CHMA11 ~ TEXTBOOK
NOTES
10.2 VESPR Theory: The Five Basic Shapes (400-403)
·
valence shell electron pair repulsion (VESPR)
theory: allows prediction of shapes of molecules
based on idea that electrons- either as lone pairs or as bonding pairs- repel
one another.
·
based on idea that electron
groups repel one another through coulombic forces
¾
electron groups:
lone pairs, single bonds, multiple bonds, or lone electrons in a molecule
·
repulsions between electron
groups on interior atoms of a molecule determine geometry of molecule
·
preferred geometry of a
molecule is one where electron groups have max separation (min energy) possible
·
for molecules with one interior
(central) atom, molecular geometry depends on
¾
number of electron groups
around central atom
¾
how many of those electron
groups are bonding and how many are lone pairs
Two
Electron Groups: Linear Geometry
·
linear geometry:
molecular geometry of 3 atoms with 180° bond angle due to repulsion of 2
electron groups
·
molecules that form only 2
single bonds, with no lone pairs, rare because they don’t follow octet rule
·
same geometry observed in all
molecules that have 2 electron groups and no lone pairs
·
Three
Electron Groups: Trigonal Planar Geometry
·
trigonal planar geometry: molecular geometry of four atoms with 120° bond angles in a plane
·
double bond contains more
electron density than single bond, exerts slightly greater repulsion on single
bonds
·
Four
Electron Groups: Tetrahedral Geometry
·
tetrahedral geometry:
molecular geometry of 5 ato ms with 109.5° bond angles
·
Five
Electron Groups: Trigonal Bipyramidal Geometry
·
trigonal bipyramidal geometry: molecular geometry of 6 atoms with 120° bond angles between 3
equatorial electron groups and 90° bond angles between the two axial electron
groups and the trigonal plane
·
Six
Electron Groups: Octahedral Geometry
·
octahedral geometry: molecular geometry of 7 atoms with 90° bond angles
·
10.8 Molecular Orbital Theory: Electron Delocalization (432-)
·
11.2, 11.3, 11.4, 11.5,
11.6, 11.7, 11.8, 11.11, 11.12, 12.2, 12.3, 12.4, 12.5, 12.6, 12.7
CHAPTER
13
13.2 The Rate of a Chemical Reaction
(564-569)
·
if chemical reaction has slow
rate, only relatively small fraction of molecules react to form products in
given period of time; fast rate means large fraction of molecules react
·
rate of chemical reaction
measured as change in amounts of reactants or products (usually in
concentration units) divided by change in time
·
reaction rate:
negative of change in concentration of a reactant divided by change in time
¾
negative sign usually part of
definition when reaction rate is defined in terms of a reactant because
reactant concentrations decrease a reaction proceeds; change in concentration
of a reactant is negative
¾
negative sign makes overall
rate positive
·
change in concentration of a
product positive; no negative sign when rate defined in respect to product
·
definition of rate with respect
to each reactant and product must reflect stoichiometric coefficients of
reaction
The Average Rate of the
Reaction
·
calculate average rate of
reaction using: –Δ [A]/Δt (M/s)
·
rate is average rate within
given time interval
Instantaneous Rate of
the Reaction
·
rate at any one point in time;
represented by instantaneous slope of curve at that point
·
can be determined from slope of
tangent to curve at point of interest
·
rate is same whether we use one
of reactants or product for calculation
·
generic reaction: aA + bB → cC
+ dD
rate of reaction: 
·
example on page 567
Measuring Reaction Rates
·
polarimetry:
measuring degree of lighting passing thorough a reacting solution
·
spectroscopy: most
common; light of specific wavelength passed through sample, intensity of
transmitted light (depends on how much light absorbed by sample) is measured
and recorded
·
reactions in which number of
moles of gaseous reactants and products changes as reaction proceeds can be
readily monitored by measuring changes in pressure
13.3 The Rate Law: The Effect of
Concentration on Reaction Rate (569-572)
·
rate law:
relationship between rate of reaction and concentration of reactants; rate = k [A]n
¾
k= rate constant:
constant of proportionality
¾
n= reaction order:
determines how rate depends on concentration of reactants
®
if n=0, reaction is zero order,
rate independent of concentration of A
®
if n=1, reaction is first order,
