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unit 8 Flashcards

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1
Q

the rate of a chemical reaction measures

A

how fast the reaction occurs

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2
Q

what happens when a chemical reaction happens at a fast rate?

A

large fraction of molecules react to form products in a period of time

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3
Q

what happens when a chemical reaction happens at a slow rate?

A

relatively small fraction of molecules react to form products per unit time

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4
Q

rate definition

A

change in a quantity (concentration of reactant/product) per unit time

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5
Q

rate (reactant) equation

A
  • (change in concentration) / (change in time)
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6
Q

reactant concentration _ over time

A

decreases

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7
Q

product concentration _ over time

A

increases

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8
Q

rate (product) equation

A

(change in concentration) / (change in time)

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9
Q

what happens to the rate of reaction when a coefficient is before the element/compound in a balanced equation?

A

multiply the rate by the positive reciprocal of the coefficient

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10
Q

reactant concentration _ with time as reactants are consumed

A

decreases

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11
Q

product concentration _ with time as products are formed

A

increases

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12
Q

4 factos that affect the rate of the reaction

A

physical state
reactant concentration
reaction temp
presence of a catalyst

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13
Q

the more readily reactant molecules collide with each other, the more _ they react

A

rapidly

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14
Q

reactant concentration and reaction rate proportion

A

directly proportional

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15
Q

temp and reaction rate proportion

A

direct (generally)

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16
Q

catalyst definition

A

agents that increase reaction rates without themselves being used up

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17
Q

2 methods of measuring reaction concentrations

A

spectroscopy
monitoring the pressure of gases

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18
Q

how to measure the reactant concentrations from spectroscopy:

A

measure the change in absorbance of light
calculating concentration using beer’s law

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19
Q

rate law definition

A

relationship between rate of reaction and concentration of the reactant

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20
Q

rate law formula

A

rate = k[A]^n

k: rate constant
n: reaction order

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21
Q

zero order reaction

A

n = 0
-molarity of reactant decreases linearly
-reaction does not slow down as [A] decreases
-rate order is the same at any [A]
-amount of reactant available for reaction is unaffected by overall quantity of reactant

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22
Q

first order

A

n = 1
-rate slows as reaction proceeds
-rate is slower as reaction proceeds
-rate is directly proportional to concentration

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23
Q

second order

A

n = 2
-rate is proportional to the square of [reactant]
-rate is more sensitive to [reactant]
-rate slows faster than first order reaction

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24
Q

the order of a reaction can only be determined by _

A

experiment

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25
Q

method of initial rates

A

initial rate is measured by running the reaction several different times with different initial [reactant] to determine the effect of [reactant] on rate

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26
Q

initial rate definition

A

a short period of time at the beginning of a reaction

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27
Q

if reaction is 0 order in A

A

-initial rate is independent of [reactant]
-rate is the same for all measured initial concentrations
-rate = k
-units: M x s^-1

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28
Q

if the reaction is 1st order in A

A

-when [A] doubles, rate doubles
-inital rate is directly proportional to initial concentration
-rate = k[A]^1
-units: s^-1

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29
Q

if the reaction is second order in A

A

-when [A] doubles, rate quadruples
-relationship between concentration and rate is quadratic
-rate=k[A]^2
-units: M^-1 x s^-1

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30
Q

overall order

A

sum of exponents m and n

m = A’s reaction order
n = B’s reaction order

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31
Q

differential rate law (first order) equation

A

k[A] = -(Delta [A]) / (delta t)

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32
Q

integrated rate law (first order) equation

A

ln [A]t = -kt + ln [A]0

or

ln ([A]t) / ([A]0) = -kt

[A]t is [A] at any time
k is rate constant
[A]0 is the initial [A]

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33
Q

integrated rate law graph

A

straight line
y=mx+b
slope = -k
time on x axis
ln[A]0 on y axis
slope is not negative, but rate constant is always positive

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34
Q

differential rate law (second order) equation

A

k[A]^2 = -(Delta [A]) / (delta t)

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35
Q

integrated rate law (second order) equation

A

(1) / ([A]t) = kt + (1 / [A]0)

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36
Q

integrated rate law (zero order) equation

A

k = -(Delta [A]) / (delta t)

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37
Q

differential rate law (zero order) equation

A

[A]t = -kt + [A]0

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38
Q

half life definition

A

time required for a reactant to fall 1/2 of its original value

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39
Q

first order half life equation

A

0.693 / k

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40
Q

for first order reactions, t1/2 is

A

independent of initial concentration

-even though [reactant] changes as reaction proceeds, half life is constant

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41
Q

second order half life equation

A

1 / (k[A]0)

