Test 2 Flashcards
1
Q
chemical kinetics
A
studies the speed and mechanisms of chemical reactions
2
Q
reaction rate
A
measures the speed of a chemical reaction, defined as the change in concentration over time
3
Q
rate
A
can be expressed as increase in product concentration or decrease in reaction concentration per unit time
4
Q
average rate
A
rate measured over time window
5
Q
are rates constant?
A
no, they typically decreasee over time
6
Q
instantaneous rate
A
found by calculating slope of tangent to concentration-time curve
7
Q
rate law
A
shows a relationship between reaction rate and reactant concentrations
8
Q
rate equation
A
rate= k[A]^m[B]^n
k = rate constant
m,n = reaction orders for respective reactants
9
Q
overall order
A
sum of individual orders
10
Q
how are reaction orders determined
A
experimentally
11
Q
determining initial rates
A
- vary concentration of one reactant while keeping others constant
- measure initial rates of each reaction
- compare rates to determine order
- solve system of rate law equations
12
Q
graphical/integrated rate law
A
- monitor concentrations over time
- plot data according to integrated rate laws
- shape of curve reveals reaction order
13
Q
zero order reaction
A
- linear plot of [A] vs time
- half life depends on initial concentration
- gets shorter over reaction course
14
Q
first order reaction
A
- linear plot of ln[A] vs time
- half life independant of concentration
t1/2=ln(2)/k
15
Q
second order reaction
A
- linear plot of 1/[A] vs time
- half life = 1/(k[A]0)
16
Q
what is slowest step
A
rate-determining
17
Q
what are intermediates
A
formed in one step and consumed in a later one
18
Q
concentration
A
higher concentration increases collision theory
19
Q
temperature
A
increases molecular motion and collision energy, follows arrhenius equation
20
Q
arrhenius equation
A
ln(k)=-Ea/RT+ln(A)
- used t find activation energy from temperature-dependent rates
21
Q
catalysts
A
- provides alternative reaction pathway with lower activation energy
- appear in rate law but regenerated in reaction
- doesn’t change overall reaction thermodynamics
- enzymes are biological catalysts showing concentration dependent behaviour
22
Q
collision theory
A
- successful reaction requires: proper molecular orientation, energy greater than activation energy
- only a fraction of collisions lead to reactions
- temperature increase exponentially increases successful collision fraction
k=Zpf
23
Q
orientation factor
A
p is usually <1
24
Q
chemical equilibrium
A
- all chemical reactions are reversible to some degree
- concentrations of reactants and products remain constant
- forward and reverse reactions occur at equal rates
- individual atoms continuously switch between products and reactants
25
Kc
represents the ratio of product to reactant concentrations at equilibrium
Kc = products/reactants
26
properties of equilibrium constants
- unitless quantities
- can use pressures (Kp) instead of concentrations
Kp= Kc(RT)(delta)n
where deltan is change in moles of gas
27
Interpretation of Kc values
Kc>103 : products favoured
Kc<10-3 : reactant favoured
10-3>Kc>103 : significant amounts of both
28
reaction quotient
similar to Kc but uses non-equilibrium concentrations, predicts reaction direction
Qc = Kc : at equilibrium
Qc < Kc : reaction proceeds forwards
Qc > Kc : reaction proceeds backwards
29
ice method
- initial change equilibrium
- systematic approach to calculate equilibrium concentrations
steps:
1. list initial concentrations
2. define changes in terms of x
3. write equilibrium expressions
4. solve for x
5. calculate final concentrations
30
le châtelier's principle
- explains system response disturbances
- concentration changes, pressure/volume changes, temperature changes
31
concentration changes
- added substance, reaction shifts to consume it
- removed substance, reaction shifts to produce it
32
pressure/volume changes
- decreased volume, shifts toward fewer gas molecules
- increased volume, shifts toward more gas molecules
- inert gas addition, no effect
33
temperature changes
- exothermic reactions, Kc decreases with increasing temperature
- endothermic reactions, Kc increases with increasing temperature
34
brønsted-lowry theory
- acids are proton donors
- bases are proton acceptors
- differs from Arrhenius theory which only considered H+ production (acids) and OH- production (bases) in water
- water is amphiprotic/amphoteric - can act as both acid and base
35
what does water undergo
- self ionization
H2O + H2O ⇌ H3O+ +OH-
36
pH equation
pH=-log[H3O+]
37
pOH equation
pOH=-log[OH-]
38
pOH + pH
14 at 25 °C
39
neutral solution
pH of 7
40
acidic solution
pH smaller than 7
41
basic solution
pH larger than 7
42
strong acids and bases
- completely dissociate in water
- strong acids: HClO4, H2SO4, HNO3, HCl, HBr, HI
- strong bases: group 1 and 2 hydroxides
43
weak acids and bases
- partially dissociate in water
- characterized by Ka (acid dissociation constant) or Kb (base dissociation constant)
44
equation for conjugate acid-base pairs
Ka x Kb = Kw
45
pKa equation
pKa = -log Ka
46
pKb equation
pKb = -log Kb
47
stronger acid effect on Ka and pKa
larger Ka, smaller pKa
48
Factors affecting acid strength
1. bond polarity: more polar H-A bonds create stronger acids, increases with electronegativity of A
2. bond strength: weaker H-A bonds create stronger acids, bond strength decreases down a group (larger atoms)
3. oxoacids (YOm(OH)n): electronegativity of central atom Y, number of oxygen atoms not bonded to hydrogen
49
salt solutions
- can be acidic, basic, or neutral depending on parent acid/base
- strong acid + strong base = neutral
- weak acid + strong base = basic
- strong acid + weak base = acidic
50
Lewis base
electron pair donor
51
lewis acid
electron pair acceptor
52
common ion effect
shift in equillibrium position when adding a substance containing an ion already present
- ex. adding acetate ions suppresses dissociation of acetic acid
53
what do buffers do
- resist pH changes when limited amounts of acid/base are added
54
buffer components
weak acid + conjugate base in appreciable concentrations
55
buffer capacity
- amount of acid/base buffer can handle without significant pH change
56
effective buffers
maintain pH within ±1 of pKa
57
henderson-hasselbalch equation
pH =pKa + log([base-]/[acid])
- used for calculating buffer pH or determining buffer compositio
58
how to make buffers
1. select weak acid with pKa near desired pH
2. mix equal amounts of acid and conjugate base OR
3. neutralize half of weak acid with strong base
4. adjust to desired pH
59
titrations
- used to determine solution concentration using standard solution
- strong acid-strong base: pH 7 at equivalence point
- weak acid-stron base: pH larger than 7 at equivalence point
- strong acid-weak base: pH smaller than 7 at equivalence point
60
titration regions
1. initial pH determined by original solution
2. buffer region forms during titration
3. equivalence point when stoichiometrically equal amounts mixed
4. post-equivalence determined by excess titrant
61
solubility produce (Ksp)
- equilibrium constant for dissolution
62
factors affecting solubility
ph changes, common ions, temperature
63
precipitation predictions
Compare ion product Q with Ksp
Q>Ksp : precipitation occurs
Q=Ksp : equilibrium
Q < Ksp : dissolution occurs
