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Chapter 10
Acids, Bases, and Salts
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Ch 10.1 Arrhenius Acid-Base Theory (also in Chapter Medley)
Arrhenius Acids produce
H+ in water
HCl
HNO3
HClO4
H2SO4
H3PO4
hydrochloric acid
nitric acid
perchloric acid
sulfuric acid
phosphoric acid
Arrhenius Bases produce
OH- in water
KOH
Ba(OH)2
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Arrhenius Acids
Sour taste
Change blue litmus paper to ___
Corrosive
Arrhenius Bases
Bitter taste
Change red litmus paper to ___
Slippery (soapy) to the touch
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Ch 10.2 Brønsted–Lowry Acid-Base Theory
Brønsted–Lowry acid = proton (H+ ion) donor
Brønsted–Lowry base = proton (H+ ion) acceptor
In aqueous solution, a proton is bonded to water
through a coordinate covalent bond:
_____________ Ion
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Base
H+ acceptor
Acid
H+ donor
H2O(l) + HCl(g) → H3O+(aq) + Cl-(aq)
Base
H+ acceptor
Acid
H+ donor
NH3(g) + HCl(g) →
[NH4+Cl-] → NH4Cl(s)
Figure 10.3 White cloud of solid _________
formed from gaseous HCl and NH3.
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Conjugate Acid-Base pairs differ by a single _____
Acid
H+ donor
Base
H+ acceptor
Acid
H+ donor
Base
H+ acceptor
HCl(g) + H2O(l) → H3O+(aq) + Cl-(aq)
Acid
Base
Conjugate
Acid
Conjugate
______
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Drill Problem. Write the chemical formula for:
1. The conjugate base of H2PO42. The conjugate acid of H2PO43. The conjugate base of H2O
4. The conjugate acid of H2O
Take note that H2PO4- and H2O can act as H+ donor
or H+ acceptor. Substances that can either donate or
accept a H+ are called amphiprotic.
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Ch 10.3 Mono-, Di-, and Triprotic Acids
Monoprotic acids can transfer 1 H+ to H2O or base
Examples: HCl and HNO3
Diprotic acids can transfer 2 H+
Example: H2SO4 + H2O → H3O+ + HSO4HSO4- + H2O
H3O+ + SO42Triprotic acids can transfer ______
Example: H3PO4 + H2O
H2PO4− + H2O
HPO42− + H2O
H3O+ + H2PO4−
H3O+ + HPO42−
H3O+ + PO43−
A polyprotic acid supplies _____________
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Only acidic H atoms are
donated:
Acetic acid is ______protic
Acidic
CH3CO2H(l) + H2O(l)
CH3CO2-(aq) + H3O+(aq)
acetate ion
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Ch 10.4 Strengths of Acids and Bases
A strong acid donates all or nearly 100% of its H+ to ____
Table 10.1
Learn the names
and formulas of
these commonly
encountered
strong acids, and
then assume that
all other acids
you encounter
are weak, unless
you are told
otherwise.
These acids are strong even in ________solution because
in water they are all or mostly ionized.
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A weak acid does not ionize completely.
Acetic acid is a weak acid; less than ____ of its molecules
are ionized:
CH3CO2H(l) + H2O(l)
CH3CO2-(aq) + H3O+(aq)
Figure 10.5 Comparison of ionized species for a strong and a weak acid.
demo
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Strong bases are
limited to the
hydroxides of
Group IA and IIA
listed in Table 10.2.
Ammonia gas (NH3) is the most common ________ base.
In water, it produces only small amounts of OH- ions:
NH3(g) + H2O(l)
NH4+(aq) + OH-(aq)
Less than _____ of ammonia is ionized.
demo
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Ch 10.5 Ionization Constants for Acids and Bases
HA(aq) + H2O(l)
H3O+(aq) + A-(aq)
acid ionization constant Ka = [H3O+][A-]
[HA]
Recall from Ch 9 that the concentrations of liquids and solids are
_______________ into the equilibrium constants.
Table 10.3 Ka and % ionization values for selected weak acids
(1.0 M) at 25oC
Acid strength increases with _____________ Ka values.
Note that for polyprotic acids, each successive step of proton
transfer occurs to a lesser extent than the previous one.
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To compare bases, we look at their tendency to accept
protons from water and generate _____________ ions:
B(aq) + H2O(l)
BH+(aq) + OH-(aq)
base ionization constant Kb = [BH+][OH-]
[B]
NH3(g) + H2O(l)
NH4+(aq) + OH-(aq)
Kb = 1.8 x 10-5 0.42% is ionized
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Acid-Base Neutralization Reactions
(Ch 10.1, 10.6 & 10.7 covered in Chapter Medley)
Acid + Base → Salt + Water
HX + BOH → BX + HOH
Double-replacement reaction
_____ ionic equation: H+ + OH- → H2O
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Ch 10.8 Self-Ionization of Water
H3O+(aq)
H2O(l) + H2O(l)
Keq = [H3O+][OH-]
[H2O]2
2
+ OH-(aq)
[H2O] = ______ M
+
-
memorize this #
Keq[H2O] = Kw = [H3O ][OH ] = 1.00 x 10-14
Kw = Ion product constant for pure H2O at 25oC
In pure H2O at 25oC
[H3O+] = [OH-] = 1.00 x 10-7 M
_________
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What happens to [H3O+] & [OH-] when an acid is added to H2O?
