1. Kinetics

2. Kinetics
Kinetics - rates of chemical reactions and the mechanisms by which they occur
Rate of a chemical reaction - change in the concentration of products and reactants in a given time

3. Collision Theory
Molecules must have effective collisions in order to react
they must have the correct amount of energy
correct orientation

4. Collision Theory Not all collisions are successful
How to increase the rate of a reaction?
Increase the number of collisions
Increase the effectiveness of the collisions

5. Factors that Affect Reaction Rates
Anything that increases the number of collisions will increase the reaction rate
Nature of the Reactants
Concentration
Temperature
Catalyst
Surface Area
Pressure (only for gases)

6. Nature of the Reactants
Reactions involving ions (or ionic compounds) are generally faster than those involving covalently bonded substances
Covalently bonded substances have more bonds that must be broken before the reaction can occur

7. Nature of the Reactants
Energy required to break bonds is proportional to the stability of the bond
More stable bonds (stronger bond) require more energy to break, slower reaction, less reactive
Weaker bonds are broken with less energy, faster reaction, more reactive

8. Concentration
When the concentration of one or more of the reactants is increased, the reaction proceeds faster
As concentration increases, there are more particles, more likely to collide

9. Pressure of a Gas
When the pressure of a gas increases, volume decreases
Making the particles closer together
Therefore there will be more collisions
P = 1.0 atm
P = 2.0 atm

10. Temperature An increase in temperature increases the rate
Higher temperatures cause particles to move faster and have more kinetic energy
Therefore, more collisions and more effective collisions, due to the increased kinetic energy

11. Catalyst
Addition of a catalyst lowers the activation energy, speeding up the reaction
Catalysts are NOT used up in a reaction

12. Surface Area More surface area, faster reaction
With more area exposed, there will be more collisions
Example: Both samples represent 2.5g of Mg, which would react faster with 25mL of 1.0M HCl?
Sample A
Sample B

13. Reaction Mechanism
Series of steps that leads from reactants to products
Describes the order in which bonds break and atoms rearrange during a reaction
Composed of a set of simple steps, which will sum to a net reaction

14. Intermediates
Substance that is formed and consumed during the course of a reaction
Intermediates do not appear in the net reaction equation

15. Reaction Mechanism Example: Step 1: NO2(g) + NO2(g)  NO3(g) + NO(g)
Step 2: NO3(g) + CO(g)  NO2(g) +CO2(g)
Net Rxn:
Intermediates:

16. Reaction Mechanism Step 1: S + O2  SO2 Step 2: 2SO2 + O2  2SO3
Step 3: SO3 + H2O  H2SO4
Step 4: SO3 + H2SO4  H2S2O7
Net Rxn:
Intermediates:

17. Entropy (S) Measure of randomness/disorder
Greater disorder, higher entropy
Spontaneous reactions tend to go towards higher entropy (more disorder)
Examples:
Solid  Liquid  Gas
(great order  great disorder)
Messy bedroom  high entropy
More moles  more disorder higher entropy
(only look at the moles if the phases are all the same)

18. Entropy Examples
Given the following reactions, indicate if entropy increases, decreases, or remains the same
1. H2O(l) H2O(g)
2. 3H2(g) + N2(g) 2NH3(g)

19. Potential Energy Diagrams
PE of activated complex
Reaction coordinate
700kJ
Activation Energy
Potential Energy
500kJ H
PE of reactants
300kJ
PE of products

20. Activation Energy
Minimum energy required to initiate a chemical reaction (energy to break bonds)
Energy required to form an activated complex
Equal to the difference between the Potential Energy of activated complex and potential energy of the reactants
All reactions require Activation Energy
The Larger the Activation Energy, the slower the reaction

21. Activated Complex Highest point on Potential Energy curve
It represents a transition state between the products and reactants, NOT an intermediate product
If no activated complex is formed – NO REACTION

22. Heat of Reaction (Enthalpy)
Difference between the potential energy of the products and the potential energy of the reactants
Energy given off or absorbed by the reaction
H = Hp – Hr
Found on Reference Table I

24. 2H2(g) + O2(g)  2H2O(l) H = -571.6 kJ
Heat of Reaction
If H is negative
The reaction is EXOTHERMIC
The Potential Energy of the products is less then the Potential Energy of the reactants
Energy is leaving (exiting) the system
Ex.
2H2(g) + O2(g)  2H2O(l) H = kJ
2H2(g) + O2(g)  2H2O(l) kJ

25. Heat of Reaction If H is positive The reaction is ENDOTHERMIC
The Potential Energy of the products is more then the Potential Energy of the reactants
Energy is entering the system (gain of heat)
Example
N2(g) + 2O2(g)  2NO2(g) H = kJ
N2(g) + 2O2(g) +66.4kJ  2NO2(g)

26. Does this graph show an exothermic or endothermic reaction?

27. Classify the following examples as Endothermic or Exothermic?
CH4 + 2O2 → CO2 + H2O KJ
2C + 2H KJ → C2H4
4Al + 3O2 → 2Al2O KJ
H+(aq) + OH- (aq) → H2O +55.8KJ
KNO KJ → K+(aq) + NO3-(aq)

28. Table I Examples Which reaction releases the most energy?
Which reaction absorbs the most energy?
What is the heat of reaction for the synthesis of 2 moles of Al2O3?
What is the heat of reaction for the synthesis of 1 mole of Al2O3?
What is the heat of reaction for the decomposition of 1 mole of NH3(g)?
As sodium hydroxide ionizes in water, what happens to the temperature of the water?

29. Spontaneous Reactions
A reaction is most likely to occur when there is a:
Change to a condition of less energy
Exothermic (lower AE), Negative H
Change to a condition of greater randomness
Greater entropy, Positive S