unit 8 Flashcards
1
Q
the rate of a chemical reaction measures
A
how fast the reaction occurs
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
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
4
Q
rate definition
A
change in a quantity (concentration of reactant/product) per unit time
5
Q
rate (reactant) equation
A
- (change in concentration) / (change in time)
6
Q
reactant concentration _ over time
A
decreases
7
Q
product concentration _ over time
A
increases
8
Q
rate (product) equation
A
(change in concentration) / (change in time)
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
10
Q
reactant concentration _ with time as reactants are consumed
A
decreases
11
Q
product concentration _ with time as products are formed
A
increases
12
Q
4 factos that affect the rate of the reaction
A
physical state
reactant concentration
reaction temp
presence of a catalyst
13
Q
the more readily reactant molecules collide with each other, the more _ they react
A
rapidly
14
Q
reactant concentration and reaction rate proportion
A
directly proportional
15
Q
temp and reaction rate proportion
A
direct (generally)
16
Q
catalyst definition
A
agents that increase reaction rates without themselves being used up
17
Q
2 methods of measuring reaction concentrations
A
spectroscopy
monitoring the pressure of gases
18
Q
how to measure the reactant concentrations from spectroscopy:
A
measure the change in absorbance of light
calculating concentration using beer’s law
19
Q
rate law definition
A
relationship between rate of reaction and concentration of the reactant
20
Q
rate law formula
A
rate = k[A]^n
k: rate constant
n: reaction order
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
22
Q
first order
A
n = 1
-rate slows as reaction proceeds
-rate is slower as reaction proceeds
-rate is directly proportional to concentration
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
24
Q
the order of a reaction can only be determined by _
A
experiment
25
method of initial rates
initial rate is measured by running the reaction several different times with different initial [reactant] to determine the effect of [reactant] on rate
26
initial rate definition
a short period of time at the beginning of a reaction
27
if reaction is 0 order in A
-initial rate is independent of [reactant]
-rate is the same for all measured initial concentrations
-rate = k
-units: M x s^-1
28
if the reaction is 1st order in A
-when [A] doubles, rate doubles
-inital rate is directly proportional to initial concentration
-rate = k[A]^1
-units: s^-1
29
if the reaction is second order in A
-when [A] doubles, rate quadruples
-relationship between concentration and rate is quadratic
-rate=k[A]^2
-units: M^-1 x s^-1
30
overall order
sum of exponents m and n
m = A's reaction order
n = B's reaction order
31
differential rate law (first order) equation
k[A] = -(Delta [A]) / (delta t)
32
integrated rate law (first order) equation
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]
33
integrated rate law graph
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
34
differential rate law (second order) equation
k[A]^2 = -(Delta [A]) / (delta t)
35
integrated rate law (second order) equation
(1) / ([A]t) = kt + (1 / [A]0)
36
integrated rate law (zero order) equation
k = -(Delta [A]) / (delta t)
37
differential rate law (zero order) equation
[A]t = -kt + [A]0
38
half life definition
time required for a reactant to fall 1/2 of its original value
39
first order half life equation
0.693 / k
40
for first order reactions, t1/2 is
independent of initial concentration
-even though [reactant] changes as reaction proceeds, half life is constant
41
second order half life equation
1 / (k[A]0)
42
zero order half life equation
[A]0 / 2k
43
half lives of zero and second order reactions depend on
initial concentration
44
rate law zero order
k[A]⁰
45
rate law first order
k[A]¹
46
rate law second order
k[A]²
47
zero order unit of k
M x s⁻¹
48
first order unit of k
s⁻¹
49
second order unit of k
M⁻¹ s⁻¹
50
the potential energy diagram represents the ___ ___ that occur during a chemical reaction
energy changes
51
y axis in a potential energy diagram represents
potential energy in kJ/mol
52
x axis in a potential energy diagram represents
progression of the reaction from reactants activated complex to products
53
chemical energy is sometimes referred to as
bond energy
54
chemical energy/bond energy is a form of _ energy
potential
55
bond breaking requires
energy input
-breaking electrostatic forces
56
bond forming results in
energy release
-results in more stable situation
57
endothermic reaction
more energy is absorbed as bonds break than is released when bonds formed
positive delta H
58
exothermic reaction
more energy is released as bonds form than absorbed as they break
negative delta H
59
activation energy (Ea) definition
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
activation energy formula
Ea = PE(activated complex) - PE(reactants)
61
enthalpy change (delta H) definition
difference between the total potential energy of the products and the total potential energy of the reactants
62
enthalpy change formula
delta H = PE(products) - PE(reactants)
63
collision theory
reaction rate is equal to frequency of effective collisions between reactants
64
in order for a reaction to occur, molecules must
-collide with sufficient energy
-properly oriented so products can form
65
the fraction of collisions with minimum energy (activation energy) increases with
rising temp
-avg kinetic energy of molecules increases with temp
66
are reactions more likely to occur with molecules or single atoms?
single atoms
-greater chance of forming a product due to its orientation
-spherical
67
generally, what happens when gases collide at normal temp?
