Unit 3 (ch 15,17) Flashcards
1
Q
K»>1
A
favor product
2
Q
K«<1
A
favor reactant
3
Q
Does K have units?
A
no
4
Q
K positive or negative?
A
positive
5
Q
Kc=
A
[products]/[reactants]
6
Q
Kp=
A
Pproducts/Preactants
7
Q
Kc=Kp when
A
same number moles of gases
8
Q
Kp=
A
Kc(RT)^delta n
9
Q
K for the forward reaction = (blank) for the reverse
A
1/K
10
Q
K overall =
A
K1xK2
11
Q
If K is always at equilibrium, Q can be
A
any concentration
12
Q
In non-equilibrium conditions find
A
Q
13
Q
Q tells us
A
where we are relative to equilibrium and in which direction the reaction will proceed
14
Q
Q
A
reaction to right/products/forward
15
Q
If you go from a normal reaction to half the reaction, what process do you use?
A
K^1/2
16
Q
Q=K
A
equilibrium
17
Q
Q>K
A
reaction to left/reactants/reverse
18
Q
Heterogeneous equilibrium
A
solid and liquid excluded
19
Q
LeChatelier’s principle
A
stressed/preturbed reactions work to regain equilibrium
20
Q
what is LC useful for?
A
drive reaction to form more products
21
Q
Effect of pressure change
A
increase pressure favors side with fewer moles, decreasing favors side with more moles
22
Q
increasing temp
A
help endo move forward and hinder exo
23
Q
What do catalysts change?
A
time to reach equilibrium, not K
24
Q
How can we determine how much K changes with T?
A
ln(K2/K1)= delta H/R(1/T2-1/T1)
25
if you have a strong acid or base, the concentration of the acid or base will be
same as concentration of ions it dissociates into
26
is water included in Ka expressions?
no
27
What is equilibrium?
rate forward= rate reverse
28
What happens to concentrations and reaction rates when equilibrium is reached?
concentrations of reactants and products remain constant over time and the rate of a reaction in the forward matches reverse
29
How do we arrive at the relationship for the equilibrium constant, K?
Rate law forward/rate law reverse
30
What information does K>>>1 or K<<<1 give us about the ratio of products and reactants?
K>>1, concentrations of products larger than reactant
31
What are the effects of volume changes?
-increase volume shifts equilibrium toward side of reaction with moles of gas
32
How can we relate K to enthalpy, entropy, and free energy?
value of K decreases as temp increase for exo
33
What are acids according to the Bronsted-Lowry definition?
acids donate protons (H+)
34
What are bases according to the Bronsted-Lowry definition?
bases accept protons (H+)
35
What is the difference between a strong and weak acid?
strong acid fully dissociates
36
What is the difference between a strong and weak base?
strong base fully dissociates
37
How do we calculate Ka for an acid?
[A-][H3O+]/[HA]
38
Kb=
[HB+][OH-]/[B]
39
strong base kb
>>1
40
fewer moles in products, (blank) pressure
increase-favors products
41
Q
reaction proceeds forward
42
Q=K
reaction at equilibrium
43
Q>K
reaction proceeds reverse
44
remove reactant =
slower
45
increase partial pressure of reactant/product shifts equilibrium so
more substance consumed
46
decrease in partial pressure shifts toward
products
47
decrease volume shifts equilibrium toward
side of reaction with fewer moles of gas
48
is water included in ka expressions?
no
49
pH=
-log[H+]
50
pOH=
-log[OH-}
51
pH+pOH=
14.00
52
polyprotic acids
have more than 1 ionizable proton
53
protons and polyprotic acids
it gets harder to pull of protons from these (1st easiest) and sometimes second or third can be hard enough to remove that we can ignore it
54
acids and bases are weak or strong because of
structure
55
if conjugate is stable, then acid or base will be more likely to (blank)
give up or accept proton/H+ (resonance and anions with more resonance structures= more stable)
56
if we dissolve salt what happens
change pH
57
what happens if the salt produces the conjugate base/acid of a weak acid/base
it can react with water to produce OH- or H3O+ and change pH
58
what can be used to relate pKa of a proton to pH of the solution
pKa+log(base/acid)
59
pKa=
-log(Ka) and the point where [base]=[acid]
60
when [base]=[acid] what does it mean?
