Exam II Chemistry

Chemical kinetics
how the molecular world changes with time
Thermal
__ energy produces constant molecular motion, causing molecules to repeatedly collide with one another.
-nuclei
-kinetics
In a small fraction of the collisions caused by thermal energy, the electrons on one molecule or atom are attracted to the __ of another. Some bonds weaken and new bonds form – a chemical reaction occurs. Chemical __ is the study of how these kinds of changes occur in time
Rate of reaction
__ = how fast a reaction occurs
life
Reaction rates are important to __
enzymes
Rate of reaction is controlled through the use of __
Wilhelmy
__ = first person to measure the rate of a chemical reaction
molecular
The rate of a reaction can tell us much about how the reaction occurs on the __ scale
reaction
The rate of a chemical reaction is a measure of how fast the __ occurs
Slow rate of reaction
__ = small fraction of molecules react to form products in a given period of time
Fast rate of reaction
__ = large fraction of molecules react to form products in a given period of time
change in some quantity per unit of time
When measuring the rate at which something occurs, we express the measurement as a __
-measuring
-time
When measuring the rate at which something occurs, we report the rate in units that represent the change in what we are __ divided by the change in __
-reactants or products (usually in concentration units)
-time
The rate of a chemical reaction is measured as a change in the amounts of __ divided by the change in __
Reaction rate
__ – the negative of the change in concentration of a reactant divided by the change in time
-reactant
-decrease
The negative sign is part of the definition when the reaction rate is defined with respect to a __ because reactant concentrations __ as a reaction proceeds; therefore the change in the concentration of a reactant is negative, which makes the overall rate positive
positive
Reaction rates are reported as __ quantities
product
The reaction rate can also be defined with respect to the __ of the reaction
-increase
-positive
Product concentrations __ as the reaction proceeds, the change in concentration of a product is __
-don’t include
– positive
When the rate is defined with respect to a product, __ a negative sign – the rate is naturally __
stoichiometric coefficients
To have a single rate for an entire reaction, the definition of the rate with respect to each reactant and product must reflect the __ of the reaction
-decreases
-consumed
The reactant concentration __ with time because reactants are __ in a reaction
-increases
-formed
The product concentration __ with time because products are __ in a reaction
Rate = – ?[molecule] / ?t
Calculate the average rate of the reaction for any time interval using the equation: __
– decreases
-slows down
The average rate __ as the reaction progresses, the reaction __ as it proceeds
concentrations of the reactants
The average rate of the reaction depends on the __
-decrease
-slows down
As reactants transform to products, their concentrations __, and the reaction __
Instantaneous rate of the reaction
__ = the rate at any one point in time, and is represented by the instantaneous slope of the curve at that point
slope of the tangent to the curve
Determine the instantaneous rate from the __ at the point of interest
-(1 / a)(?[A] / ?t) = -(1 / b)( ?[B] / ?t) = +(1 / c)( ?[C] / ?t) = +(1 / d)( ?[D] / ?t)

A and B

C and D

a, b, c, and d

Rate of the reaction > Rate = __
__ are reactants, __ are products, and __ are the stoichiometric coefficients
any other reactant or product at that point in time (from the balanced equation)
Knowing the rate of change in the concentration of any one reactant or product at a point in time allows us to determine the rate of change in the concentration of __
not possible
Predicting the rate at some future time is __ from just the balanced equation
kinetics
Must have an experimental way to measure the concentration of at least one of the reactants or products as a function of time to study the __ of a reaction
Polarimetry (experimental way to measure the concentration of at least one of the reactants or products as a function of time)
__ – measuring the degree of polarization of light passing through a reacting solution to determine the relative concentrations of the reactants and products as a function of time
spectroscopy (experimental way to measure the concentration of at least one of the reactants or products as a function of time)
Most common way to study the kinetics of a reaction is through __
Spectrometer (experimental way to measure the concentration of at least one of the reactants or products as a function of time)
__ – device that passes light through a sample and measures how strongly the light is absorbed.
spectrometer (experimental way to measure the concentration of at least one of the reactants or products as a function of time)
In using a __, if the sample contains the reacting mixture, the intensity of the light absorption will decrease as the reaction proceeds, providing a direct measure of the concentrations of the molecule as a function of time
pressure (experimental way to measure the concentration of at least one of the reactants or products as a function of time)
Reactions in which the number of moles of gaseous reactants and products changes as the reaction proceeds can be monitored by measuring changes in __
-increases
-rises
-rise
As a reaction proceeds and the amount of gas __, the pressure steadily __. The __ in pressure can be used to determine the relative concentrations of reactants and products as a function of time
polarimetry, spectroscopy, and pressure measurement
Three techniques – __ – can be used to monitor the reaction as it occurs in the reaction vessel
progress of the reaction
Some reactions occur slowly enough that samples, or aliquots can be periodically withdrawn from the reaction vessel and analyzed to determine the __
sample
__ = aliquot
aliquot
Instrumental techniques – gas chromatography, mass spectrometry – can be used to measure the relative amounts of reactants or products in the __
reactants and products
Taking aliquots at regular time intervals, can determine the relative amounts of __ as a function of time
concentration
Rate of a reaction depends on the __ of one or more of the reactants
rate law
As long as the rate of the reverse reaction (in which the products return to reactants) is negligibly slow, we can write a relationship – called the __ – between the rate of the reaction and the concentration of the reactant
k[A]
Rate law equation = __

 k = constant of proportionality called the rate constant

n = reaction order. The value of this determines how the rate depends on the concentration of the reactant

in the rate law equation k[A]n , k = __, n = __

zero order

independent of the concentration of A

if n = 0, the reaction is __ and the rate is __

first order

directly proportional to the concentration of A

ifn = 1, the reaction is __ and the rate is __

-second order

-proportional to the square of the concentration of A

if n = 2, the reaction is __, and the rate is __

concentration
in a zero-order reaction, the rate of the reaction is independent of the __ of the reactant
k[A]0 = k
if the reaction is a zero-order reaction, rate =