rate directly proportional to concentration of A
®
if n=2, reaction is second order,
rate proportional to square of concentration of A
®
other orders possible,
including non-integers, but these most common
Zero-Order Reaction
·
rate = k[A]0 = k
·
rate constant because reaction
doesn’t slow down as concentration of A decreases
·
rate same at any concentration
of A
·
occur under conditions where
amount of reactant available for reaction unaffected by changes in overall
quantity of reactant
First-Order Reaction
·
rate = k[A]1
·
rate slows down as reaction
proceeds because concentration of reactant decreases
·
rate directly proportional to
concentrations
Second-Order Reaction
·
rate = k[A]2
·
rate more sensitive to reactant
concentration
Determining the Order of
Reaction
·
order of reaction can be
determined only by experiment
·
method of initial rates: initial rate measured by running reaction several times with
different initial reactant concentrations to determine effect of concentration
on rate
·
can determine value of k by
solving rate law for k and substituting concentration and initial rate from any
one of the measurements
·
for 0 order reaction
concentration doubles as initial rate stays constant, for 1st order
reaction, concentration doubles as initial rate doubles, for 2nd
order concentration doubles as initial rate quadruples
·
·
rate constant for 0 order
reaction is Ms-1, 1st order reaction is s-1, 2nd
order reaction is M-1s-1
·
example on page 570
Reaction Order for Multiple
Reactants
·
consider reaction aA + bB → cC
+ dD
·
as long as reverse reaction
negligibly slow, rate = k [A]m [B]n
·
overall order: sum
of orders of all reactants in a chemical reaction
·
rate law for any reaction must
be determined experimentally
·
example of rate law and rate
order on page 572
13.4 The Integrated Rate Law: The
Dependence of Concentration on Time (573-580)
·
integrated rate law: relationship between concentrations of the reactants in a chemical
reaction and time
First-Order Integrated Rate Law
·
in A → products, rate directly
proportional to concentration
·
rate = k [A]
·
; aka differential rate law
; aka differential rate law
·
first order integrated rate
law: ln[A]t
= -kt + ln[A]0 or 
¾
[A]t = concentration
of A at any time, k= rate constant, [A]0= initial concentration of A
·
for 1st order
reaction, plot of natural log of reactant concentration as function of time
yields straight like with slope of –k and y-intercept of ln[A]0
·
example on page 575-576
Second-Order Integrated Rate Law
·
in A → products, rate
proportional to square of concentration of A
·
rate = k [A]2
·
·
second order integrated rate
law: 
·
for straight line, plot inverse
of concentration of reactant as function of time; slope of k, intercept of
1/[A]0
·
example on page 577
Zero Order Integrated Rate Law
·
rate proportional to constant
·
rate = k [A]0 = k
·
·
zero order integrated rate law:
[A]t
= -kt + [A]0
·
for straight line, plot
concentration of reactant as function of time; slope of k, intercept of [A]0
The Half-Life of a Reaction
·
half-life (t1/2): time required for concentration of reactant or amount of
radioactive isotope to fall to one-half of its initial value
·
half-life expression: defines dependence of half-life on rate constant and initial
concentration; different for different reactions
First-Order Reaction Half-Life
·
integrated rate law: 
·
half-life of first-order
reaction: 
·
t1/2 independent of
initial reaction
·
constant half-life unique to
first order reactions
·
example on page 579
Second-Order Reaction Half-Life
·
integrated rate law: 
·
half-life of second-order
reaction: 
·
half-life depends on initial
concentration
·
half-life continues to get
longer as concentration decreases
Zero-Order Reaction Half-Life
·
integrated rate law: [A]t
= -kt + [A]0
·
half-life of first order
reaction: 
·
half-life depends on initial
concentration
13.5 The Effect of Temperature on Reaction
Rate (581-588)
·
temperature dependence of
reaction rate contained in k
·
k only constant when
temperature constant
·
increase in temperature
generally results in increase in k,
resulting in faster rates
·
Arrhenius equation:
relates rate constant of reaction to temperature, activation energy, and
frequency factor
¾
; k= rate constant,
T=temperature (K)
; k= rate constant,
T=temperature (K)
¾
R= gas constant= 8.34 J/molK
