forces lead to vibratory motion in solids, ↔ the particles are very limited in motion. Solids have definite volume and shape.
In liquids the intermolecular forces are weaker and allow particles to slide by one another – think of marbles being poured. The forces are strong enough that the particles cannot escape from the other particles. There is a greater degree of freedom of movement – liquids do not have definite shape but have definite volume.
Gases consist of small particles that are far apart in comparison to their own size. There are no attractive forces between gas particles or between particles and the sides of the container with which they collide. Gases will expand to fill their container. Because of the great degree of freedom gases do not have definite shape or volume.
3. Particles collide with one another and the sides of the container.
4. Energy is conserved in these collisions, although one particle may gain energy at the expense of the other. Energy can be transferred in collisions among particles. Energy can be transferred in collisions among molecules. The average kinetic energy of all the molecules is proportional to the absolute temperature (K). The higher the temperature the faster the particles move.
The Collision Theory states:
1. Reacting particles must collide with each other. In order to react, reactant particles must collide with each other.
2. The reactant particles must collide hard enough to break old chemical bonds so that new chemical bonds will form. Bond breaking is an endothermic process. The minimum energy that colliding particles must have in order to react is called the activation energy.
3. The colliding reactants particles must have the correct spatial orientation so that the old bonds break and new bonds can form. During the reaction there is a temporary arrangement of the particles in which the old chemical bonds are breaking and the new chemical bonds are forming. This temporary arrangement of particles, more complex than either the reactants or products, is called the activated complex. The activated complex is also referred to as the transition state. The activated complex will only form if the collision has sufficient energy called the activation energy. The activated complex breaks down to produce either new products or the original reactants back. If new products are produced the collision was effective. An effective collision occurs when reactant particles collide with sufficient energy and correct spatial orientation to create the activated complex that breaks down into products.
Things that Effect Rate
The greater the number of collisions the greater the chance there will be for chemical reactions to occur. So whatever affects the number of collisions will also affect the rate at which chemical reactions occur. The greater the number of collisions the faster the reaction, extremely fast reactions are called explosions. Some reactions are so slow that we are not aware that they are occurring – rusting of cars.
1. Temperature is a measure of the average translational kinetic energy. A higher temperature means faster moving particles which leads to more and harder collisions that in turn yields faster reactions.
2. Concentration is an expression of moles of a substance (number of particles) per volume. When the reactants become more concentrated, the reactant particles are closer together, the particles collide more often causing a faster reaction.
3. Particle size refers to the size of the reacting particles. Particles of a solid can only collide when on the surface. Smaller particles yield a bigger surface area, allowing more reactant particles to collide. The smaller the reacting particles, the faster the reaction occurs. The smallest possible particles are molecules, atoms or ions. One way to get these individual particles is to dissolve the reactants so that they exist in the dissociated form. Dissolving speeds up the rate of the reaction. If the reactants are kept in solid form the reaction will be slow.
4. A catalyst is a substance that speeds up a reaction without being used up (Example: an enzyme). The catalyst speeds up reaction by giving the reaction a new alternate pathway.
The new pathway has a new activated complex that involves the catalyst. The new activated complex has lower activation energy. More molecules will have this energy of activation and the reaction goes faster. An inhibitor is a substance that blocks a catalyst.
An example of a catalyst is the use of platinum with reactions with hydrogen.
2H2(g) + O2(g) → 2H2O(l)
The hydrogen molecule bonds to surface of the platinum metal and break H-H bonds.
The oxygen bond starts to break when the hydrogen atoms start to bond with the oxygen molecule. The platinum reduces the energy requirement to break the hydrogen molecules into hydrogen atoms.
Reaction Rates
Rate equals the decrease in the concentration of a reactant (A) per unit of time. Because the concentration of A ([A]) is decreasing, the rate is given a negative sign (R = -[A]/t).
Rate is proportional to the molar concentration of A.
R α [A] Square brackets [ ] means concentration in Molarity (moles/liter)
This proportionality can be change to equality by introducing the proportionality constant (k). R = k[A] This relationship is shown in the graph:
Reaction Mechanisms
A reaction mechanism is a description of how a chemical reaction occurs. The mechanism is a path, a series of steps that a chemist uses to describe a chemical reaction. An elementary reaction occurs when the reactants are converted to products in a single step.
Most chemical reactions are complex and are viewed as a series of elementary reactions. Each step (elementary reaction) in the series has its own activated complex. Each activated complex has a different energy of activation. As each of these activated complex breaks down they create intermediate products that are immediately consumed as reactants in the next step. The intermediates do not appear in the net chemical equation, as the intermediates are consumed to create the final products. The net chemical equation does not give any information about the reaction mechanism.
A reaction mechanism is proposed based on experimental evidence. The chemist varies the concentration of reactants to study the effect on the rate of the chemical reaction. A potential energy profile for a three-step reaction shows the activation energy peak for each elementary reaction. The valleys indicate the formation of intermediates.
In any process that occurs in a series of steps, the rate of the process will be determined by the slowest step. If a group of people went for a walk and wanted to stay as a group, the group could go no faster than the slowest walker in the group. This reaction takes place in three steps; there is an activation energy for each elementary step, the first step is fast (low activation energy). Step 2 has the highest energy of activation (Ea2), this would be the slowest step in the reaction mechanism and becomes the rate-determining step for the overall reaction.
If the rate determining step is represented by the elementary reaction:
aA + bB → cC + dD
The rate of the reaction is dependent on both substances so the rate law expression would be: R = k[A]a[B]b
The order of the reaction in each reactant is the value of the exponent associated with that reactant. The overall order of the reaction is the sum of the exponents for each reactant. When reaction rate is proportional to the concentration to one reactant, the exponent is equal to one and the reaction is First-order in that reactant. Example [A] is doubled and the rate of the reaction doubles, the reaction is First-order in A. If the [B] is doubled and the rate increases by a factor of four, the exponent must be two and the reaction is Second-order in B. If the [B] is doubled and the reaction rate increases by a factor of eight, the exponent is three and the reaction is Third-order in B.