decreases

constant

does not

in a zero-order reaction,the concentration of the reactant __ linearly with time at a __ rate because the reaction ___ slow down as the concentration of A decreases
same
The rate of a zero-order reaction is the __ at any concentration of A
Zero-order
__ reactions occur under conditions where the amount of reactant actually available for reaction is unaffected by changes in the overall quantity of reactant
zero order
Sublimation is normally a __ reaction because only molecules at the surface can sublime, and their concentration does not change when the amount of subliming substance decreases
directly proportional
in a first-order reaction,the rate of the reaction is __ to the concentration of the reactant
k[A]1
in a first-order reaction, rate = __

slows down

decreases

For first-order reaction the rate __ as the reaction proceeds because the concentration of the reactant __
first-order
in a __ reaction,the rate is directly proportional to the concentration
proportional to the square
in a second-order reaction,the rate of the reaction is __ of the concentration of the reactant
k[A]2
for a second-order reaction, rate = __
sensitive
For second-order reaction, the rate is __ to the reactant concentration
second-order
for a __ reaction,the rate is proportional to the square of the concentration
experiment
The order of a reaction can be determined only by __
Method of initial rates
___ = common way to determine reaction order method of initial rates 
the initial rate
  __ – the rate for a short period of time at the beginning of the reaction
method of inital rates

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

initial rate

initial concentration


For a reaction that is a first-order reaction, the __ is directly proportional to the __
k[A]1
for a reaction that is a first-order reaction, Rate = __
solving the rate law for k and substituting the concentration and the initial rate from any one of the measurements found through experiment
For a reaction that is a first-order reaction, determine the value of the rate constant, k, by __
s-1
For a first-order reaction, the rate constant has units of __

same

initial

For a zero-order reaction, the initial rate is independent of the reactant concentration – the rate is the __ at all measured __ concentrations

quadruples

doubling

quadratic

For a second-order reaction, the initial rate __ for a __ of the reactant concentration – the relationship between concentration and rate is __
second-order

 For a __ reaction you can substitute any two initial concentrations and the corresponding initial rates into a ratio of the rate laws to determine the order (n): rate 2 / rate 1 = k[A]n2 / k[A]n1

M x s-1

 The rate constant for a zero-order reaction has units of __

M-1 x s-1

The rate constant for a second-order reaction has units of __

multiplied
reaction order for multiple reactants:As long as the reverse reaction is negligibly slow, the rate law is proportional to the concentration of [A] raised to the m __ by the concentration of [B] raised to the n
k[A]m[B]n
Reaction order for multiple reactants; Rate = __

-A and B

-m

-n

in the reaction order for multiple reactions,Rate = k[A]m[B]n

__ are reactants

 

__ is the reaction order with respect to A

__ is the reaction order with respect to B

 

sum of the exponents (m + n)
in the reaction order for multiple reactants,the overall order is the __

experiment

method of initial rates

the rate law for any reaction must always be determined by __, often by the __
rate law
no simple way to look at a chemical equation and determine the __ for the reaction
independently
in determining the reaction order for multiple reactants,when there are two or more reactants, the concentration of each reactant is usually varied __ of the others to determine the dependence of the rate on the concentration of that reactant
Integrated rate law
__ for a chemical reaction is a relationship between the concentration of the reactants and time
order of the reaction
The integrated rate law for a reaction depends on the __
ln[A]t = -kt + ln[A]0 OR     ln ([A]t / [A]0) = -kt
what is the first-order integrated rate law?

[A]t

k

[A]0

for all of the integrated rate laws:

__ = the concentration of A at any time, t

__ = rate constant

__ = initial concentration of A

straight line

ln[A]t = -kt + ln[A]0 –>

y       = mx +  b

 

the integrated rate law has the form of an equation for a __

straight line

–k

ln[A]0

for a first-order reaction, a plot of the natural log of the reactant concentration has a function of time yields a  __ with a slope of __and a y-intercept of __

negative

positive

for a first-order reaction, slope is __, but the rate constant is always __
1 / [A]t = kt + 1 / [A]0
what is the second-order integrated rate law?

straight line

 1 / [A]t = kt + 1 / [A]0 –>

y         = mx +  b

 

The second-order integrated rate law is also in the form of an equation for a __

inverse

straight line

k

1 / [A]0

for a second-order reaction, plot the __ of the concentration of the reactant as a function of time, which yields a __ with a slope of __ and an intercept of __
[A]t = -kt + [A]0
what is the zero-order integrated rate law?

straight line

[A]t = -kt + [A]0 –>

y     = mx +  b

The zero-order integrated rate law is also in the form of an equation for a __

straight line

–k

[A]0

for a zero-order reaction, a plot of the concentration of the reactant as a function of time yields a __ with a slope of __ and an intercept of __
Half-life (t1/2)
__ of a reaction is the time required for the concentration of a reactant to fall to one-half of its initial value

rate constant

initial concentration

The half-life expression defines the dependence of half-life on the __ and the __ 
different
the half-life expression is __ for different reaction orders
t1/2 = 0.693 / k
what is the half-life of a first-order reaction?
independent
For a first-order reaction, t1/2 is __ of the initial concentration
constant
Even though the concentration is changing as the reaction proceeds, the half-life (how long it takes for the concentration to halve) is __ for a first-order reaction half-life 
first-order
Constant half-life is unique to __ reactions
t1/2 = 1 / k[A]0
what is the half-life of a second-order reaction?
initial concentration
For a second-order reaction, the half-life depends on the __
decreases
The half-life continues to get longer as the concentration __ for second-order reactions
t1/2 = [A]0 / k[A]0
what is the half-life or a zero-order reaction?
initial concentration
For a zero-order reaction, the half-life depends on the __

reaction order

rate law

The __ and __ must be determined experimentally
reactant(s)
The rate law relates the rate of the reaction to the concentration of the __

reactant(s)

time

The integrated rate law (which is mathematically derived from the rate law) relates the concentration of the __ to __

one-half of its initial value
The half-life is the time it takes for the concentration of a reactant to fall to __
first-order
The half-life of a __ reaction is independent of the initial concentration

zero-order

second-order

The half-lives of __ and __ reactions depend on the initial concentration
C) the reaction is most likely second order because its rate depends on the concentration (therefore it cannot be zero order), and its half-life depends on the initial concentration (therefore is cannot be first order). For a second-order reaction, a doubling of the initial concentration results in the quadrupling of the rate

A decomposition reaction, with a rate that is observed to slow down as the reaction proceeds, is found to have a half-life that depends on the initial concentration of the reactant. Which of the following is most likely to be true of this reaction?