¾
frequency factor (A): aka pre-exponential factor; number of times that reactants
approach activation energy per unit time
¾
activation energy (Ea): energy barrier in chemical reaction that must be overcome for
reactants to be converted into products
The Activation Energy
·
activated complex (transition state): high-energy intermediate state between reactant and product
·
activation energy is energy
required to reach activated complex
·
higher Ea, slower reaction
rate (at given temperature)
The Frequency Factor
·
approaching activation barrier
not equivalent to surmounting it
·
most approaches don’t have
enough total energy to make it over activation barrier
The Exponential Factor
·
exponential factor:
number from 0-1 that represents fraction of molecules that have enough energy
to pass activation barrier on given approach
·
fraction of approaches that are
successful and result in product
·
·
as temperature increases,
number of molecules having enough thermal energy to pass activation barrier
increases
Arrhenius Plots:
Experimental Measurements of the Frequency Factor and the Activation Energy
·
·
Arrhenius plot: plot
of natural log of rate constant (ln k) versus inverse of temperature in kelvins
(1/T) that yields a straight line with slope of –Ea/R and
y-intercept of ln A
·
example on page 584-585
·
when either data are limited or
plotting capabilities absent, can calculate activation energy if we know rate
constant at 2 different temperatures
·
2-point form of Arrhenius
equation: 
·
example on page 585-586
The Collision Model: A Closer Look at the
Frequency Factor
·
consider 2 gas phase reactants:
A(g) + B(g) → products
·
collision model:
model of chemical reactions in which reaction occurs after a sufficiently
energetic collision between two reactant molecules
·
each approach to activation
barrier is a collision between the reactant molecules
·
frequency factors of most
gas-phase chemical reactions tend to be smaller than number of collisions that
occur per second
·
¾
orientation factor (p): fraction of sufficiently energetic collisions
¾
collision frequency (z): number of collisions that occur per unit time; can be calculated
for a gas-phase reaction from pressure of gases and temperature of reaction
mixture
·
under typical conditions,
single molecule undergoes on the order of 109 collisions every
second
·
small orientation factor
indicates that orientational requirements for this reaction very stringent
·
reactions between individual
atoms usually have orientation factor of ~1 because atoms spherically symmetric
·
if orientation factor greater
than 1, collisions aren’t needed for reaction
·
harpoon mechanism:
positive charge on potassium and negative charge on bromine cause 2 species to
attract each other and form a bond without a collision
13.6 Reaction Mechanisms (588-591)
·
overall equation doesn’t show
intermediate steps
·
reaction mechanism:
series of individual chemical steps by which an overall chemical reaction
occurs
·
elementary step:
an individual step in a reaction mechanism; can’t be broken down into simpler
steps
·
individual steps in mechanism
add to overall reaction
·
reaction intermediates: species that are formed in step of a reaction mechanism and
consumed in another
·
can piece together a reaction
mechanism by measuring kinetics of overall reaction and working backward to
write a mechanism consistent with measured kinetics
Rate Laws for Elementary Steps
·
molecularity:
number of reactant particles involved in elementary step; unimolecular and
bimolecular most common
·
unimolecular:
describes a reaction that involves only one particle that goes on to form
products
·
bimolecular: an elementary step in a reaction that
involves two particles, either the same species or different, that collide and
go on to from products
·
termolecular: an
elementary step of a reaction in which three particles collide and go on to
form products; rare because probability of three particles colliding
simultaneously is small
·
rate law for overall chemical
reaction can’t be deduced from balanced chemical equation
·
rate law proportional to product of
concentration of particles in elementary step
Rate-Determining
Steps and Overall Reaction Rate Laws
·
rate-determining step: step in a reaction mechanism that occurs much slower than any of
the other steps’ limits overall rate of reaction, determining rate law for
