1. Chapter 7 Reaction Rates and Chemical Equilibrium

2. Reaction Rates The rate at which a chemical reaction occurs
In order for a reaction to occur, particles have to collide  collision theory of chemical reactions

3. Chemical Kinetics
Chemical kinetics: The study of the rates of chemical reactions.
Consider the reaction that takes place when chloromethane and sodium iodide are dissolved in acetone; the net ionic equation for this reaction is:
To determine the rate of this reaction, we measure the concentration of iodomethane at periodic time intervals, say every 10 minutes.

4. Chemical Kinetics
The rate of reaction is the increase in concentration of iodomethane divided by the time interval.
For example, the concentration might increase from 0 to 0.12 mol/L over a 30 minute time period.
The reaction rate over this period is:
This unit is read mole per liter per minute.

5. Chemical Kinetics
Figure 7.11 Changes in the concentration of B in the system A—> B with respect to time.

6. Rates of Reaction
The rates of chemical reactions are affected by the following factors:
Molecular collisions
Activation energy
Nature of the reactants
Concentration of the reactants
Temperature
Presence of a catalyst
On the following screens, we examine these factors one at a time.

7. Molecular Collisions
In order for two species, A and B (they may be molecules or ions), to react, they must collide.
It is possible to calculate how many collisions will take place between A and B in a given period of time.
Such calculations indicate that the rate at which A and B collide is far greater than the rate at which they react.
The conclusion is that most collisions do not result in a reaction.
A collision that results in a reaction is called an effective collision.
There are two main reasons why some collisions are effective and others are not; activation energy and the relative orientations of the colliding particles.

8. Molecular Collisions
Activation energy: The minimum energy required for a reaction to take place.
In most chemical reactions, one or more covalent bonds must be broken and energy is required for this to happen.
This energy comes from the collision between A and B.
If the collision energy is large, there is sufficient energy to break the necessary bonds, and reaction takes place.
If the collision energy is too small, no reaction occurs.

9. Molecular Collisions Orientation at the time of collision
The colliding particles must be properly oriented for bond breaking and bond making.
For example, to be an effective collision between H2O and HCl molecules, the oxygen of H2O must collide with the H of HCl so that the new O-H bond can form and the H-Cl bond can break.

10. Energy Diagrams
The reaction of H2 and N2 to form ammonia is exothermic:
In this reaction, six covalent bonds are broken and six new ones formed. In this calculation, the triple bond is counted as three bonds.
Breaking a bond requires energy, and forming a bond releases energy.
In this reaction, the energy released in making the six new bonds is greater than the energy required to break the six original bonds. The reaction is exothermic.

11. Energy Diagrams
Figure 7.4 Energy diagram for an exothermic reaction.

12. Energy Diagrams
Figure 7.5 Energy diagram for an endothermic reaction.

13. Energy Diagrams Transition state: A maximum on an energy diagram.
The transition state for the reaction between H2O and HCl probably looks like this. In the transition state, the new O-H bond is partially formed and the H-Cl bond is partially broken.

14. Factors Affecting Rate
Nature of reactants
In general, reactions between ions in aqueous solution are very fast (activation energies are very low).
In general, reaction between covalent compounds, whether in water or another solvent, are slower (their activation energies are higher).
Concentration
In most cases, the reaction rate increases when the concentration of either or both reactants increases.
For many reactions, there is a direct relationship between concentration and reaction rate. When concentration doubles the rate doubles.

15. Factors Affecting Rate
Temperature
In virtually all reactions, rate increases as temperature increases.
An approximate rule for many reactions is that for a 10°C increase in temperature, the reaction rate doubles.
When temperature increases, molecules move faster (they have more kinetic energy), which means that they collide more frequently. More frequent collisions mean higher reaction rates.
Not only do molecules move faster at higher temperatures, but the fraction of molecules with energy equal to or greater than the activation energy also increases.

16. Factors Affecting Rate
Figure 7.8 The distribution of kinetic energies (molecular velocities) at two temperatures.

17. Factors Affecting Rate
Catalyst: A substance that increases the rate of a chemical reaction without itself being used up.

18. Factor Affecting Rate
One way a catalyst can affect the rate of a reaction is to provide a surface on which reactants can meet.
The reaction of ethylene with hydrogen is an exothermic reaction.
If these two reagents are mixed, there is no visible reaction even over long periods of time.
When they are mixed and shaken with a finely divided transition metal catalyst, such as Pd, Pt, or Ni, reaction takes place readily at room temperature.

