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ap chem > unit 10 > Flashcards

unit 10 Flashcards

1
Q

general properties of acids

A

sour taste
able to dissolve many metals and neutralize bases

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2
Q

general properties of bases

A

bitter taste
slippery feel
ability to neutralize acids

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3
Q

Arrhenius definition of acid

A

a substance that produces H+ ions in solution

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4
Q

arrhenius definition of base

A

substance that produces OH- ions in solution

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5
Q

acids and bases combine to form

A

H2O and a salt

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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

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7
Q

conjugate acid-base pair

A

2 substances related to each other by the transfer of a proton

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8
Q

conjugate acid

A

any base to which a proton has been added

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9
Q

conjugate base

A

any acid from which a proton has been removed

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10
Q

strength of acids and bases proportion

A

inversely proportional

(stronger acid, weaker base and vice versa)

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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.

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12
Q

what is the first element usually written in an acid compound?

A

H

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13
Q

what happens to organic acids when they ionize?

A

-COOH —-} -COO- + H+

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14
Q

amphoteric

A

capable of reacting as an acid or base

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15
Q

binary acids

A

contain H and one other atom

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16
Q

in aq solns, the name of every binary acid starts with _ and ends with _

A

hydro-, -ic

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17
Q

HF

A

hydrofluoric acid

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18
Q

HCl

A

hydrochloric acid

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19
Q

HBr

A

hydrobromic acid

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20
Q

HI

A

hydroiodic acid

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21
Q

H2S

A

hydrosulfuric acid

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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

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23
Q

sulfate changes to

A

sulfuric acid

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24
Q

sulfite changes to

A

sulfurous acid

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25
Q

nitrate changes to

A

nitric acid

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26
Q

nitrite changes to

A

nitrous acid

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27
Q

hypochlorite changes to

A

hypochlorous acid

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28
Q

chlorite changes to

A

chlorous acid

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29
Q

chlorate changes to

A

chloric acid

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30
Q

perchlorate changes to

A

perchloric acid

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31
Q

hydroxide bases

A

have hydroxide ions in their formulas, named using ionic compound nomenclature

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32
Q

NaOH name

A

sodium hydroxide

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33
Q

nitrogen bases definition

A

related to NH3 are amines

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34
Q

acid or base? HNO3

A

acid

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35
Q

acid or base? NH4+

A

acid

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36
Q

acid or base? KOH

A

base

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37
Q

acid or base? HC2H3O2

A

acid

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38
Q

strong acid

A

-equilibrium lies far to right
-acid ionizes completely

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39
Q

monoprotic acid

A

acids with only one ionizable proton

ex. HCl, HBr, HI

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40
Q

diprotic acid

A

acids with 2 ionizable protons

ex. H2SO4

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41
Q

weak acid

A

equilibrium lies far to left
-only small percentage of acid molecules ionize

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42
Q

nitric acid

A

HNO3

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43
Q

sulfuric acid

A

H2SO4

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44
Q

hydrofluoric acid

A

HF

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45
Q

acetic acid

A

HC2H3O2

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46
Q

sulfurous acid

A

H2SO3

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47
Q

carbonic acid

A

H2CO3

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48
Q

phosphoric acid

A

H3PO4

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49
Q

acid ionization constant

A

equilibrium constant for the ionization reaction of a weak acid

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50
Q

autoionization

A

water is amphoteric and can act as a base or acid with itself

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51
Q

autoionization reaction

A

H2O —-> H+ + OH-

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52
Q

ion product constant expression or dissociation constant (Kw)

A

ionization constant for water

[H3O+] x [OH-] = 1.0x10^-12

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53
Q

acidic solution

A

contains acid that increases [H3O+]

[OH-] decreases (= 1.0x10^-11 M)

[H3O+] > [OH-]

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54
Q

basic solution

A

contains a base that increases [OH-]

[H30+] decreases (= 1.-x10^-12 M)

[OH-] > [H3O+]

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55
Q

pH scale definition

A

specifies acidity of a solution

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56
Q

pH =

A
  • log [H30+]
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57
Q

pH < 7

A

acidic

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58
Q

pH > 7

A

basic

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59
Q

pH = 7

A

neutral

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60
Q

change of 1 pH unit corresponds to

A

10-fold change in [H30+]

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61
Q

if the [H30+] has two sig figs, report answer to

A

2 decimal points

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62
Q

pOH scale definition

A

defined with respect to [OH-]