overall reaction
·
for proposed reaction mechanism
to be valid:
¾
elementary steps in mechanism
must sum to overall reaction
¾
rate law predicted by mechanism
must be consistent with experimentally rate law
·
valid mechanism not proven
mechanism because other mechanisms may also fulfill both above requirements
Mechanism with a Fast Initial
Step
·
with slow initial step, rate
law predicted by mechanism normally contains only reactants involved in overall
reaction
·
when mechanism begins with fast
initial step, some other subsequent step in mechanism is rate limiting step
·
rate law predicted by rate
limiting step may contain reaction intermediates
·
rate law containing intermediates
cannot generally correspond to experimental law
·
products of first step can
build up, rate at which they’re consumed limited by slower step further down;
as products build up, can begin to react with one another to re-form reactants
·
as long as first step is fast
enough compared to rate-limiting stem, first-step reaction will reach
equilibrium
·
indicate equilibrium:
reactants products
·
if equilibrium reached, rate of
forward reaction equals rate of reverse reaction
·
example on page 592-593
13.7 Catalysis (593-598)
·
catalyst:
substance that is not consumed in a chemical reaction, but increases rate of
reaction by providing an alternate mechanism in with rate-determining step has a
smaller activation energy
·
activation energy for rate-limiting
step much smaller than for first, uncatalyzed pathway, so faster reaction
Homogenous and Heterogeneous
Catalysis
·
homogenous catalysis: catalysis in which the catalyst exists in same phase as reactants
·
heterogeneous catalysis: catalysis in which catalyst and reactants exist in different
phases
·
hydrogenation: the
catalyzed addition of hydrogen to alkene double bonds to make single bonds;
heterogeneous
·
ex. of hydrogenation: in
presence of finely divided platinum, palladium, or nickel, reaction between
ethene and hydrogen (relatively slow at normal temperature) happens rapidly
¾
1. adsorption reactants are
adsorbed onto metal surface
¾
2. diffusion: reactants diffuse
on the surface until they approach each other
¾
3. reaction: the reactants
react to form the products
¾
4. desorption: the products
desorb from the surface into the gas
¾
large activation energy of
hydrogenation reaction greatly lowered when reactants adsorb onto surface
Enzymes: Biological Catalysts
·
enzymes:
biochemical catalyst made of protein that increases rates of biochemical
reactions; usually large protein molecules with complex 3D structures
·
active site: the
specific area of an enzyme at which catalysis occurs; properties and shape just
right to bind reactant molecule
·
substrate: the
reactant molecule of a biochemical reaction that bonds to an enzyme at the
active site
¾
fits into active site like key
fits into lock
·
when substrate binds to active
site of enzyme, activation energy of reaction greatly lowered, allowing
reaction to occur at much faster rate
·
by allowing otherwise slow
reactions to occur at reasonable rates, enzymes give living organisms
tremendous control over which reactions occur, and when they occur
·
each enzyme very specific
(catalyzes only single reaction) and efficient
·
many substances that inhibit
action of enzymes are highly toxic
CHAPTER
14
14.2 The Concept of Dynamic Equilibrium
(615-617)
·
reversible: as
applied to a reaction, the ability to proceed in either the forward or reverse
direction
·
reversible reaction: a reaction that achieves the theoretical limit with respect to
free energy and will change direction upon an infinitesimally small change in a
variable related to the reaction
·
dynamic equilibrium: the point at which the rate of the reverse reaction or process
equals the rate of the forward reaction or process; forward and reverse
occurring at same rate
·
concentrations of reactants and
products no longer change at equilibrium, doesn’t meant concentration of
reactants and products are equal to one another at equilibrium
14.3 The Equilibrium Constant (K)
(618-622)
·
equilibrium constant (K): ratio, at equilibrium, of concentrations of products of a reaction
raised to their stoichiometric coefficients, to concentrations of reactants
raised to their stoichiometric coefficients
·
law of mass action:
