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Unit 13. Reaction Rates. Unit 13 Goals. Understand collision theory Understand how concentration, surface area & temperature, effect reaction rate. Understand potential energy diagrams Predict reaction rate change when concentration, temperature or surface area ∆’s

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Unit 13

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## Unit 13

Reaction Rates

### Unit 13 Goals

• Understand collision theory

• Understand how concentration, surface area & temperature, effect reaction rate.

• Understand potential energy diagrams

• Predict reaction rate change when concentration, temperature or surface area ∆’s

• Construct and interpret PE diagrams

### Thinker:

• Make a list of RATES that you have heard about in your daily lives.

• Rate of growth

• Rate of pay

• Rate of decay

• Rate of aging

• Rate of flow

• Rate of power usage

• Rate of speed

### Rates of Chemical Change

• Rate is how fast something changes with time

• Sprinters

• Human 28.9 mph

• Cheetah 62 mph

• Quarter horse 55 mph

• Reaction Rate – rate at which reactants are changed to products

• Chemical Kinetics – study of Rx rates

http://en.wikipedia.org/wiki/Footspeed

Karl Wenzel

German Scientist 1777

First to study Rx rates

Claude-Louis Berthollet

French – 1803

Published first book on Rx rates & concentration

### Chemical Kinetics

http://en.wikipedia.org/wiki/Claude-Louis_Berthollet

### Rate = Change over time

• CH3OCH3(g) CH4(g) + CO(g) + H2(g)

• At 500 C, dimethyl ether decomposes

• As Rx continues:

• Conc. of dimethyl ether will decrease

• Time will continue to change

• Therefore:

• Note: all reactants are -

### CH3OCH3(g) CH4(g) + CO(g) + H2(g)

• We can talk about the rates of reactants OR products:

• Consider this Rx:

• 2N2O5(s) 4NO2(g) + O2(g)

• How do coefficients affect reactant/product?

• Divide by each coefficient:

• Think about how many per time unit

### Measuring Reactions

• Need to keep track of at least 1 concentration:

• Color change

• Pressure change

• Temperature change

• Rates is measured as is indicated by the rate you write.

Calculate the rate of reaction between 0 and 20.0 s:

0.0000585 M/s

Calculate the rate of reaction between 20.0 and 40.0 s

0.0000528 M/s

### Will Reaction Rate always be the same?

• 2N2O5(s) 4NO2(g) + O2(g)

• No!

• Reactant is constantly decreasing

• Product is constantly increasing

slope of graph = rate of Rx

### Factors Affecting Rate

• Concentration

• Pressure

• Temperature

• Surface Area

### Concentration

• Almost always:

• Increase in concentration of reactants yields in increase of reaction rate

NO2(g) + CO(g) NO(g) + CO2(g)

• If [NO2] is doubled

• More collisions with CO will occur

• More collisions means more successful collisions

• More successful collisions means more product

• The inverse is also true

### Non-collision reactions

• Decomposition does not require collisions

(CH2)3(g) CH2=CH-CH3(g)

• The more cyclopropane you have:

• The more is available to decompose

• Thus a faster reaction rate

### Pressure

• In liquid or solid, has almost no affect

• In gasses:

• Doubling pressure essentially doubles the concentration.

•  by increasing pressure, we increase reaction rates in gasses.

• Inverse is true

### Temperature

• In General:

• Increase in temperature will increase reaction rate.

• Often increase of 10% temp yields in 2 times the reaction rate!

• Our bodies work best around 37 C

• Small changes cause great distress!

• Snakes, lizards & other reptiles

### Surface Area

• Increased surface area yields an increase in reaction rate.

• Wood & fires

• Grain Elevator

### Questions?

• P. 585 Section Review:

• 1-3, 6, 8, 11 – 13

### Rate Law

• Describes how reactant concentration affect the reaction rate.

• Provides the ability to form a reaction mechanism.

• Model of how the reaction occurs

• Discusses # of steps necessary

### Determining a General Rate Law Equation

• When there is a single reactant:

• Rate of reaction is proportional to the concentration raised to some power (n)

• The power is the order of the Rx.

Rate = k[reactant]n

• k is the rate constant

• Proportionality that varies with temperature

• Order is usually a whole number (1, 2)

• Can be fractions

• Occasionally is 0

• Reaction rate is independent of concentration

### Determining Order (example)

• Experiments were preformed to measure the initial rate of the reaction 2HI(g)  H2(g) + I2(g). Conditions were identical except that the HI concentrations were varied.

• rate = k[HI]n

• n=?

• Find ratio of the reactant concentrations b/t experiments.

• Find ratio of the reaction rates.

• Plug into rate = k[HI]n

• Verify the Results (compare 3 and 1)

### Other questions

• Given the previous reaction:

• What will happen if [HI] is increased from 4 M to 8 M?

• What will happen if [HI] is decreased from 6 M to 2 M?

• What will happen if [HI] is quadrupled?

• Homework:

• Worksheet

### Rate Laws for Several Reactants

• When there are several reactants the rate law applies to each.

•  an order applies to each reactant

NO(g) + O3(g) NO2(g) + O2(g)

• rate = k[NO3]n1[O3]n2

• It turns out:

• n1 = n2 = 1

• Since each order is equal, it suggests that this reaction has a 1 step mechanism.

• If they weren’t 1, then the mechanism could be longer than 1 step.

### Rate Determining Step

• Despite being written as 1 step, many Rx, aren’t quite that simple.

• The mechanism is more complex.

• Example:

2Br-(aq) + H2O2(aq) + 2H3O+(aq) Br2(aq) + 4H2O(l)

• Has 4 separate steps (each reactants order 1)

• Br-(aq) + H3O+(aq) HBr(aq) + H2O(l)

• HBr(aq) + H2O2(aq) HOBr(aq) + H2O(l)

• Br-(aq) + HOBr(aq)  Br2(aq) + OH-(aq)

• OH-(aq) + H3O+(aq)  2H2O(l)

• Each step produces intermediates.

• Species that are produced, but consumed

• If one of the steps is slower, it is known as the rate determining step

• In this case it is step 2

### Collision Theory:

• Reactions occur because:

• Substances (atoms, ions, molecules) must collide

• As particles get closer, they repel, so they must have enough KE to actually collide

• Collisions must occur in correct orientation

• Collisions must have sufficient energy to form the activated complex

• This energy is known as activation energy (Ea)

A + B  C + D

Imagine trying to bowl uphill

This is what molecules need to do

Activation energyis required to formactivated complexat top of curve.

# peaks = # of steps

### Collision Theory

• Reacting substances (atoms, ions, or molecules) must collide

• Collisions must occur with the correct orientation

• Collisions must have sufficient energy to form the activated complex.

• Activated complex is a temporary unstable arrangement of atoms that may form products or break apart to re-form the reactants

### Successful Collision

Catalyst – chemical in a reaction that lowers the activation energy, but never is used.

Save energy

Usually costly

Biological catalysts are called enzymes

### Catalysts & Potential Energy

http://en.wikipedia.org/wiki/Image:Activation_energy.svg

### Homework:

• PE Diagram Worksheet