1. Unit 1: Kinetics First Unit Quiz I.1-1.4 (pages 1-11 Hebden)
Definition
Definition of Reaction rates
Units
measuring reaction rates
Factors effecting reaction rates *
Calculating Reaction rates

2. 1.Definition of Kinetics
Kinetics is the study of reaction rates and factors that effect reaction rates

3. 2. Reaction Rates

4. 3. Units
Reaction Rates can be expressed in many units, however a quantity unit must be in the numerator, and a time unit must be in the denominator
Examples
G/s (grams per second)
L/Hr. (liters per hour)
Mol/min. ( moles per minute)
Pa/s (pascals per second)
Candles/Hr.
C/min (degrees Celsius per min

5. 4. Measuring Reaction Rates
To measure a reaction rate, you need to be able to identify a quantifiable property that will change as the reaction takes place.
What quantifiable property is chosen depends or the chemical system and whether the experiment is preformed as an open system or closed system experiment.

6. 4. Measuring reaction rates
Examples of Quantifiable properties
Mass (change in mass)
volume (change in volume)
Temperature (change in temperature)
Pressure (change in pressure)
Color (change in color)
pH (change in H+ concentration)

7. Measuring Reaction Rates
Problem: identify 3 methods for measuring the rate of this reaction

8. Calculating Reaction Rates
Reaction rates are always expressed as positive numbers.

9. Factors Effecting Reaction Rates (Quiz 2)
Temperature: Generally, the greater the temperature the faster a reaction occurs
Rule of 10 for many reactions, if you increase the temperature by 10 degrees C, you get a doubling of reaction rates
Concentration of reactants: The more concentrated reactants are the faster the reaction occurs.

10. Factors Effecting Reaction Rates
Pressure: When Gases react, the more pressure they are under, the faster they will react.
Nature of reactants: Some chemical bonds are stronger than others, chemicals with stronger bonds tend to react slower than ones with weak bonds.
Mixing and Surface area: For a reaction to occur, reactants must come in contact, reactants that have a greater surface area have more points of contact and can react faster.

11. Factors Effecting Reaction Rates
Catalysts: a catalyst is something (usually a chemical) that is added to the reaction which increases the rate of the reaction. **The catalyst is not consumed by the reaction.
Inhibitor: is something (usually a chemical) that is added to the reaction which slows down the rate of reaction. ** The inhibitor is not consumed by the reaction

12. Factors Effecting Reaction Rates
Nature of Reactants
Type of reactant (fastest to slowest RR)
Ions
Gasses
Liquids
Solids
For Heterogeneous reactions, the reaction can only occur at the interface. I.e., at the surface where the two phases meet. This is why surface area is an important factor affecting rates in a heterogeneous reaction.

13. Collision Theory and Potential Energy Diagrams Quiz topics
Kinetic Molecular theory review
Collision Theory
Requirements for successful reactions
Thermo dynamic definitions
The potential energy diagram
Kinetic energy distributions

14. Kinetic Molecular Theory
All matter is made of small particles
The amount of space between particles depends on what State the matter is in
The particles that make up matter are always moving
Particles of matter are attracted to each other, the strength of attraction depends on the type of particle.

15. Collision theory
For a reaction to occur between molecules they must collide
The collision must have sufficient energy
The collision must have the correct orientation

16. Laws of Thermodynamics
Zeroth laws: Conservation laws
Matter is neither created nor destroyed
Energy is neither created nor destroyed
Charge is neither created nor destroyed
First Law (law of laziness) All systems tend towards there lowest energy state
Second Law (law of messiness)All systems tend towards a state of maximum disorder
Third Law Heat flows from a hot object to cold, never from a cold object to a hot object.

17. Thermodynamics Basic Definitions
Kinetic energy is the energy a particle has due to its motion
Potential energy is the energy stored in the chemical bonds of the molecule and the energy due to its position in space.
Enthalpy (H): The total kinetic and potential energy in an OPEN SYSTEM
H is the change in enthalpy and represents the amount of heat energy absorbed or given off by the system

18. Reaction Mechanisms (key Concepts)
Definition
Drawing multi-step reactions on a PE Diagram
Identifying and distinguishing between: activated complex, intermediate products and catalysts in multi-step reactions
Writing net reactions
Concept of rate determining step

19. Definition
Reaction Mechanisms: are the actual steps by which a reaction occurs.
Some reactions proceed in only 1 step, however many reactions involve multi-step processes, especially reactions involving more than 2 reactants or reactions that use a catalyst.
For reactions that involve more than 2 reagents, the reaction process will have many steps. This is because the probability of mare than 2 reactants coming together all with the proper orientation and kinetic energy is very low.

20. Catalysts (practical applications)
All catalysts increase the rate of reaction with out being consumed. They do this by providing an alternative route for the reaction mechanism that lowers the energy of activation for
Homogenous Catalysts are in the same phase as reactants
Hetrogenous catalysts are in a different phase than the reactants and and offer a path of lower energy by providing a surface on which reactants can be adsorbed

21. Industrial Applications of Catalysts
Production of H2SO4 (heterogeneous , V2O5)
2S + 3O2 +2H2O => 2H2SO4
Haber Process (heterogeneous , FeO, RhO)
N2 + 3H2 => 2NH3
Catalytic Converters (heterogeneous , Pt, Rh, Ir)
2NO=>N2 +O2
CO +CxH2x+2 + O2 =>CO2 +H2O
Starch=> sugar (homogenous maltase)