A) a plot of the natural log of the concentration of the reactant as a function of time will be linear

B) the half-life of the reaction increases as the initial concentration increases

C) A doubling of the initial concentration of the reactants results in a quadrupling of the rate

temperature
The rates of chemical reactions are highly sensitive to __
rate constant, k (which is actually a constant only when the temperature remains constant)
The rate law for a reaction is Rate = k[A]n , the temperature dependence of the reaction rate is contained in the __

increase

faster

Increase in temperature generally results in an __ in k, which results in a __ rate
Arrhenius equation
__ – shows the relationship between the rate constant (k) and the temperature in kelvins (T)
Kelvins
The Arrhenius equation has to have the temperature in __
k = Ae-Ea / RT
what is the Arrhenius equation?

A

-Ea

e-Ea / RT

R

in the Arrhenius equation,

__ = constant called the frequency factor (or pre-exponential factor)

__ = activation energy (or activation barrier)

__ = exponential factor

__ = gas constant

8.314 J / mol x K
the gas constant (R) for the Arrhenius equation is what?
Activation energy (Ea)
__ is an energy barrier or hump that must be surmounted for the reactants to be transformed into products
Frequency factor (A)
__ is the number of times that the reactants approach the activation barrier per unit time
activated complex, or transition state
To get from the reactant to the product, the molecule must go through a high-energy intermediate state called the __
initially weaken the chemical bonds
The overall reaction is energetically downhill (exothermic), but it must first go uphill to reach the activated complex because energy is required to __
activation energy
The energy required to reach the activated complex is the __
slower
The higher the activation energy, the __ the reaction rate (at a given temperature)
approaches
Frequency factor represents the number of __ to the activation barrier per unit time
not equivalent
Approaching the activation barrier is __ to surmounting it
do not
Most of the approaches to the activation barrier __ have enough total energy to make it over the activation barrier
Exponential factor
__ is a number between 0 and 1 that represents the fraction of molecules that have enough energy to make it over the activation barrier on a given approach

successful

product

Exponential factor is the fraction of approaches that are actually __ and result in the __

temperature (T)

activation energy (Ea)

Exponential factor depends on both the __ and the __ of the reaction
one
Low activation energy and a high temperature make the negative exponent small, so that the exponential factor approaches __

one (e-0 = 1)

successful

If the activation energy is zero, then the exponent is zero, and the exponential factor is exactly __. every approach to the activation barrier is __

negative number

very small

A large activation energy and a low temperature make the exponent a very large __, so that the exponential factor become __
zero (einfinitey = 0)
As the temperature approaches 0 K, the exponent approaches an infinitely large number, and the exponential factor approaches __
increases
As the temperature increases, the number of molecules having enough thermal energy to surmount the activation barrier __
distribution
At any given temperature, a sample of molecules will have a __ of energies
small

Commonly, only a __ fraction of molecules have enough energy to make it over the activation barrier

large
A small change in temperature results in a __ difference in the number of molecules having enough energy to surmount the activation barrier
frequency factor
The __ is the number of times that the reactants approach the activation barrier per unit time
exponential factor
The __ is the fraction of approaches that are successful in surmounting the activation barrier and forming products

increases

decreases

The exponential factor__ with increasing temperature, but __ with increasing activation energy
ln k = – Ea / R (1 / T) + ln A
what is the equation for the Arrhenius plots?
straight line
The Arrhenius plots equation is in the form of a __
Arrhenius plot
A plot of the natural log of the rate constant (ln k) versus the inverse of the temperature in kelvins (1 / T) yields a straight line with a slope of –Ea / R and a y-intercept of ln A
kinetic
Arrhenius plot used in the analysis of __ data
rate constant
When either data are limited or plotting capabilities are absent, we can calculate the activation energy if we know the __ at just two different temperatures
ln (k2 / k1) = Ea / R (1 / T1 – 1 / T2)
what is the two-point form of the Arrhenius equation?
two different
the two-point form of the Arrhenius equation isused to calculate the activation energy from experimental measurements of the rate constant at __ temperatures
Collision model
__ – a chemical reaction occurs after a sufficiently energetic collision between two reactant molecules
Collision theory
__ – each approach to the activation barrier is a collision between the reactant molecules. So the value of the frequency factor should be the number of collisions that occur per second
smaller
Frequency factors of most gas-phase chemical reactions tend to be __ than the number of collisions that occur per second

 

 

k = Ae-Ea / RT = pze-Ea / RT

p = orientation factor

z = collision frequency

 

 

 

 

what two separate parts can the frequency factor be separated into?
collision frequency
__ – number of collisions that occur per unit time, can be calculated for a gas-phase reaction from the pressure of the gases and the temperature of the reaction mixture
109
single molecule undergoes on the order of __ collisions every second (typical conditions)
sufficient energy
for reaction to occur, molecules must collide with __
properly oriented
not all collisions with sufficient energy will lead to products because the reactant molecules must also be __
break and form

if two molecules are to react with each other, they must collide in such a way that allows the necessary bonds to __

small orientation factor
__ = orientational requirements are very stringent – the molecules must be aligned in a very specific way for the reaction to occur
1
reactions between individual atoms usually have orientation factors of approximately __, because atoms are spherically symmetric and thus any orientation can lead to the formation of products
don’t have to collide to react
a few reactions have orientation factors greater than one, which means that reactants __

nuclei

weaken

form

products

when two molecules with sufficient energy and the correct orientation collide the electrons on one of the atoms or molecules are attracted to the __ of the other; some bonds begin to __ while other bonds begin to __ and the reactants go through the transition state and are transformed into the __ –> how chemical reaction occurs
C) since the reactants in part A) are atoms, the orientation factors should be about one. The reactants in parts B) and C) are both molecules, so we expect orientation factors of less than one. Since the reactants in B) are symmetrical, we would not expect the collisions to have as specific an orientation requirement as in C), where the reactants are asymmetrical and must therefore collide in such way that a hydrogen atom is in close proximity to another hydrogen atom. Therefore, we expect C) to have the smallest orientation factor

Which reaction would you expect to have the smallest orientation factor?