half the acid is deprotonated
61
pKa+pKb=
14.00
62
what is the Henderson-Hasselbalch equation used for
figure out how much of an acid and base we need to give a particular pH or how much pH will change when we add a strong acid or base to our buffered solution
63
how do buffers work?
weak acid/base reacting with the strong base/acid you're adding to neutralize them and limit the pH change
64
indicators show
pH change happening in system
65
when do indicators work best?
within +/- 1 pH unit of pKa
66
Why do indicators change color?
near pKa, losing or gaining a proton which changes their structure and electronic properties, which effects what wavelengths of light they absorb
67
indicators used in titration how?
look for when we've added enough acid or base to neutralize the other and determine concentration of unknown
68
determine concentration of unknown in titration problem
MaVa=MbVb
69
Ksp
solubility product constant- measure of how much of a solid will dissolve- usually just products
70
Kw is specific to
water
71
Kw always=
1.00x10^-14
72
pH and autoionization of H2O
H2O(l) base+H2O(l) acid--> H3O+ (aq) conjugate acid+ OH- (aq) conjugate base
73
pH=
potential of hydrogen ion
74
pH=
-log[H+]
75
[ ] =
10^-pH
76
sig figs with pH
first # not considered significant
77
if we have a weak and strong acid at some [ ], which is more acidic/has lower pH?
strong acid
78
pOH=
-log[OH-]
79
pH + pOH
14, pKw, -log[1x10^-14]
80
strong acid
complete dissociation, no ice, no K, no reverse
81
strong base
complete dissociation, no K, no ice, no reverse
82
weak acid
partly dissociates
| HA---<> H+ + A-, Ka= [H+][A-]/[HA]
83
the larger Ka is, the stronger
weak acid is- wants to donate H+
84
weak bases
partially accept an H+ from water
B+H2O(l)-->< HB+ + OH-
Kb= [OH-]{HB+]/[B]
85
The larger Kb is, the stronger
the weak base, the more it wants to accept a H+ from water
86
[H+][OH-]=
Kw
87
Ka=
kw/kb
88
ka x kb=
kw (1,00x10^-14)
89
larger Ka is, smaller (blank)
kb for conjugate base
90
weakest weak bases include
Cl-, Br-, I-, NO2-, VSO4-, ClO4-
91
pKa=
-log[ka]
92
pkb=
-log[kb]
93
small pka means
larger Ka, strong acid
94
small pkb means
larger kb, strong base
95
polyprotic
acids with more than 1 H+ to donate, bases which will accept more than 1 H+
96
oxoacids
atom (usually nonmetal) bonded to 1 or more O with some H+ as well
97
in a series of oxoacids, what happens?
acid strength goes up with more O
98
examples of organic amines
nitrogen, ammonium
99
organic amines
ammonia type areas of organic (c) molecules
| N surrounded by 3 other atoms/ R groups with a lone pair (will accept H+)
100
anions tend to be
bases
101
Buffers
resist change in pH, mixture of a weak acid and its conjugate base
102
weak acids
eat OH-
103
weak bases
eat H+
104
Henderson-Hassel Balch equation pKa
-log[H+]+-log([A-]/[HA])
105
Henderson-Hassel Balch equation pH
pKa + log [A-]/[HA]
106
common ion has to do with
LeP and stress on system
107
ksp helps determine
how soluble a compound is
108
ksp=
[M^n+]^M [Z^y-]^Z
109
pH indicators work by
being ionized (gain or lose H+)
110
How can we use pH and pKa to find the acid base ratio?
pH=pKa + log (base/acid)
111
acid base titration
pH refers to a solution while pKa refers to a specific molecule/H+ on a molecule
112
[base]=[acid]
pH = pKa
113
equivalence point for strong acid/strong base
7.00
114
equivalence point for strong acid/weak base
pH < 7.00