relationship between balanced chemical equation and expression of equilibrium
constant
¾
products/ reactants
products/ reactants
Expressing Equilibrium Constants
for Chemical Reactions
·
to express equilibrium constant
for chemical reaction, apply law of mass action to balanced chemical equation
·
coefficients in chemical
equation become exponents in expression of equilibrium constant
·
example on page 618
The Significance of the
Equilibrium Constant
·
large equilibrium constant
(K>>1) indicates that numerator (number of products at equilibrium) is
larger than denominator (amount of reactants at equilibrium); reaction
favoured; forward reaction essentially completes
·
equilibrium constant only says
how far reaction has proceeded once equilibrium is reached, not how fast
·
(K<<1) indicates reverse
reaction favoured; high concentration of reactants, low concentrations of
products; forward reaction doesn’t proceed very far
·
K ≈ 1 means neither direction
favoured; forward reaction proceeds about halfway
Relationships between the
Equilibrium Constant and the Chemical Equation
·
if chemical equation modified
in some way, equilibrium constant for equation must be changed to reflect
modification
·
if you reverse equation, invert
equilibrium constant
¾
consider A + 2B ↔ 3C 
¾
reversed to 3C ↔ A + 2B 
·
if you multiply coefficients in
equation by a factor, raise equilibrium constant to same factor
¾
consider A + 2B ↔ 3C 
¾
multiply by n to get nA + 2nB ↔
3nC 
·
if you add 2 or more individual
chemical equations to obtain an overall equation, multiply corresponding equilibrium
constants by each other to obtain overall equilibrium constant
¾
consider A↔ 2B 
2B
↔ 3C 
2B
↔ 3C
¾
2 equations sum as follows
A↔ 2B
2B
↔ 3C 
<-
product of K1 and K2
<-
product of K1 and K2
A ↔ 3C
·
example on page 622
14.4 Expressing the Equilibirum Constant
in Terms of Pressure (622-624)
·
for gaseous reactions, partial
pressure of a gas is proportional to its concentrations of reactants and
products
·
can express equilibrium
constant in terms of partial pressures of reactants and products
·
consider 2SO3(g) ↔ 2SO2(g) +
O2(g) 
; KC= K with respect to
concentration in molarity
; KC= K with respect to
concentration in molarity
¾
KP= K in respect to partial
pressures in atmospheres; partial pressures of each
gas instead of conc.
®
·
since partial pressure of a gas
in atmospheres not same as concentration in molarity, value of KP
for a reaction necessarily equal to value of KC
·
[A]= concentration of
ideal gas, nA=moles of A, V= volume in litres
[A]= concentration of
ideal gas, nA=moles of A, V= volume in litres
·
ideal gas law: 
PA= partial pressure of A R= gas constant= 8.34 J/molK
PA= partial pressure of A R= gas constant= 8.34 J/molK
·
¾
if total number of moles of gas
is same before and after reaction, then Δn = 0, and KP equal to KC
·
example on page 624
Units of K
·
as long as concentration units
expressed in molarity for KC and pressure units expressed in
atmospheres for KP, we can enter quantities directly into
equilibrium expression, dropping their corresponding units
14.5 Heterogeneous Equilibria: Reactions
Involving Solids and Liquids (625-626)
·
concentration of a solid
doesn’t change because solid doesn’t expand to fill its containiner;
concentration depends only on density
·
pure solids and pure liquids
not included in equilibrium expression (KC equation)
·
example on page 626
14.6 Calculating the Equilibrium Constant
from Measured Equilibrium Concentrations (626-628)
·
most direct way to obtain
experimental value for equilibrium constant of a reaction is to measure
concentrations of reactants and products in a reaction mixture at equilibrium
·
equilibrium constant will
always be same at given temperature, regardless of initial concentrations
·
in most cases, only need to
know initial concentrations of reactant(s) and equilibrium concentration of any
one reactant or product; other equilibrium concentrations can be deduced from
stoichiometry of reaction
·
consider A(g) ↔ 2B(g)
;
¾
initial [A] is 1.00 M, initial
[B] is 0.00 M; when equilibrium reached, concentration of A is 0.75 M
¾
A changed by -0.25 M, so based
on stoichiometry, [B] changed by 2 x (+ 0.25 M) or +0.50 M
[A]
|
[B]
|
|
Initial
|
1.00
|
0.00
|
Change
|
-0.25
|
+0.50
|
Equilibrium
|
0.75
|
0.50
|
¾
to calculate equilibrium
constant, use balanced equation to write expression for equilibrium constant,