A) H(g) + I(g) –> HI(g)

B) H2(g) + I2(g) –> 2 HI (g)

C) HCl(g) + HCl(g) –> H2(g) + Cl2(g)



overall reaction
In writing a chemical equation to represent a chemical reaction, we usually represent the __, not the series of individual steps by which the reaction occurs
Reaction mechanism
__ – series of individual chemical steps by which an overall chemical reaction occurs
elementary step
Each step in a reaction mechanism is an __
cannot
Elementary step – step that __ be broken down into simpler steps, they occur as they are written (they represent the exact species that are colliding in the reaction)
equal
For a reaction mechanism to be valid the individual steps in the mechanisms have to __ the overall reaction
Reaction intermediate
__ – forms in one elementary step and is consumed in another
Reaction mechanism
__ – complete, detailed description of the reaction at the molecular level – it specifies individual collisions and reactions that result in the overall reaction
measuring the kinetics of the overall reaction
Can put a reaction mechanism together by __ and working backward to write a mechanism consistent with the measured kinetics
molecularity
Elementary steps = characterized by their __ – number of reactant particles involved in the step

unimolecular

bimolecular

Most common molecularities = __ and __
unimolecular
A –> products = __
bimolecular
A + A –> products = __
bimolecular
A + B –> products = __
Termolecular
__ = elementary steps in which three reactant particles collide; very rare because of low probability of three particles colliding simultaneously
balanced chemical equation
Rate law for an elementary step can be deduced from the __ (rate law can’t be found for an overall chemical reaction with same method)
product
Elementary step occurs though the collision of the reactant particles so the rate law is proportional to the __ of the concentrations of those particles
A + B –> products; Rate = k[A][B]
rate for the bimolecular elementary step in which A reacts with B is proportional to the concentration of A multiplied by the concentration of B:
A + A –> products; Rate = k[A]2
rate law for the bimolecular step in which A reacts with A is proportional to the square of the concentration of A:

1

k[A]

A —> products ; molecularity = __ ; Rate = __

2

k[A]2

A+A –> products ; molecularity = __ ; Rate = __

2

k[A][B]

A+B –> products ; molecularity = __ ; Rate = __

3

k[A]3

A+A+A –> products; molecularity= __; Rate = __

3

k[A]2[B]

A+A+B –> products; molecularity=__; Rate = __

3

k[A][B][C]

A+B+C –> products; molecularity=__; Rate = __
overall order
The molecularity of the elementary step is equal to the __ of the step
rate-determining step
In most chemical reactions, one of the elementary steps – called the __ – is much slower than the others

limits

rate law

Rate-determining step in a reaction mechanism __ the overall rate of the reaction and determines the __ for the overall reaction

 

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

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

what two conditions must be met for a reaction mechanism to be valid?
validated
reaction mechanisms can only be __ not proven
is not
valid mechanism __ a proven mechanism because other mechanisms may also fulfill both of the requirements; therefore, we can only say that a given mechanism is consistent with kinetic observations of the reaction and therefore possible
validity
other types of data – such as the experimental evidence for a proposed intermediate – can further strengthen the __ of a proposed mechanism
overall reaction
When proposed mechanism for a reaction has a slow initial step, the rate law predicted by the mechanism normally contains only reactants involved in the __
rate limiting step
When a reaction mechanism begins with a fast initial step, some other subsequent step in the mechanism is the __.
reaction intermediates
When a reaction mechanism begins with a fast initial step, the rate law predicted by the rate-limiting step may contain __
experimental rate law
Reaction intermediates do not appear in the overall reaction equation, so a rate law containing intermediates cannot generally correspond to the __
concentrations of the reactants
We can express the concentration of intermediates in terms of the __ of the overall reaction
limited
In multistep mechanism where the first step is fast, the products of the first step can build up, because the rate at which they are consumed is __ by some slower step further along
re-form
In mechanism with a fast first step the build up of products from the first step react with one another to __ the reactants
equilibrium
In mechanism with a fast first step, as long as the first step is fast enough compared to the rate-limiting step, the first-step reaction will reach __
equals
both the forward reaction and the reverse reaction occur. If equilibrium is reached, then the rate of the forward reaction __ the rate of the reverse reaction

concentration of the reactants

temperature

Can speed up the rate of a reaction by increasing the __ or by increasing the __
catalyst
Reaction rates can be increased by using a __ – substance that increases the rate of a chemical reaction but is not consumed by the reaction
alternative mechanism
Catalyst provides an __ for the reaction – one in which the rate-determining step has a lower activation energy
homogeneous and heterogeneous
catalysis can be categoriaed into what two types?
Homogeneous catalysis
__ = the catalyst exists in the same phase (or state) as the reactants
Heterogeneous catalysis
__ = the catalyst exits in a different phase than the reactants
heterogeneous catalysis