19. Reversible Reactions
Reversible reaction: A reaction that can be made to go in either direction.
If we mix CO and H2O in the gas phase at high temperatures, CO2 and H2 are formed:
We can also make the reaction take place the other way by mixing CO2 and H2.
The reaction is reversible, and we can discuss both a forward reaction and a reverse reaction.

20. Reversible Reactions
Equilibrium: A dynamic state in which the rate of the forward reaction is equal to the rate of the reverse reaction.
At equilibrium there is no change in concentration of either reactants or products.
Reaction, however, is still taking place. Reactants are still being converted to products and products to reactants, but the rates of the two reactions are equal.

21. Equilibrium Constants
Equilibrium constant, Keq: The product of the concentrations of products of a chemical equilibrium divided by the concentrations of reactants, each raised to the power equal to its coefficient in the balanced chemical equation.
A quantitative way of viewing chemical equilibrium
Value of Keq allows us to determine whether equilibrium favors the reactants or products
For the general reaction:
The equilibrium constant expression is:

22. Equilibrium Constants
Problem: Write the equilibrium constant expression for this reversible reaction:
Solution: The equilibrium constant expression is:
Note that no exponents are shown in this equilibrium constant expression; by convention the exponent “1” is understood but not written.

23. Equilibrium constants
Write the equilibrium constant for the reaction
O2(g) + 4ClO2(g) ↔ 2Cl2O5(g)

24. Equilibrium Constants
Problem: When H2 and I2 react at 427°C, the following equilibrium is reached:
The equilibrium concentrations are [I2] = 0.42 mol/L, [H2] = mol/L, and [HI] = 0.76 mol/L. Using these values, calculate the value of K.
Solution:
This K has no units because molarities cancel.

25. Equilibrium and Rates
There is no relationship between a reaction rate and the value of K.
Reaction rate depends on the activation energy of the forward and reverse reactions. These rates determine how fast equilibrium is reached but not its position.
It is possible to have a large K and a slow rate at which equilibrium is reached.
It is also possible to have a small K and a fast rate at which equilibrium is reached.
It is also possible to have any combination of K and rate in between these two extremes.

26. Le Chatelier’s Principle
Le Chatelier’s Principle: When a stress is applied to a chemical system at equilibrium, the position of the equilibrium shifts in the direction to relieve the applied stress.
We look at three types of stress that can be applied to a chemical equilibrium:
addition of a reaction component
removal of a reaction component
change in temperature

27. Le Chatelier’s Principle
Addition of a reaction component
Suppose this reaction reaches equilibrium:
Suppose we now disturb the equilibrium by adding some acetic acid.
The rate of the forward reaction increases and the concentrations of ethyl acetate and water increase.
As this happens, the rate of the reverse reaction also increases.
In time, the two rates will again become equal and a new equilibrium will be established.

28. Le Chatelier’s Principle
At the new equilibrium, the concentrations of reactants and products again become constant, but with different values than before the addition of acetic acid.
The concentrations of ethyl acetate and water are now higher, and the concentration of ethanol is lower.
The concentration of acetic acid is also higher, but not as high as it was immediately after we added the extra amount.
The system has relieved the stress by increasing the components on the right side of the equilibrium.
We say that the system has shifted to minimize the stress.

29. Le Chatelier’s Principle
Removal of a reaction component
Removal of a component shifts the position of equilibrium to the side that produces more of the component that has been removed.
Suppose we remove some ethyl acetate from this equilibrium:
If ethyl acetate is removed, the position of equilibrium shifts to the right to produce more ethyl acetate and restore equilibrium.
The effect of removing a component is the opposite of adding one.

30. Le Chatelier’s Principle
Problem: When acid rain (H2SO4(aq))attacks marble (calcium carbonate), the following equilibrium can be written:
How does the fact that CO2 is a gas influence the equilibrium?
Solution: CO2 gas diffuses from the reaction site, and is removed from the equilibrium mixture. The equilibrium shifts to the right and the marble continues to erode.

31. Le Chatelier’s Principle
Change in temperature
The effect of a change in temperature on an equilibrium depends on whether the forward reaction is exothermic or endothermic.
Consider this exothermic reaction:
We can consider heat as a product of the reaction.
Adding heat (increasing the temperature) pushes the equilibrium to the left.
Removing heat (decreasing the temperature) pushes the equilibrium to the right.

32. Le Chatelier’s Principle
Summary of the effects of change of temperature on a system in equilibrium

33. Chapter 7 Equilibrium and Reaction Rates
End
Chapter 7