unit 10 Flashcards
1
Q
general properties of acids
A
sour taste
able to dissolve many metals and neutralize bases
2
Q
general properties of bases
A
bitter taste
slippery feel
ability to neutralize acids
3
Q
Arrhenius definition of acid
A
a substance that produces H+ ions in solution
4
Q
arrhenius definition of base
A
substance that produces OH- ions in solution
5
Q
acids and bases combine to form
A
H2O and a salt
6
Q
Bronsted-Lowry definition of Acids and Bases
A
acids (proton donors) and bases (proton acceptor s) always occur together in an acid-base reaction
7
Q
conjugate acid-base pair
A
2 substances related to each other by the transfer of a proton
8
Q
conjugate acid
A
any base to which a proton has been added
9
Q
conjugate base
A
any acid from which a proton has been removed
10
Q
strength of acids and bases proportion
A
inversely proportional
(stronger acid, weaker base and vice versa)
11
Q
modern definition of acid
A
any compound having one or more hydrogen atoms that are weakly bound to the rest of the molecule. when dissolved in water, H+ ions ionize from the rest of the molecule.
12
Q
what is the first element usually written in an acid compound?
A
H
13
Q
what happens to organic acids when they ionize?
A
-COOH —-} -COO- + H+
14
Q
amphoteric
A
capable of reacting as an acid or base
15
Q
binary acids
A
contain H and one other atom
16
Q
in aq solns, the name of every binary acid starts with _ and ends with _
A
hydro-, -ic
17
Q
HF
A
hydrofluoric acid
18
Q
HCl
A
hydrochloric acid
19
Q
HBr
A
hydrobromic acid
20
Q
HI
A
hydroiodic acid
21
Q
H2S
A
hydrosulfuric acid
22
Q
polyatomic anion acid rules
A
-ending of polyatomic anion name is changed, word acid is added
-if the polyatomic anion ends in -ate, it’s changed to -ic
-if the ending is -ite, its changed to -ous
23
Q
sulfate changes to
A
sulfuric acid
24
Q
sulfite changes to
A
sulfurous acid
25
nitrate changes to
nitric acid
26
nitrite changes to
nitrous acid
27
hypochlorite changes to
hypochlorous acid
28
chlorite changes to
chlorous acid
29
chlorate changes to
chloric acid
30
perchlorate changes to
perchloric acid
31
hydroxide bases
have hydroxide ions in their formulas, named using ionic compound nomenclature
32
NaOH name
sodium hydroxide
33
nitrogen bases definition
related to NH3 are amines
34
acid or base? HNO3
acid
35
acid or base? NH4+
acid
36
acid or base? KOH
base
37
acid or base? HC2H3O2
acid
38
strong acid
-equilibrium lies far to right
-acid ionizes completely
39
monoprotic acid
acids with only one ionizable proton
ex. HCl, HBr, HI
40
diprotic acid
acids with 2 ionizable protons
ex. H2SO4
41
weak acid
equilibrium lies far to left
-only small percentage of acid molecules ionize
42
nitric acid
HNO3
43
sulfuric acid
H2SO4
44
hydrofluoric acid
HF
45
acetic acid
HC2H3O2
46
sulfurous acid
H2SO3
47
carbonic acid
H2CO3
48
phosphoric acid
H3PO4
49
acid ionization constant
equilibrium constant for the ionization reaction of a weak acid
50
autoionization
water is amphoteric and can act as a base or acid with itself
51
autoionization reaction
H2O ----> H+ + OH-
52
ion product constant expression or dissociation constant (Kw)
ionization constant for water
[H3O+] x [OH-] = 1.0x10^-12
53
acidic solution
contains acid that increases [H3O+]
[OH-] decreases (= 1.0x10^-11 M)
[H3O+] > [OH-]
54
basic solution
contains a base that increases [OH-]
[H30+] decreases (= 1.-x10^-12 M)
[OH-] > [H3O+]
55
pH scale definition
specifies acidity of a solution
56
pH =
- log [H30+]
57
pH < 7
acidic
58
pH > 7
basic
59
pH = 7
neutral
60
change of 1 pH unit corresponds to
10-fold change in [H30+]
61
if the [H30+] has two sig figs, report answer to
2 decimal points
62
pOH scale definition
defined with respect to [OH-]
63
pOH =
- log [OH-]
64
pOH < 7
basic
65
pOH > 7
acidic
66
pH + pOH =
14.00
67
pK
way of expressing strength of acid or base
68
pKa =
-log (Ka)
69
pKb =
-log (Kb)
70
Ka =
10^-pKa
71
Kb=
10^-pKb
72
acid strength and pKa relation
stronger acid = smaller pKa
73
base strength and pKb relation
stronger base = smaller pKb
74
Ka and pKa relation
larger Ka = smaller pKa
because it is the -log
75
Kb and pKb relation
larger Kb = smaller pKb
76
what are the sources of H3O+ in a solution containing a strong or weak acid?