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63
Q

pOH =

A
  • log [OH-]
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64
Q

pOH < 7

A

basic

65
Q

pOH > 7

A

acidic

66
Q

pH + pOH =

A

14.00

67
Q

pK

A

way of expressing strength of acid or base

68
Q

pKa =

A

-log (Ka)

69
Q

pKb =

A

-log (Kb)

70
Q

Ka =

A

10^-pKa

71
Q

Kb=

A

10^-pKb

72
Q

acid strength and pKa relation

A

stronger acid = smaller pKa

73
Q

base strength and pKb relation

A

stronger base = smaller pKb

74
Q

Ka and pKa relation

A

larger Ka = smaller pKa

because it is the -log

75
Q

Kb and pKb relation

A

larger Kb = smaller pKb

76
Q

what are the sources of H3O+ in a solution containing a strong or weak acid?

A

ionization of the acid
autoionization of H2O

77
Q

how much [H3O+] is produced during the autoionization of water at 25 degrees C? what happens to the additional [H3O+]?

A

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
Q

strong acids

A

completely ionize in solution

79
Q

weak acids

A

does not completely ionize in solution

80
Q

percent ionization =

A

[ionized acid] / initial [acid] x100%

81
Q

equilibrium [H30+] of a weak acid _ with increasing initial [acid]

A

increases

82
Q

percent ionization of a weak acid _ with increasing [acid]

A

decreases

83
Q

group 1A metal hydroxides are _ soluble in water and form

A

highly, strongly basic solutions

84
Q

group 2A metal hydroxides are _____ in water and produce

A

slightly soluble, 2 mol of OH- per mole of base

85
Q

group 2A metal hydroxides have the general formula

A

M(OH)2

86
Q

base ionization constant (Kb)

A

quantized extent of ionization of a weak base

87
Q

pKb =

A

-log (Kb)

88
Q

generally, anions form _ solutions

A

basic or neutral

89
Q

generally, cations form _ solutions

A

acidic or neutral

90
Q

an anion that is the conjugate base of a strong acid is

A

pH - neutral

91
Q

an anion that is the conjugate base of a weak acid is

A

a weak base

92
Q

general acid and conjugate base relation

A

weaker acid = stronger conjugate base

93
Q

3 categories of cations acting as weak acids

A

-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
Q

polyprotic acid

A

contains 2+ ionizable protons
-typically ionize in successive steps, each with its own Ka

95
Q

is Ka1 or Ka2 smaller? why

A

Ka2, first H+ separates from a neutral molecule while the second separates from an anion

96
Q

in which Ka step is the most amount of [H3O+] formed? why

A

first step
formation of H3O+ in the first step inhibits formation of additional H3O+ in subsequent steps

97
Q

strength of an acid and strength of the bond between H and molecule proportion

A

inversely

98
Q

bond polarity

A

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
Q

bond strength

A

strength of H-Y bond on a binary acid affects strength of acid

stronger bond (higher electronegativity on Y) = weaker acid

100
Q

combined effect of bond polarity and strength trends on PT

A

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
Q

oxyacids contain

A

H, O, and other element

102
Q

oxyacids bond structure

A

O atoms bonded to central atom, H atoms bonded to O atoms

103
Q

bond strength in oxyacids depends on

A

-electronegativity of central atom
-number of O atoms bonded to the central atom
-number of O atoms per H atom

104
Q

electronegativity and acid strength proportion in oxyacids. explain

A

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
Q

number of O atoms and acid strength proportion. explain

A

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
Q

which is stronger, an acid with more O atoms per H atoms or an acid with more H atoms per O atoms? why?

A

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
Q

organic acids

A

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
Q

buffers

A

resists pH change by neutralizing added acid or added base

109
Q

buffers contain either:

A

Significant amounts of a weak acid and its conjugate base

Significant amounts of a weak base and its conjugate acid

110
Q

how do buffers work?

A

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
Q

common ion effect

A

solution contains two substances that share a common ion

112
Q

In any buffer where the acid and conjugate base concentrations are equal, [H3O+] =

A

Ka

113
Q

Henderson-Hasselbalch Equation

A

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
Q

what does the Henderson-Hasselbalch Equation allow us to calculate?

A

pH of a buffer solution from the initial concentrations of the buffer components as long as the x is small approximation is valid

115
Q

What is the pH of a buffer solution when the concentrations of both buffer solutions (weak acid and its conjugate base) are equal?

A

pH = pKa

116
Q

What happens to the pH when the buffer contains more of the weak acid than the conjugate base?

