Chem 2 - Chapter 13: Kinetics

Front 1

Chemical Kinetics

Back 1

The study of reaction rates, and tells us the change in concentrations of reactants or products over a period of time

Front 2

Simplest form of a chemical Reaction:

Back 2

Reactants break down to form products

Front 3

Rate

Back 3

a change in some quantity per unit time; average

Front 4

Why Chemists Study Rates:

Back 4

1. To control chemical reactions
2. to give information about mechanisms

Front 5

Factors that CAN effect Rate:

Back 5

1. Concentration:
2. Temperature
3. Particle Size (Surface Area):
4. Catalyst:

Front 6

Concentration:

Back 6

Molecules must collide to react.

Increasing the number of molecules in a container increase their collisions and therefore causes the rate to increase

Front 7

Temperature:

Back 7

Increasing the temperature increases the reaction rate by increasing the number and energy of collisions.

Front 8

Particle Size (surface area):

Back 8

the frequency of collisions increases with increased surface area

Linear (acyclic) structures have more surface area than cyclic structures

Front 9

Cyclic

Back 9

in a ring

less surface area

less collisions

Front 10

Catalyst:

Back 10

A catalyst increases the rate of a reaction by lowering the activation energy

not used up during a reaction (regenerates)

Front 11

Average Rate:

Back 11

The rate of a chemical reaction is the change in some variable over a period of time

Avg. Rate = Δ [A]/ Δtime

A = reactant or product

[ ] = concentration or Molarity

Front 12

Initial Rate

Back 12

Instantaneous rate

The slope of the tangent line

Initial rate is at t = 0

Front 13

Ways to Measure Concentration

Back 13

1. Evolution of gas (change in P)
2. Color (spectroscopic absorbance )
---Appearance/Disappearance of color or light
3. Electrical Conductivity

Front 14

Continuous Monitoring:

Back 14

1. Spectrophotometry
2. Total Pressure

Front 15

Spectrophotometry:

Back 15

measuring the amount of light of a particular wavelength absorbed by one component over time
-- the component absorbs its complementary color

Front 16

Total Pressure:

Back 16

the total pressure of a gas mixture is stoichiometrically related to partial pressures of the gases in the reaction

Front 17

Gas Chromatography:

Back 17

Can measure the concentrations of various components in a mixture
--For samples that have volatile components
-- Separates mixture by adherence to a surface

Front 18

Drawing off Periodic Aliquots:

Back 18

Drawing off periodic aliquots from the mixture and doing quantitative analysis
-- Titration for one of the components
-- Gravimetric analysis

Front 19

Rate Laws:

Back 19

must be experimentally derived (not necessarily related to stoichiometry) unless it is a slow step

ignore the reverse reaction and depends only on REACTANT concentrations and the rate constant

Rate = -Δ[A]/Δt = k[A]^x*[B]^y

[A] and [B] are initial concentrations
k = rate constant
x and y = reaction orders

Front 20

At equilibrium:

Back 20

the rate of formation for reactants and products are equal

Front 21

Zero Order

Back 21

n = 0

Double the concentration --> rate stays the same

Δconcentration is independent of the rate
Rate is independent of the Δconcentration

Rate = K

Front 22

First Order

Back 22

n = 1

The rate is directly proportional to the concentration of A

Rate = K[A]

Front 23

Second Order

Back 23

n = 2

The rate is directly proportional to the square of the concentration of A

Rate = k[A]^2

Front 24

Units of K:

Back 24

M/time

K = M^(-n+1) * time^-1

n = Overall order

Front 25

Steps to Solve the Rate Law:

Back 25

1. Examine Experimental Data
2. Select 2 experiments which vary only one reactant concentration
3. Solve the ration of the two experiments to find exponents
4. Return to step 2 until all reactants have been solved
5. Solve for K

Front 26

Problems in Measuring Rates:

Back 26

Rate law doesn't include time

1. Rates chance
2. Need instantaneous rates
3. Instantaneous rate at t=0 is defined as the initial rate

Front 27

Integrated Rate Law:

Back 27

helps to answer:

How long will it take x moles per liters of A to be consumed?

factors in time

Front 28

1st Order Equations:

(for Integrated Rate Law and half-life)

Back 28

ln [A]t = -kt + ln[A]0
y = mx + b

ln[A] vs. time = linear

t1/2 = ln2/k = 0.693/k

Front 29

Zero Order Equation:

(for Integrated Rate Law:)

Back 29

[A] = -kt + [A]0
y = mx + b

[A]0 = initial

t1/2 = [A]0/2k

Front 30

Second Order Equation:

(for Integrated Rate Law)

Back 30

1/[A] = kt + 1/[A]0
y =mx + b

t1/2 = 1/k[A]0

Front 31

Collision Theory:

Back 31

Reactions occur because reactant molecules collide

Collisions can cause the bonds of reactant molecules to break

This allows the atoms to be arranged in new combinations therefore forming products

Front 32

Simplified Collision Theory:

Back 32

The rate can be effected by factors that:

1. Increase the number of effective collisions per unit time (Results in product)

2. Increase the force of collisions

Front 33

Properties of Collisions:

Back 33

1. Molecules must collide
2. Must collide with sufficient energy to pass the activation energy
3. Must collide in the correct orientation

Front 34

Reasons Why Collisions are NOT Successful:

Back 34

1. Geometry of Collisions
--- Not the right orientation

2. Threshold Energies
--- Not enough energy

Front 35

Temperature and The Reaction Rate:

Back 35

Temperature affects rate by affecting the rate constant (k)

As Temperature Increases, K increases.

Front 36

Arrhenius Equation:

Back 36

k = Ae^(-Ea/RT)

A = Frequency Factor
-Ea = Activation Energy
R = 8.314 J/mol*K
T = Temperature in Kelvin

Front 37

The slower the rate of the reaction, the _________ the Activation Energy. And the higher the temperature, the ______ the value of K (rate constant).

Back 37

larger, larger

Front 38

Two Point Form of Arrhenius Equation:

Back 38

ln K2/K1 = Ea/R [1/T1 - 1/T2]

Front 39

Reaction Mechanism:

Back 39

a sequence of single reaction steps (elementary steps) that add up to the overall chemical reaction.

The overall reaction gives the stoichiometric reaction equation.

Front 40

Unimolecular

Back 40

1 mole of reactant

rate = k[A]

Front 41

Bimolecular

Back 41

2 moles of reactant

rate = k[A]^2

or

rate = k[A][B]

Front 42

Tetramolecular

Back 42

4 moles of reactant

Front 43

Intermediates:

Back 43

Begin as products then are found as reactants

Cancel out

A reaction intermediate is formed in one step and consumed in the next
Not shown in the net equation for the complete reaction

Front 44

Catalysts:

Back 44

Begin as one of the first reactants and then as one of the last products in the last reaction

Also cancel out

Speed up the reaction (doesn't cause reaction) by lowering the Activation Energy (Ea)

Makes an alternative mechanism possible

Is NOT part of the stoichiometry of the reaction (b/c its not a reactant or a product) but it is part of the rate law

Front 45

Slow Step

Back 45

the rate-determining step

Front 46

Energy Diagram

Back 46

a way of illustrating how the energy levels in a chemical reaction change from reactants to products.

The number of steps involved are shown by "hills"

Front 47

Exothermic

Back 47

exergonic

products are lower than your reactants

energy (heat) is released

-ΔH

Front 48

Endothermic

Back 48

endogonic

products are higher up than reactants

energy (heat) is absorbed