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42
Q

zero order half life equation

A

[A]0 / 2k

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43
Q

half lives of zero and second order reactions depend on

A

initial concentration

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44
Q

rate law zero order

A

k[A]⁰

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45
Q

rate law first order

A

k[A]¹

46
Q

rate law second order

A

k[A]²

47
Q

zero order unit of k

A

M x s⁻¹

48
Q

first order unit of k

A

s⁻¹

49
Q

second order unit of k

A

M⁻¹ s⁻¹

50
Q

the potential energy diagram represents the ___ ___ that occur during a chemical reaction

A

energy changes

51
Q

y axis in a potential energy diagram represents

A

potential energy in kJ/mol

52
Q

x axis in a potential energy diagram represents

A

progression of the reaction from reactants activated complex to products

53
Q

chemical energy is sometimes referred to as

A

bond energy

54
Q

chemical energy/bond energy is a form of _ energy

A

potential

55
Q

bond breaking requires

A

energy input

-breaking electrostatic forces

56
Q

bond forming results in

A

energy release

-results in more stable situation

57
Q

endothermic reaction

A

more energy is absorbed as bonds break than is released when bonds formed

positive delta H

58
Q

exothermic reaction

A

more energy is released as bonds form than absorbed as they break

negative delta H

59
Q

activation energy (Ea) definition

A

difference between potential energy of the activated complex and the total potential energy of the separated reactant molecules

represents the amount of energy the reactant molecules must gain to form an activated complex

60
Q

activation energy formula

A

Ea = PE(activated complex) - PE(reactants)

61
Q

enthalpy change (delta H) definition

A

difference between the total potential energy of the products and the total potential energy of the reactants

62
Q

enthalpy change formula

A

delta H = PE(products) - PE(reactants)

63
Q

collision theory

A

reaction rate is equal to frequency of effective collisions between reactants

64
Q

in order for a reaction to occur, molecules must

A

-collide with sufficient energy
-properly oriented so products can form

65
Q

the fraction of collisions with minimum energy (activation energy) increases with

A

rising temp

-avg kinetic energy of molecules increases with temp

66
Q

are reactions more likely to occur with molecules or single atoms?

A

single atoms

-greater chance of forming a product due to its orientation
-spherical

67
Q

generally, what happens when gases collide at normal temp?

A

-most collisions do not have enough energy to overcome activation barrier
-atoms bounce off each other
-the molecules that do overcome the activation barrier do not have the orientation for the reaction to occur

-electrons from one atom are attracted to the nucleus of the other atom
-bonds weaken
-new bonds form

68
Q

transition state theory

A

-the orbitals of molecules interact and distort each other, weakening bonds
-bonds break
-kinetic energy becomes potential energy
-new bonds form
-energy is released
-activated complex forms

69
Q

activated complex/transition state

A

forms when potential energy is at a maximum and kinetic energy is maximized

-intermediate stage in reaction
-high energy, unstable arrangement of atoms is formed

70
Q

temperature dependence of the reaction is contained in what part of the equation?

A

rate constant, k

increase in temp = increase in k

71
Q

Arrhenius equation

A

k = Ae^(-E0/RT)

R: gas constant (8.314 J/mol x K)
Ea: energy barrier that must be surmounted for reactants to become products
A: constant that describes the number of times that reactants approach the activation barrier per unit time
E: number between 0-1 that represents fraction of successful approaches

72
Q

overall reaction is _. but first must absorb _ to reach the activated state. this _ _ is _ _.

A

exothermic, energy, absorbed energy, activation energy

73
Q

higher activation rate = _ reaction

A

slower

74
Q

units of frequency factor are _ as the rate constant depending on

A

the same, order of the reaction

75
Q

e: exponential factor definition

A

depends on temp and activation energy
-low Ea and high temp result in small negative exponent, approaches 1 (successful)
-vice versa for opposite conditions
-as temp increases, number of molecules that overcome barrier increases

76
Q

reaction mechanism

A

series of individual chemical steps by which an overall chemical reaction occurs

77
Q

elementary step

A

each step in a reaction mechanism

-cannot be broken down into simpler steps
-represents colliding species in an experiment

78
Q

individual steps in the mechanism add up to:

A

overall reaction

79
Q

reaction intermediate

A

forms in one elementary step, consumed in another

80
Q

molecularity

A

number of reactant particles involved in the step

most common are unimolecular and bimolecular

81
Q

how can we deduce the rate law for an overall reaction?

A

by using an elementary step from its equation

82
Q

rate and reactant proportion

A

rate is proportional to concentrations of reactant particles

83
Q

rate determining step

A

slowest of all elementary step
-first step
-limits overall rate of reaction
-determines rate law

84
Q

comparison between step 1 and 2 elementary steps

A

1: larger activation energy, smaller rate constant, determines overall rate of reaction

2: smaller activation energy

85
Q

valid mechanism conditions:

A
  1. elementary steps must add up to overall equation
  2. rate law predicted must be consistent with experimentally observed rate law
86
Q

mechanism with a slow initial step:

A

rate law contains only reactants involved in overall reaction

87
Q

mechanism with a fast initial step:

A

some other subsequent step in the mechanism is the rate-limiting step

88
Q

rate law predicted by rate-limiting step may contain

A

reaction intermediates

89
Q

what happens when there is a fast first elementary step?