ionization of the acid
autoionization of H2O
77
how much [H3O+] is produced during the autoionization of water at 25 degrees C? what happens to the additional [H3O+]?
1.0 x 10^-7
additional [H3O+] from acid shifts equilibrium to the left. autoionization of H2O produces less H3O+ than in pure water and can be ignored
78
strong acids
completely ionize in solution
79
weak acids
does not completely ionize in solution
80
percent ionization =
[ionized acid] / initial [acid] x100%
81
equilibrium [H30+] of a weak acid _ with increasing initial [acid]
increases
82
percent ionization of a weak acid _ with increasing [acid]
decreases
83
group 1A metal hydroxides are _ soluble in water and form
highly, strongly basic solutions
84
group 2A metal hydroxides are _____ in water and produce
slightly soluble, 2 mol of OH- per mole of base
85
group 2A metal hydroxides have the general formula
M(OH)2
86
base ionization constant (Kb)
quantized extent of ionization of a weak base
87
pKb =
-log (Kb)
88
generally, anions form _ solutions
basic or neutral
89
generally, cations form _ solutions
acidic or neutral
90
an anion that is the conjugate base of a strong acid is
pH - neutral
91
an anion that is the conjugate base of a weak acid is
a weak base
92
general acid and conjugate base relation
weaker acid = stronger conjugate base
93
3 categories of cations acting as weak acids
-cations that are the counterions of strong bases
-cations that are the conjugate bases of weak acids
-cations that are small, highly charged metals
94
polyprotic acid
contains 2+ ionizable protons
-typically ionize in successive steps, each with its own Ka
95
is Ka1 or Ka2 smaller? why
Ka2, first H+ separates from a neutral molecule while the second separates from an anion
96
in which Ka step is the most amount of [H3O+] formed? why
first step
formation of H3O+ in the first step inhibits formation of additional H3O+ in subsequent steps
97
strength of an acid and strength of the bond between H and molecule proportion
inversely
98
bond polarity
for a binary acid with bond H-Y, the H atom should be the positive pole for the H-Y bond to make acid acidic
partial positive charge on H makes it easier for H to separate and become an ion
99
bond strength
strength of H-Y bond on a binary acid affects strength of acid
stronger bond (higher electronegativity on Y) = weaker acid
100
combined effect of bond polarity and strength trends on PT
from left to right on PT, hydrides become more acidic (H-Y bond becomes more polar an EN increases)
from top to bottom, hydrides become more acidic (as we go down a group, anion size increases so bond length between anion-H increases, weakens bond)
101
oxyacids contain
H, O, and other element
102
oxyacids bond structure
O atoms bonded to central atom, H atoms bonded to O atoms
103
bond strength in oxyacids depends on
-electronegativity of central atom
-number of O atoms bonded to the central atom
-number of O atoms per H atom
104
electronegativity and acid strength proportion in oxyacids. explain
direct
explanation: the more electronegative the central atom is, the more it weakens and polarizes the H-O bond and the more acidic the oxyacid
105
number of O atoms and acid strength proportion. explain
direct
explanation: because each O atom bonded to the central atom is electronegative, they draw electron density away from the central atom and from the O-H bond, further weakening and polarizing it, leading it to increased acidity
106
which is stronger, an acid with more O atoms per H atoms or an acid with more H atoms per O atoms? why?