Use of solid catalysts with gas-phase or solution-phase reactants is the most common type of __

alkenes
Second example of heterogeneous catalysis involves the hydrogenation of double bonds with __
1. adsorption, 2. diffusion, 3. reaction, 4. desorption
what are the four steps of a heterogeneous catalysis?
adsorption
__ is the first step of a heterogeneous catalysis. the reactants are absorbed onto the metal surface
diffusion
__ is the second step of a heterogeneous catalysis. the reactants diffuse on the surface until they approach each other
reaction
__ is the third step of a heterogeneous catalysis. the reactants react to form the products
desorption
__ is the fourth step of a heterogeneous catalysis. the products desorb from the surface into the gas phase
enzymes
Living organisms rely on __, biological catalysts that increase rates of biochemical reactions
Enzymes
__ – large protein molecules with complex 3-D structures. Within their structures there is a specific area called the active site
substrate
properties and shape of the active site are just right to bind the reactant molecule, called the __

lowered

faster

substrate fits into the active site perfectly, when substrate binds to the active site of the enzyme – through intermolecular forces – the activation energy of the reaction is greatly __, allowing the reaction to occur at a much __ rate

which

when

enzymes give living organisms control over __reactions occur, and __ they occur

specific

efficient

enzymes are extremely __ and __, speeding up reaction rates by factors of as much as a billion
catalyze
if living organism wants to turn a particular reaction on, it produces or activates the correct enzyme to __ the reaction
kinetics
Speed of a chemical reaction is determined by __
thermodynamics
Extent of a chemical reaction is determined by __
large
Reaction with a __ equilibrium constant proceeds nearly to completion – nearly all the reactants react to form products
small
Reaction with a __ equilibrium constant barely proceeds at all – nearly all the reactants remain as reactants, hardly forming any products
extent
Equilibrium constant is an experimentally measureable quantity and be used to predict and quantify the __ of a reaction
equilibrium constant, K
The concentrations of the reactants and products in a reaction at equilibrium are described by the __

right

high

low

 Large value of K = reaction lies far to the __ at equilibrium = __ concentration of products and a __ concentration of reactants

left

high

low

small value of K = reaction lies far to the __ at equilibrium = __ concentration of reactants and a __ concentration of products
how far a reaction proceeds
value of K = measure of __
toward the products
the larger the value of K, the more the reaction proceeds __

maintain equilibrium

counteract

any system at equilibrium, responds to changes in ways that _; if any of the concentrations of the reactants or products change, the reaction shifts to __ that change

larger

product

equilibrium for fetal hemoglobin is __ than the equilibrium constant for adult hemoglobin = the reaction tends to go farther in the direction of the __

increase

decrease

Reaction rates generally __ with increasing concentration of the reactants (unless the reaction order is zero) and __ with decreasing concentration of the reactants
reversible
A reaction that can proceed in both the forward and reverse directions is said to be __
equals
Dynamic equilibrium for a chemical reaction is the condition in which the rate of the forward reaction __ the rate of the reverse reaction

occurring

occurring at the same rate

In dynamic equilibrium, the forward and reverse reactions are still __; however, they are __
form at the same rate that they are depleted
When dynamic equilibrium is reached, the concentrations of the reactants and products no longer change. They remain the same because the reactants and products __
equal to one another
Because the concentrations of reactants and products no longer change at equilibrium does not mean that the concentrations of reactants and products are __ at equilibrium

most

small fraction

Some reactions reach equilibrium only after __ of the reactants have formed products

Some reactions reach equilibrium when only a __ of the reactants have formed products

reaction
When a reaction reaches equilibrium depends on the __
so small that it can be ignored
Nearly all chemical reactions are at least theoretically reversible. In many cases, the reversibility is __
no longer change
Equilibrium is reached in a chemical reaction when the concentrations of the reactants and products __
Dynamic equilibrium
__: rate of forward reaction = rate of reverse reaction. Concentrations of reactant(s) and product(s) no longer change

slows down

speeds up

As concentration of product increases, and concentrations of reactants decrease, rate of forward reaction __ and rate of reverse reaction __
constant
When equilibrium is reached, both the forward and reverse reactions continue, but at equal rates, so the concentrations of the reactants and products remain __
concentrations
When a chemical reaction reaches dynamic equilibrium the __ of reactants and products will not necessarily be equal at equilibrium

not equal

equal

The concentrations of reactants and products are __ at equilibrium; rather, the rates of the forward and reverse reactions are __

quantify

concentrations

The equilibrium constant is a way to __ the __ of the reactants and products at equilibrium

ratio

products

reactants

The equilibrium constant (K) for the reaction is defined as the __ – at equilibrium – of the concentrations of the __ raised to their  stoichiometric coefficients divided by the concentrations of the __ raised to their stoichiometric coefficients (Law of Mass Action)
K = [C]c[D]d / [A]a[B]b
what is the expression for the equilibrium constant (K) also known as the Law of Mass Action?

molarity of products

molarity of reactants

molar concentration of A (M; molarity)

in the equilibrium constant (K) equation (also known as the law of mass action) –>K = [C]c[D]d / [A]a[B]b

[C]c[D]d = __

[A]a[B]b = __

[A] = __

law of mass action
the relationship between the balanced chemical equation and the expression of the equilibrium constant is known as the __

balanced chemical equation

law of mass action

To express an equilibrium constant for a chemical reaction, examine the __ and apply the __
exponents
The coefficients in the chemical equation become the __ in the expression of the equilibrium constant
larger
A large equilibrium constant (K>>1) indicates that the numerator (which specifies the amounts of products at equilibrium) is __ than the denominator (which specifies the amounts of reactants at equilibrium). 
forward
when equilibrium constant is large (K>>1) the __ reaction is favored

far to the right

products

reactants

When the equilibrium constant is large, the equilibrium point for the reaction lies __ – high concentration of __, low concentrations of __
long time
A reaction with a large equilibrium constant may be kinetically very slow and take a __ to reach equilibrium
smaller
A small equilibrium constant (K<<1) indicates that the numerator (which specifies the amounts of products at equilibrium) is __ than the denominator (which specifies the amounts of reactants at equilibrium) 
reverse
when the equilibrium constant is small (K<<1) the __ reaction is favored

far to the left

reactants

products

When the equilibrium constant is very small, the equilibrium point for the reaction lies __ – high concentrations of __, low concentrations of __
K << 1
__ = reverse reaction is favored; forward reaction does not proceed very far
K approximately equal to 1
__ = neither direction is favored; forward reaction proceeds about halfway
K >> 1
__ = forward reaction is favored; forward reaction proceeds essentially to completion

(b) the reaction mixture will contain 1 mol of A and 10 mol of B so that [B]/[A] = 10

…it will contain more moles of B when compared to A because the equilibrium constant (K) is ‘large’ so the forward reaction is favored and the concentration of products is large and the concentration of reactants is small

The equilibrium constant for the reaction A (g) <–> B (g) is 10. A reaction mixture initially contains 11 mol of A and 0 mol of B in a fixed volume of 1 L. When equilibrium is reached, which statement is true?