A

-products build up
-products react with each other and re-form the reactants
-first step reaction reaches equilibrium where (rate of forward reaction = rate of reverse reaction)

90
Q

2 ways to speed up the rate of a reaction

A

increase concentration of reactants
increase the temperature

91
Q

catalyst

A

substance that increases the rate of reaction, not consumed by the reaction

provides alternate mechanism for a reaction

92
Q

homogeneous catalyst

A

catalysts exist in the same phase (state) as reactants

93
Q

heterogenous catalyst

A

catalyst is in a different phase than the reactants

94
Q

enzymes

A

biological catalysts that are shape-specific to their substrate
-active site that attracts substrate by intermolecular forces

95
Q

why is the reaction rate for reactants defined as the negative change in reactants concentration with respect to time (a positive value?)

A

-reactant concentration decreases as the reaction progresses, so the change in concentration has a negative value
-negative sign in definition makes reaction rate positive in terms of reactants

96
Q

whats the difference between avg rate of reaction and instantaneous rate of reaction?

A

avg: average rate in a time interval
instantaneous: rate determined at one point in time

97
Q

How is the order of a reaction generally determined?

A

By experiment-usually the method of initial rates. The initial rate is measured by running the reaction several times with differing concentrations to determine the effect of concentration on rate

98
Q

For a reaction with multiple reactants, how is the overall order of the reactants determined?

A

add the orders for each of the reactants

99
Q

Explain the difference between the rate law for a reaction and the integrated rate law for a reaction.

A

The rate law for a reaction is the relationship between the rate of the reaction and the [reactant]

The integrated rate law is the relationship between the [reactant] and time

100
Q

What are the two conditions that must be satisfied in order for a reaction to occur?

A

Molecules must have sufficient kinetic energy when they collide

Must have the proper orientation

101
Q

Explain how an activated complex forms

A

Reactant molecules collide. If they have enough kinetic energy, bonds between the atoms in the molecules will break

As the kinetic energy of the atoms decreases, the kinetic energy released is converted to potential energy in a new unstable arrangement of atoms.

Due to the high energy of this arrangement, the atoms rearrange to a new lower potential energy state to form the products.

102
Q

How do reaction rates typically depend on temperature?

A

Reaction rates are sensitive to temperature. An increase in temperature usually leads to an increased rate of reaction.

103
Q

What part of the rate law is temperature dependent?

A

The rate constant is temperature dependent. An increase in temperature results in an increase in k.

104
Q

Explain the meaning of each term within the Arrhenius equation: activation energy

A

Activation energy (Ea): energy barrier that must be surmounted for reactants to become products.

105
Q

Explain the meaning of each term within the Arrhenius equation: frequency factor

A

a constant that describes the number of times that reactants approach the activation barrier per unit time.

106
Q

Explain the meaning of each term within the Arrhenius equation: exponential factor

A

number between 0 and 1 that represents the fraction of approaches that are successful and result in product.

A small decrease in temperature makes the value of the exponential factor lower in the Arrhenius equation and so a lower value of k. Less approaches surmount the activation barrier. A small increase in temperature has the opposite effect.

107
Q

Explain how a chemical reaction occurs according to the collision model. Explain the meaning of the exponential factor in this model

A

A chemical reaction occurs after a sufficiently energetic collision between two reactant molecules. The two molecules must collide with sufficient energy to overcome the activation barrier and must also have the correct orientation.

The exponential factor has a value between 0 and 1 and represents the fraction of collisions with an orientation that allows reactions to occur.

108
Q

Explain the difference between a normal chemical equation for a chemical reaction and the mechanism of that reaction

A

A normal chemical equation shows only reactants and products. It does not show intermediate steps.

A reaction mechanism is a complete, detailed description of the reaction at the molecular level. It specifies the individual collisions and reactions that result in the overall reaction.

109
Q

In a reaction mechanism, what is an elementary step? What does it represent?

A

a step in an overall reaction and cannot be broken down into simpler steps.

It represents the species that are colliding in the reaction.

110
Q

Write down the three most common elementary steps and the corresponding rate law for each one.

A

Unimolecular
Rate = k [A]
Bimolecular
Rate = k [A]2 or Rate = k [A] [B]
Termolecular
Rate = k [A]3 or Rate = k [A]2 [B] or rate = k [A] [B] [C]

111
Q

What are the two requirements for a proposed mechanism to be valid for a given reaction?

A

The elementary steps in the mechanism must sum to the overall reaction

The rate law predicted by the mechanism must be consistent with the experimentally observed rate law.

112
Q

What is an intermediate within a reaction mechanism?

A

A species that forms in one elementary step and is consumed in another

ap chem (11 decks)

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