A

pH decreases

117
Q

What happens to the pH when the buffer contains more of the conjugate base than the weak acid?

A

pH increases

118
Q

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?

A

Some of the conjugate base is used to neutralize the acid. So the [weak acid] > [conjugate base]

119
Q

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?

A

Some of the weak acid is used to neutralize the added base so [weak acid] < [conjugate base]

120
Q

how can you find the pH of a solution composed of a base and its conjugate acid?

A

calculate pKa for the conjugate acid of the weak base by subtracting pKb of the weak base from 14

121
Q

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?

A

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
Q

2 factors influence the effectiveness of a buffer:

A

-relative amounts of the acid and conjugate base
-absolute concentrations of the acid and conjugate base

123
Q

when is a buffer effective?

A

-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
Q

buffering capacity

A

amount of acid or base a buffer can neutralize

125
Q

buffering range

A

pH range over which the buffer can be effective

126
Q

a concentrated buffer can neutralize (more/less) added acid or base than a dilute buffer

A

more

127
Q

acid-base titration:

A

a titrant is slowly added to a solution of known concentration from a biuret until the reaction is complete

128
Q

titrant

A

solution of unknown concentration used in a titration

129
Q

indicator definition

A

chemical that changes color when the pH changes

130
Q

how do we know when we reach the endpoint of a titration?

A

an indicator will change color

131
Q

equivalence point of a titration

A

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
Q

what’s a titration curve?

A

plot of pH vs. amount of added titrant

133
Q

what’s the significance of the inflection point on a titration curve?

A

its the equivalence point of the titration

134
Q

prior to equivalence point:

A

known solution in the flask is in excess

pH is closest to pH of the known solution

135
Q

beyond the equivalence point:

A

unknown solution (titrant) added from the biuret (and OH-?) is in excess

pH approaches pH of titrant (increases)

136
Q

when will the titration curve be decreasing?

A

when the acid is in the biuret and the base is in the flask

137
Q

in a titration, initial pH =

A

pH of the weak acid solution

138
Q

how do you calculate the pH in a titration problem?

A

solve an equilibrium problem using [weak acid] as the initial concentration

139
Q

when does the solution become a buffer in a titration?

A

between the initial pH and the equivalence point

140
Q

buffer region

A

between initial pH and equivalence point (both acid and its conjugate base are present)

141
Q

main difference from the titration of a strong acid and weak acid

A

equivalence point of a titration of a strong acid is neutral

equivalence point of a titration of a weak acid is basic

142
Q

main difference between titrating a weak acid and a strong base compared to other titrations

A

curve starts basic and has an acidic equivalence point

143
Q

when a diprotic acid is titrated with a strong base:

A

-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
Q

In the titration of a strong acid with a strong base, how would you calculate initial pH?

A

–log [acid]

145
Q

In the titration of a strong acid with a strong base, how would you calculate pH at the equivalence point?

A

For the titration of a strong acid with a strong base, pH at the equivalence point is 7.00

146
Q

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

A

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
Q

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

A

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
Q

In the titration of a weak acid with a strong base, how would you calculate initial pH?

A

Use an equilibrium problem (ICE table) for the ionization of the weak acid

149
Q

In the titration of a weak acid with a strong base, how would you calculate pH at one half of the equivalence point?

A

pH = pKa

150
Q

indicators are _ acids

A

weak

151
Q

indicators establish an equilibrium with

A

H2O and H3O+ in the solution

152
Q

color of a titration solution depends on the relative concentrations of:

A

indicator : H indicator +

153
Q

when pH = pKa (indicator)

A

[ln-]/[Hln] = 10^0 = 1
intermediate color

154
Q

when pH = pKa + 1 (indicator)

A

[ln-]/[Hln] = 10^1 = 10
color of ln-

155
Q

when pH = pKa - 1 (indicator)

A

[ln-]/[Hln] = 10^-1 = 0.10
color of Hln

156
Q

at equivalence point. pH almost equals

A

pKa of H indicator+

157
Q

What is the endpoint of an indicator and what is it used for?

A

An indicator changes color to indicate the equivalence point of a titration

158
Q

An indicator is often a weak acid mixed with its conjugate base. How is the acid different from its conjugate base?

A

Acid form is a different color than its conjugate base

159
Q

Why is it important to use small amounts of indicator?

A

The indicator should have no effect on the pH or equivalence point of the neutralization reaction

ap chem (11 decks)

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