an acid with more O atoms per H atoms
explanation: increasing electronegativities of central atom cause increase in acid strength. decreasing number of H atoms increases acid strength.
107
organic acids
contains carboxyl group -COOH
electronegative atoms (F, Cl, Br, I, O, S) will pull electron density from the O-H bond and increase the strength of the acid
108
buffers
resists pH change by neutralizing added acid or added base
109
buffers contain either:
Significant amounts of a weak acid and its conjugate base
Significant amounts of a weak base and its conjugate acid
110
how do buffers work?
When additional base is added to a buffer, the weak acid neutralizes the added base. When additional acid is added to a buffer, the conjugate base neutralizes the added acid
111
common ion effect
solution contains two substances that share a common ion
112
In any buffer where the acid and conjugate base concentrations are equal, [H3O+] =
Ka
113
Henderson-Hasselbalch Equation
pH = pKa + log ( [base] / [acid] )
Where the base is the conjugate base of the acid or the acid is the conjugate acid of the base
114
what does the Henderson-Hasselbalch Equation allow us to calculate?
pH of a buffer solution from the initial concentrations of the buffer components as long as the x is small approximation is valid
115
What is the pH of a buffer solution when the concentrations of both buffer solutions (weak acid and its conjugate base) are equal?
pH = pKa
116
What happens to the pH when the buffer contains more of the weak acid than the conjugate base?
pH decreases
117
What happens to the pH when the buffer contains more of the conjugate base than the weak acid?
pH increases
118
Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong acid is added to the buffer?
Some of the conjugate base is used to neutralize the acid. So the [weak acid] > [conjugate base]
119
Suppose that a buffer contains equal amounts of a weak acid and its conjugate base. What happens to the relative amounts of the weak acid and conjugate base when a small amount of strong base is added to the buffer?
Some of the weak acid is used to neutralize the added base so [weak acid] < [conjugate base]
120
how can you find the pH of a solution composed of a base and its conjugate acid?
calculate pKa for the conjugate acid of the weak base by subtracting pKb of the weak base from 14
121
How do you use the Henderson-Hasselbach equation to calculate the pH of a buffer containing a base and its conjugate acid? Specifically, how do you determine the correct value for pKa?
Identify which buffer component is the acid and which component is the base
If the buffer is a weak acid and its conjugate base, find Ka (tabulated), calculate pKa and use the Henderson-Hasselbalch to calculate the pH of the buffer
If the buffer is a weak base and its conjugate acid, find Kb (tabulated), calculate pKb and subtract pKb from 14 to get pKa, then use the Henderson-Hasselbalch equation to calculate the pH of the buffer.
122
2 factors influence the effectiveness of a buffer:
-relative amounts of the acid and conjugate base
-absolute concentrations of the acid and conjugate base
123
when is a buffer effective?