then substitute concentrations from the ICE table
¾
·
examples on page 628
14.7 The Reaction Quotient: Predicting the
Direction of Change (629-631)
·
reaction quotient (QC): ratio, at any point in the reaction, of concentrations of products
of a reaction raised to their stoichiometric coefficients to the concentrations
of the reactants raised their stoichiometric coefficients
·
for gases with amounts measured
in atmospheres, reaction quotient uses partial pressures in place of
concentrations, and is called QP
·
reaction quotient depends on
current state of reaction and has many different values as reaction proceeds
·
consider aA + bB ↔ cC + dD 
¾
in reaction mixture containing
only reactants, QC= 0
¾
in reaction mixture containing
only products, QC= 
¾
in
reaction mixture containing both reactants and products, each at a
concentration of 1 M, QC=1
·
value
of Q relative to K is measure of reaction toward equilibrium
·
at
equilibrium, reaction quotient is equal to equilibrium constant
·
Q
< K reaction goes to right (toward products)
Q > K reaction goes
to left (toward reactants)
Q = K reaction is at
equilibrium
·
example
on page 631
14.8 Finding Equilibrium Concentrations
(631-640)
Finding Equilibrium
Concentrations When We Know the Equilibrium Constant and All but One of the
Equilibrium Concentrations of the Reactants and Products
·
can use the equilibrium
constant to calculate equilibrium concentration of one of the reactants or
products, given the equilibrium concentrations of the others
·
example on page 632
Finding Equilibrium
Concentrations When We Know the Equilibrium Constant and Initial Concentrations
or Pressures
·
set up ICE chart, chances in concentration
represented by variable x
·
consider A(g) ↔ 2B(g)
[A]
|
[B]
|
|
Initial
|
1.0
|
0.00
|
Changed
|
-x
|
+ 2x
|
Equilibrium
|
1.0-x
|
2x
|
¾
solve for x using quadratic
equation
·
examples on page 634-635
·
when initial conditions given
in terms of partial pressures (instead of concentrations) and equilibrium,
constant given as KP instead of KC, use same procedure,
but substitute partial pressures for concentrations
·
examples on page 636-637
Simplifying Approximations in
Working Equilibrium Problems
·
if equilibrium constant relatively
small, reaction won’t proceed very far to right; if initial reactant
concentration relatively large, can assume x is small relative to initial
concentration of reactant;
·
if x much smaller than 1, then
denominator can be approximated as 1.0
·
consider above example (in ICE
chart), this time, K= 3.3 x 10-5
·
to check validity: ratio of s
to number it’s subtracted from should be <0.05
·
examples page 638-640
14.9 Le Châtelier’s Principle: How a
System at Equilibrium Responds to Disturbances (641-648)
·
Le Châtelier’s Principle: when a chemical system at equilibrium is disturbed, the system
shifts in a direction that minimizes the disturbance
The Effect of a Concentration
Change on Equilibrium
·
increasing concentration of one
or more of the reactants (making Q<K) shifts reaction to right
·
increasing concentration of one
or more of the products (making Q>K) shift reaction left
·
decreasing concentration of one
or more of the reactants (making Q>K) shifts reaction left
·
decreasing concentration of one
or more of the products (making Q<K) shifts reaction right
·
example on page 644
The Effect of a Volume (or
Pressure) Change on Equilibrium
·
decrease in volume causes an
increase in pressure, vice versa
·
decreasing volume causes
reaction to shift in direction that has fewer moles of gas particles
·
increasing volume causes
reaction to shift in direction that has greater number of moles of gas
particles
·
if reaction has equal number of
moles of gas on both sides of chemical equation, change in volume produces no
effect on equilibrium
·
adding an inert gas to mixture
at a fixed volume has no effect on equilibrium
·
example on page 645
The Effect of a Temperature
Change on Equilibrium
·
changing temperature changes value
of equilibrium constant
·
increasing temperature causes
exothermic reaction to shift left; equilibrium constant (K) decreases
·
decreasing temperature causes
exothermic reaction to shift right; K increases
·
increasing temperature causes
endothermic reaction to shift right; K increases
·
decreasing temperature causes
endothermic reaction to shift left; K decreases
example
on pa