(a) the reaction mixture will contain 10 mol of A and 1 mol of B

(b) the reaction mixture will contain 1 mol of A and 10 mol of B

(c) the reaction mixture will contain equal amounts of A and B

equilibrium constant
If a chemical equation is modified in some way, then the __ for the equation must be changed to reflect the modification

1. reversing the equation

2. multiplying the coefficients in the equation by a factor

3. adding two or more individual chemical equations to obtain an overall equation

what are the three common modifications for a chemical equation?
invert
because the equilibrium constant must be changed if a chemical equation is modified in some way, if you reverse a chemical equation __ the equilibrium constant to reflect the modification of the chemical equation
same factor
because the equilibrium constant must be changed if a chemical equation is modified in some way, if you multiply the coefficients in a chemical equation by a factor, raise the equilibrium constant to the __ to reflect the modification in the chemical equation
multiply the corresponding equilibrium constants
because the equilibrium constant must be changed if a chemical equation is modified in some way, if you add two or more individual chemical equations to obtain an overall equation, __ by each other to obtain the overall equilibrium constant to reflect the modification in the chemical equation
concentration
For gaseous reactions, the partial pressure of a particular gas is proportional to its __
equilibrium constant
We can also express the __ in terms of partial pressures of the reactants and products
Kc
__ = equilibrium constant with respect to the concentration in molarity
Kp
__ = equilibrium constant with respect to partial pressures in atmospheres
partial pressure of each gas
The expression for Kp takes the form of the expression for Kc, except that we use the __ in place of its concentration
PA
__ = partial pressure of gas A in units of atmospheres
not necessarily equal
Since the partial pressure of a gas in atmospheres is not the same as its concentration in molarity, the value of Kp for a reaction is __ to the value of Kc
ideally
As long as the gases are behaving __, we can derive a relationship between two constants (K and Kp)
number of moles of A (nA) divided by its volume (V) in liters: [A] = nA/V
The concentration of an ideal gas A is the __
PAV = nART –> PA = (nA/V)RT
From the ideal gas law, we can relate the quantity nA/V (concentration of an ideal gas ‘A’) to the partial pressure of A: __
PA = [A]RT or [A] = PA / RT
[A] = nA / V –> PA = __ or [A] = __
Kp = Kc(RT)c+d-(a+b)
what is the equation for Kp (the equilibrium constant with respect to partial pressures)?
delta n
__ = c + d – (a + b), which is the sum of the stoichiometric coefficients of the gaseous products minus the sum of the stoichiometric coefficients of the gaseous reactants
Kp = Kc(RT)delta n
Kp = __

0

equal


If the total number of moles of gas is the same after the reaction as before, then delta n = __, and Kp is __ to Kc
corresponding units
As long as concentration units are expressed in molarity for Kc and pressure units are expressed in atmospheres for Kp, we can skip the formality of units and enter the quantities directly into the equilibrium expression, dropping their __

reactants

products

Many chemical reactions involve pure solids or pure liquids as __ or __
does not
The concentration of a solid __ change because a solid does not expand to fill its container.
density
A solids concentration depends only on its __, which is a constant as long as some solid is present
are not
Pure solids – those reactants or products labeled in the chemical equation with an (s) – __ included in the equilibrium expression (because their constant value is incorporated into the value of K)
does not
The concentration of a pure liquid __ change
excluded
Pure liquids – reactants or products labeled in the chemical equation with an (l) – are also __ from the equilibrium expression
Heterogeneous
__ equilibrium – the concentration of solid carbon (the number of atoms per unit volume) is constant as long as some solid carbon is present. The same is true for pure liquids. Thus, the concentrations of solids and pure liquids are not included in equilibrium constant expressions
(b) since delta n for gaseous reactants and products is zero, Kp equals Kc

For which reaction does Kp = Kc

(a) 2Na2O2(s) + 2CO2(g) « 2Na2CO3(s) + O2(g)

(b) NiO(s) + CO(g) « Ni(s) + CO2(g)

(c) NH4NO3(s) « N2O(g) + 2H2O(g)

concentrations of the reactants and products
Most direct way to obtain an experimental value for the equilibrium constant of a reaction is to measure the __ in a reaction mixture at equilibrium
has no formal part in the calculation
Since equilibrium constants depend on temperature, many equilibrium problems will state the temperature even though it __

moles per liter (M)

unitless

The concentrations within Kc should always be written in __; however, the units are not normally included when expressing the value of the equilibrium constant, so Kc is __
equilibrium concentrations
The __ of the reactants and products will depend on the initial concentrations (and in general vary from one set of initial concentrations to another)
same
The equilibrium constant will always be the __ at a given temperature, regardless of the initial concentrations
the same
Whether you start with only reactants or only products, the reaction reaches equilibrium concentrations in which the equilibrium constant is __
equilibrium constant, K
No matter what the initial concentrations are, the reaction will always go in a direction so that the equilibrium concentrations – when substituted into the equilibrium expression – give the same __
the stoichiometry of the reaction
We need only know the initial concentrations of the reactant(s) and the equilibrium concentration of any one reactant or product. The other equilibrium concentrations can be deduced from __
ICE table
__ – table summarizing the initial conditions, the changes, and the equilibrium conditions; 
I = initial, C = change, E = equilibrium
for the ICE table, I = __, C = __, E = __