-the relative [acid] and [base] should not differ by more than a factor of 10
-The effective range for a buffering system is one pH unit on either side of the pKa
-when [acid] and [base] are high
most effective when [acid] = [conjugate base]
124
buffering capacity
amount of acid or base a buffer can neutralize
125
buffering range
pH range over which the buffer can be effective
126
a concentrated buffer can neutralize (more/less) added acid or base than a dilute buffer
more
127
acid-base titration:
a titrant is slowly added to a solution of known concentration from a biuret until the reaction is complete
128
titrant
solution of unknown concentration used in a titration
129
indicator definition
chemical that changes color when the pH changes
130
how do we know when we reach the endpoint of a titration?
an indicator will change color
131
equivalence point of a titration
moles H30+ = moles OH-
equal concentrations of buffer components
pH = pKa
all acid has been converted to its conjugate base (mol acid = mol base)
132
what's a titration curve?
plot of pH vs. amount of added titrant
133
what's the significance of the inflection point on a titration curve?
its the equivalence point of the titration
134
prior to equivalence point:
known solution in the flask is in excess
pH is closest to pH of the known solution
135
beyond the equivalence point:
unknown solution (titrant) added from the biuret (and OH-?) is in excess
pH approaches pH of titrant (increases)
136
when will the titration curve be decreasing?
when the acid is in the biuret and the base is in the flask
137
in a titration, initial pH =
pH of the weak acid solution
138
how do you calculate the pH in a titration problem?
solve an equilibrium problem using [weak acid] as the initial concentration
139
when does the solution become a buffer in a titration?
between the initial pH and the equivalence point
140
buffer region
between initial pH and equivalence point (both acid and its conjugate base are present)
141
main difference from the titration of a strong acid and weak acid
equivalence point of a titration of a strong acid is neutral
equivalence point of a titration of a weak acid is basic
142
main difference between titrating a weak acid and a strong base compared to other titrations
curve starts basic and has an acidic equivalence point
143
when a diprotic acid is titrated with a strong base:
-pH curve has two equivalence points if Ka1 and Ka2 are sufficiently different
-volume of base to reach first equivalence point = volume to reach second equivalence point
144
In the titration of a strong acid with a strong base, how would you calculate initial pH?
–log [acid]
145
In the titration of a strong acid with a strong base, how would you calculate pH at the equivalence point?
For the titration of a strong acid with a strong base, pH at the equivalence point is 7.00
146
The pH at the equivalence point of the titration of a strong acid with a strong base is 7.0. However, the pH at the equivalence point of a weak acid with a strong base is above 7.0. Explain
The pH at the equivalence point for titration of a strong acid is 7.0 because [H3O1+] = [OH1-]
In the titration of a weak acid, the acid is converted to its conjugate base. The increase in [conjugate base] raises the pH above 7.00
147
The volume required to reach the equivalence point of an acid-base titration depends on the volume and concentration of the acid or base to be titrated and on the concentration of the acid or base used to do the titration. It does not, however, depend on whether or not the acid or base being titrated is strong or weak. Explain
At the equivalence point mol acid = mol base regardless if the acid or base being titrated is strong or weak.
Calculation of mol acid or base requires the concentration and the volume of the solution
148
In the titration of a weak acid with a strong base, how would you calculate initial pH?
Use an equilibrium problem (ICE table) for the ionization of the weak acid
149
In the titration of a weak acid with a strong base, how would you calculate pH at one half of the equivalence point?
pH = pKa
150
indicators are _ acids
weak
151
indicators establish an equilibrium with
H2O and H3O+ in the solution
152
color of a titration solution depends on the relative concentrations of:
indicator : H indicator +
153
when pH = pKa (indicator)
[ln-]/[Hln] = 10^0 = 1
intermediate color
154
when pH = pKa + 1 (indicator)
[ln-]/[Hln] = 10^1 = 10
color of ln-
155
when pH = pKa - 1 (indicator)
[ln-]/[Hln] = 10^-1 = 0.10
color of Hln
156
at equivalence point. pH almost equals
pKa of H indicator+
157
What is the endpoint of an indicator and what is it used for?
An indicator changes color to indicate the equivalence point of a titration
158
An indicator is often a weak acid mixed with its conjugate base. How is the acid different from its conjugate base?
Acid form is a different color than its conjugate base
159
Why is it important to use small amounts of indicator?
The indicator should have no effect on the pH or equivalence point of the neutralization reaction