-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
transition state theory
-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
activated complex/transition state
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
temperature dependence of the reaction is contained in what part of the equation?
rate constant, k
increase in temp = increase in k
71
Arrhenius equation
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
overall reaction is _. but first must absorb _ to reach the activated state. this _ _ is _ _.
exothermic, energy, absorbed energy, activation energy
73
higher activation rate = _ reaction
slower
74
units of frequency factor are _ as the rate constant depending on
the same, order of the reaction
75
e: exponential factor definition
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
reaction mechanism
series of individual chemical steps by which an overall chemical reaction occurs
77
elementary step
each step in a reaction mechanism
-cannot be broken down into simpler steps
-represents colliding species in an experiment
78
individual steps in the mechanism add up to:
overall reaction
79
reaction intermediate
forms in one elementary step, consumed in another
80
molecularity
number of reactant particles involved in the step
most common are unimolecular and bimolecular
81
how can we deduce the rate law for an overall reaction?
by using an elementary step from its equation
82
rate and reactant proportion
rate is proportional to concentrations of reactant particles
83
rate determining step
slowest of all elementary step
-first step
-limits overall rate of reaction
-determines rate law
84
comparison between step 1 and 2 elementary steps
1: larger activation energy, smaller rate constant, determines overall rate of reaction
2: smaller activation energy
85
valid mechanism conditions:
1. elementary steps must add up to overall equation
2. rate law predicted must be consistent with experimentally observed rate law
86
mechanism with a slow initial step:
rate law contains only reactants involved in overall reaction
87
mechanism with a fast initial step:
some other subsequent step in the mechanism is the rate-limiting step
88
rate law predicted by rate-limiting step may contain
reaction intermediates
89
what happens when there is a fast first elementary step?
-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
2 ways to speed up the rate of a reaction
increase concentration of reactants
increase the temperature
91
catalyst
substance that increases the rate of reaction, not consumed by the reaction
provides alternate mechanism for a reaction
92
homogeneous catalyst
catalysts exist in the same phase (state) as reactants
93
heterogenous catalyst
catalyst is in a different phase than the reactants
94
enzymes
biological catalysts that are shape-specific to their substrate
-active site that attracts substrate by intermolecular forces
95
why is the reaction rate for reactants defined as the negative change in reactants concentration with respect to time (a positive value?)
-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
whats the difference between avg rate of reaction and instantaneous rate of reaction?
avg: average rate in a time interval
instantaneous: rate determined at one point in time
97
How is the order of a reaction generally determined?
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
For a reaction with multiple reactants, how is the overall order of the reactants determined?
add the orders for each of the reactants
99
Explain the difference between the rate law for a reaction and the integrated rate law for a reaction.
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
What are the two conditions that must be satisfied in order for a reaction to occur?
Molecules must have sufficient kinetic energy when they collide
Must have the proper orientation
101
Explain how an activated complex forms
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
How do reaction rates typically depend on temperature?
Reaction rates are sensitive to temperature. An increase in temperature usually leads to an increased rate of reaction.
103
What part of the rate law is temperature dependent?
The rate constant is temperature dependent. An increase in temperature results in an increase in k.
104
Explain the meaning of each term within the Arrhenius equation: activation energy
Activation energy (Ea): energy barrier that must be surmounted for reactants to become products.
105
Explain the meaning of each term within the Arrhenius equation: frequency factor
a constant that describes the number of times that reactants approach the activation barrier per unit time.
106
Explain the meaning of each term within the Arrhenius equation: exponential factor
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
Explain how a chemical reaction occurs according to the collision model. Explain the meaning of the exponential factor in this model
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
Explain the difference between a normal chemical equation for a chemical reaction and the mechanism of that reaction
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
In a reaction mechanism, what is an elementary step? What does it represent?
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
Write down the three most common elementary steps and the corresponding rate law for each one.
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
What are the two requirements for a proposed mechanism to be valid for a given reaction?
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
What is an intermediate within a reaction mechanism?
A species that forms in one elementary step and is consumed in another