equilibrium concentrations

ICE table

to calculate the equilibrium constant, we can use the balanced  equation to write an expression for the equilibrium constant and the substitute the __ from the __
right (toward the products)
When the reactants of a chemical reaction mix, they generally react to form products – we say that the reaction proceeds to the __
magnitude of the equilibrium constant
The amount of products formed when equilibrium is reached depends on the __
reaction quotient
In order to compare the progress of a reaction to the equilibrium state of the reaction, we use a quantity called the __
any point in the reaction
Reaction quotient (Qc) – the ratio – at __ – of the concentrations of the products raised to their stoichiometric coefficients divided by the concentrations of the reactants raised to their stoichiometric coefficients
Qp
For gases with amounts measured in atmospheres, the reaction quotient uses the partial pressures in place of concentrations and is called __

equilibrium constant

reactant quotient

The difference between the reaction quotient and the equilibrium constant is that, at a given temperature, the __ has only one value and it specifies the relative amounts of reactants and products at equilibrium

The __ depends on the current state of the reaction and has many different values as the reaction proceeds

zero (Qc = 0)
In a reaction mixture containing only reactants, the reaction quotient is __
infinite (Qc = infinitey)
In a reaction mixture containing only products, the reaction quotient is __
one (Qc = 1)
In a reaction mixture containing both reactants and products, each at a concentration of 1M, the reaction quotient is __
reaction toward equilibrium
The reaction quotient is useful because the value of Q relative to K is a measure of the progress of the __
equal
At equilibrium, the reaction quotient is __ to the equilibrium constant

less than


If Q is__ K, Q must therefore get larger as the reaction proceeds toward equilibrium. Q becomes larger as the reactant concentration decreases and the product concentration increases – the reaction proceeds to the right
greater than
If Q is __ K, Q must therefore get smaller as the reaction proceeds toward equilibrium. Q gets smaller as the reactant concentration increases and the product concentration decreases – the reaction proceeds to the left
equal to
If Q is __ K the reaction is at equilibrium – the reaction will not proceed in either direction
progress of a reaction
The reaction quotient (Q) is a measure of the __ toward equilibrium

goes to the right (toward products)

goes to the left (toward reactants)

is at equilibrium

Q < K ; reaction __

Q > K ; reaction __

Q = K ; reaction __

reactant or product
Calculating equilibrium concentrations of reactants or products from the equilibrium constant allow us to calculate the amount of a __ at equilibrium
(1) finding equilibrium concentrations when we know the equilibrium constant and all but one of the equilibrium concentrations of the reactants and products; and (2) finding equilibrium concentrations when we know the equilibrium constant and only initial
Calculating equilibrium concentrations of reactants or products from the equilibrium constant can be divided into what two categories?

  1. Using the balanced equation as a guide, prepare a table showing the known initial concentrations of the reactants and products (ICE table)

2. Use the initial concentrations to calculate the reaction quotient (Q) for the initial concentrations. Compare Q to K to predict the direction in which the reaction will proceed.

3. Represent the change in the concentration of one of the reactants or products with the variable x. Define the changes in the concentrations of the other reactants or products in terms of x

4. Sum each column for each reactant and product to determine the equilibrium concentrations in terms of the initial concentrations and the variable x

5. Substitute the expressions for the equilibrium concentrations (from step 4) into the expression for the equilibrium constant. Using the given value of the equilibrium constant, solve the expression for the variable x.

6. Substitute x into the expressions for the equilibrium concentrations of the reactants and products (from step 4) and calculate the concentrations

7. Check your answer by substituting the computed equilibrium values into the equilibrium expression. The calculated value of K should match the given value of K

what is the procedure for finding equilibrium concentrations from initial concentrations and the equilibrium constant?

 1. Using the balanced equation as a guide, prepare a table showing the known initial partial pressures of the reactants and products

2. Use the initial partial pressures to calculate the reaction quotient (Q). compare Q to K to predict the direction in which the reaction will proceed

3. Represent the change in the partial pressure of one of the reactants or products with the variable x. Define the changes in the partial pressures of the other reactants or products in terms of x

4. Sum each column for each reactant and product to determine the equilibrium partial pressures in terms of the initial partial pressures and the variable x

5. Substitute the expressions for the partial pressures (from step 4) into the expression for the equilibrium constant. Use the given value of the equilibrium constant to solve the expression for the variable x

6. Substitute x into the expressions for the equilibrium partial pressures of the reactants and products (from step 4) and calculate the partial pressures

7. Check your answer by substituting the calculated equilibrium partial pressures into the equilibrium expression. The calculate value of K should match the given value of K

what is the procedure for finding equilibrium partial pressures when you are given the equilibrium constant and initial partial pressures?

1. Using the balanced equation as a guide, prepare a table showing the known initial concentrations of the reactants and products

2. Use the initial concentrations to calculate the reaction quotient (Q). Compare Q to K to predict the direction that the reaction will proceed

3. Represent the change in the concentration of one of the reactants or products with the variable x. Define the changes in the concentrations of the other reactants or products with respect to x

4. Sum each column for each reactant and product to determine the equilibrium concentrations in terms of the initial concentrations and the variable x

5. Substitute the expressions for the equilibrium concentrations (from step 4) into the expression for the equilibrium constant. Use the given value of the equilibrium constant to solve the expression for the variable x.