How does a chemist propose reaction mechanism? The only way is to determine the rate-determining step is through experimental evidence. The rate of this reaction:
C2H4Br2(l) + 3I-(aq) → C2H4(l) + 2Br-(aq) + I3-(aq)
doubles when either [C2H4Br] or [I-] is doubled. The exponents for the concentration terms must be one. The reaction is first order with respect to C2H4Br2 and first order with respect to I-. The rate-determining step according to experimental evidence must involve one of each of these particles. There are a number of possible mechanisms that agree with the observed data; consider the following.
C2H4Br2(l) + I-(aq) → C2H4Br-(aq) + IBr(aq) slow (rate determining step)
C2H4Br-(aq) → C2H4(l) +Br-(aq) fast
IBr(aq) + I-(aq) → Br-(aq) + I2(aq) fast
I2(aq) + I-(aq) → I3-(aq) fast
When added together the net reaction is:
C2H4Br2(l) + 3I-(aq) → C2H4(l) + 2Br-(aq) + I3-(aq)
Kinetics and Potential Energy diagrams
For the forward reaction:
1. What is the value of ∆H? What type of Chemical reaction? Point at Answer to see Answer
2. What is EA ? (EAF) Answer
3. What happens to the [AB] and [C] as the reaction progresses? Answer
4. What happens to the Rate of the reaction? Answer
Answers:
1. ∆H =+65kJ Endothermic
2. Ea = 99 kJ
3. The concentrations decrease.
4. The rate decreases.
In the reverse reaction:
1. What happens to the [AC] and [B] as the reaction progresses? Answer
2. What happens to the Rate of the reaction? Answer
3. What is the value of ∆H? What type of Chemical reaction? Answer
4. What is EA ? (EAR) Answer
Answers:
1. The concentrations increase.
2. The rate of the reverse reaction increases.
3. ∆H = - 65kJ; Exothermic reaction.
4. Ea = + 34kJ.
For the forward reaction:
1. What is the value of ∆H? What type of Chemical reaction?
2. What is EA ? (EAF)
3. What happens to the [O3] and [O] as the reaction progresses?
4. What happens to the Rate of the reaction?
For the reverse reaction:
1. What happens to the [O2] as the reaction progresses?
2. What happens to the Rate of the reaction?
3. What is the value of ∆H? What type of Chemical reaction?
4. What is EA ? (EAR)
Kinetics and Chemical Equilibrium
Eventually the rate of the forward reaction becomes equal to the rate of the reverse reaction.
The reactants are consumed as fast as they are being reformed.
The products are being produced as fast as they are reacting to reform the reactants.
The [reactants] becomes constant.
The [products] becomes constant.
When this occurs we have Chemical Equilibrium. (RForward = RReverse)
At equilibrium
The RF = RR.
RF = k[O3][O] and RR = k[O2]2.
Therefore:
k[O3][O] = k[O2]2 (because RF = RR)
k = [O2]2 . (k/k = K)
k [O3][O]
Keq = [O2]2 . (Equilibrium Constant)
[O3][O]
For a general reaction
aA + bB ↔ cC + dD
Keq = [C]c [D]d. (Products)
[A]a [B]b (Reactants)
If Keq is greater than 1 there are more products than reactants; if Keq is less than 1 there are more reactants than products when equilibrium is established.
Chemical Kinetics Questions
- State the collision theory for chemical reactions.
- Why don’t reactant particles just have to collide in order to produce products?
- Why is the energy requirement needed for a successful collision?
- Why is the geometry or orientation requirement needed for a successful collision?
- In spite of the extremely small percentage of successful collisions, why is it that chemical reactions are still observed to take place at a reasonable rate?
- Is it possible for a slow moving particle to have a successful collision with a very fast moving particle? Explain.
- Why does the rate of a chemical reaction slow with time?
- Sketch a graph of
a) concentration of reactant versus time.
b) concentration of product versus time
- Calculate the average rate of a chemical reaction if the concentration of reactant Y is 0.45 M at 3.0 min and 0.20 M after 8.0 min.
- Why is it referred to as average rate that we are calculating for a reaction time interval?
- Analyze the activation energy diagram shown below for the hypothetical reaction:
E + 2F ® G + H and answer the following questions:
i) What is the activation energy for the forward reaction? The reverse reaction?
ii) What is the value of DH for the forward reaction? The reverse reaction?
iii) What is the energy of the activated complex?
- The following hypothetical reaction has activation energy of 120 kJ and a DH of 80 kJ.
2A + B ® 2C + D
i) Draw and label a potential energy diagram for this reaction.
ii) Calculate the activation energy for the reverse reaction.
- Explain how concentration affects the rate of a reaction.
- Explain how temperature affects the rate of a reaction.
- Explain how amount of surface area affects the rate of a reaction.
- Explain how a catalyst affects the rate of a reaction.
- Explain why a match has little effect on lighting a lump of coal, yet a spark can cause coal dust to ignite with explosive force.
- Which of the factors that affect the rate of a reaction affect the collision frequency (number of collisions per second)?
- Why is the following reaction not likely to take place in one step:
4HBr(g) + O2(g) ® 2H2O(g) + 2Br2(g) - The above reaction in question number 19 is believed to take place in the following three steps:
i) HBr(g) + O2(g) ® HOOBr(g) Slow
ii) HOOBr(g) + HBr(g) ® 2HOBr(g) Fast
iii) 2(HOBr(g) + HBr(g) ® H2O(g) + Br2(g)) Fast