 

in this case, the resulting equation is cubic in x. Although cubic equations can be solved, the solutions are not usually simple. However, since the equilibrium constant is small, we know that the reaction does not proceed very far to the right. Therefore, x will be a small number and can be dropped from any quantities in which it is added to or subtracted from another number (as long as the number itself is not too small)

 

check whether your approximation was valid by comparing the calculated value of x to the number it was added to or subtracted from. The ratio of the two numbers should be less than 0.05 ( or 5%) for the approximation to be valid. If approximation is not valid proceed to step 5a

 

5a. if the approximation is not valid, you can either solve the equation exactly (by hand or with your calculator), or use the method of successive approximations. In this case, we use the method of successive approximations


 substitute the value obtained for x in step 5 back into the original cubic equation, but only at the exact spot where x was assumed to be negligible and then solve the equation for x again. Continue this procedure until the value of x obtained from solving the equation is the same as the one that is substituted into the equation

 

 

6. Substitute x into the expressions for the equilibrium concentrations of the reactants and products (from step 4) and calculate the concentrations

7. Check your answer by substituting the calculated equilibrium values into the equilibrium expression. The calculated value of K should match the given value of K. Note that the approximation method and rounding errors could cause a difference of up to about 5% when comparing values of the equilibrium constant

what is the procedure for finding equilibrium concentrations from initial concentrations in cases with a small equilibrium constant?
(a) the x is small approximation is most likely to apply to a reaction with a small equilibrium constant and an initial concentration of reactant that is not too small. The bigger the equilibrium constant and the smaller the initial concentration of reactant, the less likely that the x is small approximation will apply

For the generic reaction, A(g) « B(g), consider each value of K and initial concentration of A. For which set will the x is small approximation most likely apply?

(a) K = 1.0 x 10-5; [A] = 0.250M

(b) K = 1.0 x 10-2; [A] = 0.250M

(c) K = 1.0 x 10-5; [A] = 0.00250M

(d) K = 1.0 x 10-2; [A] = 0.00250M

minimize the disturbance
Le Chatelier’s principle states that the chemical system will respond to __ when a chemical system already at equilibrium is disturbed
Le Chatelier’s principle
__: when a chemical system at equilibrium is disturbed, the system shifts in a direction that minimizes the disturbance
changing the concentration of a reactant or product, changing the volume or pressure, and changing the temperature
how can we distrub a system in chemical equilibrium?

<

right (in the direction of the products)

If a chemical system is at equilibrium, increasing the concentration of one or more of the reactants (which makes Q __ K) causes the reaction to shift to the __

>

left (in the direction of the reactants)

If a chemical system is at equilibrium increasing the concentration of one or more of the products (which makes Q __ K) causes the reaction to shift to the __

>

left (in the direction of the reactants)

If a chemical system is at equilibrium decreasing the concentration of one or more of the reactants (which makes Q __ K) causes the reaction to shift to the __

<

right (in the direction of the products)

If a chemical system is at equilibrium decreasing the concentration of one or more of the products (which makes Q __ K) causes the reaction to shift to the __
pressure
Changing the volume of a gas (or a gas mixture) results in a change in __

inversely

increase

decrease

Pressure and volume are __ related: a decrease in volume causes an __ in pressure, and an increase in volume causes a __ in pressure
minimize that change
If the volume of a reaction mixture at chemical equilibrium is changed, the pressure changes and the system will shift in a direction to __
lower
From the ideal gas law (PV = nRT), we know that lowering the number of moles of a gas (n) results in a __ pressure (P)
higher
From the ideal gas law (PV = nRT), we know that increasing the number of moles of gas (n) results in a __ pressure (P)

increases

do not change

no

does not shift in either direction

If we keep the volume the same, but increase the pressure by adding an inert gas to the mixture, the overall pressure of the mixture __, but the partial pressures of the reactants and products __. There is __ effect and the reaction __
has the fewer moles of gas particles
If a chemical system is at equilibrium decreasing the volume causes the reaction to shift in the direction that __
greater number of moles of gas particles
If a chemical system is at equilibrium increasing the volume causes the reaction to shift in the direction that has the __
produces no effect on the equilibrium
If a chemical system is at equilibrium, if a reaction has an equal number of moles of gas on both sides of the chemical equation, then a change in volume __
no effect on the equilibrium
If a chemical system is at equilibrium, adding an inert gas to the mixture at a fixed volume has __
temperature
In considering the effect of a change in volume, we are assuming that the change in volume is carried out at a constant __
pressure
In considering the effect of a change in temperature, we are assuming that the heat is added (or removed) at constant __
counter that change
If the temperature of a system at equilibrium is changed, the system will shift in a direction to __
product in an exothermic reaction
An exothermic reaction (negative delta H) emits heat and we can think of heat as a __
reactant in an endothermic reaction
An endothermic reaction (positive delta H) absorbs heat and we can think of heat as a __

shift left

reactants

products

smaller

At constant pressure, raising the temperature of an exothermic reaction (think of this as adding heat) is similar to adding more product, causing the reaction to __. The new equilibrium mixture will have more __ and fewer __ and therefore a __ value of K
equilibrium constant
Changing the temperature does change the value of the __
exothermic
At constant pressure, lowering the temperature of an __reaction, causes the reaction to shift right, releasing heat and producing more products because the value of K has increased.

shift right

products

increased

For an endothermic reaction at constant pressure, raising the temperature (adding heat) causes the reaction to __ to absorb the added heat and producing more __ because the value of K has __

shift left

products

lowering

For an endothermic reaction at constant pressure, lowering the temperature (removing heat) of a reaction mixture causes the reaction to __, releasing heat, forming less __, and  __ the value of K

Increasing


In an exothermic chemical reaction, heat is a product. __ the temperature causes an exothermic reaction to shift left (in the direction of the reactants); the value of the equilibrium constant decreases
Decreasing
In an exothermic chemical reaction, heat is a product. __ the temperature causes an exothermic reaction to shift right (in the direction of the products); the value of the equilibrium constant increases
Increasing
In an endothermic chemical reaction, heat is a reactant. __ the temperature causes an endothermic reaction to shift right (in the direction of the products); the equilibrium constant increases
Decreasing
In an endothermic chemical reaction, heat is a reactant. __ the temperature causes an endothermic reaction to shift left (in the direction of the reactants); the equilibrium constant decreases
endothermic direction
Adding heat favors the __
exothermic direction
Removing heat favors the __
metastable
Systems that are not in
thermodynamic equilibrium
but are kinetically stable are
called __.
concentration, pressure, temperature
for a system at equilibrium, the principle can be used to qualitatively predict the effects of changes in what