Example. When HCl is added to produce a 0.010 M solution:
0.010 M HCl → 0.010 M H3O+ acidic solution
[H3O+] = 1.0 x 10-2
100,000 times more than in pure H2O
In acidic solution [H3O+] > ______
And what happens to [OH-]?
H3O+
2 H2O
Le Châtelier
+
large # added
with HCl
pushes to left
OHsome must ______
[OH-] decreases
Kw = [H3O+][OH-] = 1.00 x 10-14
[OH-] = Kw/[H3O+] = 1.00 x 10-14/1.0 x 10-2
[OH-] = __________
In acidic solution [H3O+] > [OH-]
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What happens to [H3O+] & [OH-] when a base is added to H2O?
Example.
When NaOH is added to produce a 0.0010 M solution:
0.0010 M NaOH → 0.0010 M OH- basic solution
[OH-] = 1.0 x 10-3
10,000 times more than in pure H2O
In basic solution [OH-] > ______
And what happens to [H3O+]?
2 H2O
H3O+
Le Châtelier
+
[H3O+] decreases
OHlarge # added
pushes left
Kw = [H3O+][OH-] = 1.00 x 10-14
[H3O+] = Kw/[OH-] = 1.00 x 10-14/1.0 x 10-3
[H3O+] = _________
In basic solution [OH-] > [H3O+]
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Figure 10.9 Summary of relationship between [H3O+] and [OH-]
I
n
c
r
e
a
s
e
Correction note: T = 25oC for values throughout this chapter
Table 10.4
Summary
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Ch 10.9 The pH concept
The pH scale offers an easier method of expressing acidity
pH is the negative logarithm of the ___________ of the
hydronium ion:
pH = -log[H3O+]
when [H3O+] = 1 x 10
-x
then pH = x
-9
[H3O+] = 1.0 x 10 then pH = 9.00
2 sig fig
2 digits
If the coefficient in the exponential expression is not 1, then the
pH value will be ____________:
[H3O+] = 6.3 x 10
2 sig fig
-5
pH = 4.20
2 digits
enter 6.3 x 10-5 in your calculator, press the log key, then switch the ___
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Relationships among pH
values, [H3O+], and [OH-] at 25oC
Figure 10.11
Higher [H3O+] = lower _____
∆ in 1 pH unit = 10-fold ∆ in [
]
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Fig 10.10 Most fruits
and vegetables are
acidic (tart or sour taste)
Fig 10.12 pH values
of common liquids
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Measurements of pH
Impregnated strips of paper, pH
test paper, are still widely in use.
HInd + H2O
H3O+ + Ind-
Phenolphthalein is an acid-base
indicator added to solutions.
Figure 10.13 The pH meter
measures the voltage across a
special membrane that passes
only H+ when immersed in the
solution.
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Ch 10.10 The pKa Method for Expressing Acid Strength
pKa = -logKa
example: Ka acetic acid = 1.8 x 10-5 pKa = 4.74
Acid strength increases with increasing Ka (Ch 10.5), while
it increases with ____________ pKa.
Drill question: Compare acetic acid with HF. Which is
the stronger acid?
HF Ka = 6.8 x 10-4
pKa 3.17
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Ch 10.11 The pH of Aqueous Salt Solutions
Acid + Base → Water + ______
When a strong acid reacts with a strong base, a neutral salt is
formed:
HCl + NaOH → H2O + NaCl
pH = _____
When a weak acid and/or base is involved, the resulting salts
react with water (hydrolysis reaction) to produce hydronium
ions or hydroxide ions or both.
To understand this phenomenon, let’s have another look at the
ionization of acids and bases in water:
HA + H2O
weak
acid
B + H2O
weak
base
H3O+ + A-
strong
conjugate
base
BH+ + OHstrong
conjugate
acid
Strong conjugate acids and bases react with ________!
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1. Basic Hydrolysis:
Example. When sodium nitrite dissolves in water, it dissociates
completely:
NaNO2 → Na+(aq) + NO2-(aq)
strong conjugate base of a weak acid (HNO2)
has strong attraction for H+
NO2- + H2O
HNO2 + OH-
solution has higher [OH-]
than pure H2O,
solution is ______
NaNO2 = _______ salt!
Anions of strong acids do not hydrolyze because the weak
conjugate base cannot take H+ from H2O.
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2. Acidic Hydrolysis:
Example. When ammonium chloride dissolves in water, it
dissociates completely:
NH4Cl → Cl-(aq) + NH4+(aq)
strong conjugate acid of a weak base (NH3)
has strong tendency to donate H+
NH4+ + H2O
NH3 + H3O+
NH4Cl = ________ salt!
solution has higher [H3O+]
than pure H2O,
solution is ________
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Examples of neutral, acidic, and basic ions
Ions of Neutral Salts
Cations
Na+
K+
Rb+
Cs+
Mg2+ Ca2+ Sr2+ Ba2+
Acidic Ions
Basic Ions
F
C2H3O2- NO2- HCO3-
Transition metal ions
CN-
CO32-
HSO4- H2PO4-
HPO42- PO43-
NH4
+
Al
3+
Pb
2+
Sn
2+
-
S2-
Anions
Cl
-
Br-
I-,
ClO4- BrO4- ClO3- NO3-
Table 10.6 Examples of neutral, acidic, and basic salts.
Table 10.7 Approximate pH of selected 0.1 M salt solutions.
SO42-
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Ch 10.12 Buffers
demo
A buffer is a solution capable of absorbing small amounts
of acid or base without dramatic changes in solution ____
Buffered aspirin has a buffering agent,
such as MgO, that will maintain the pH
of the aspirin as it passes through the
stomach of the patient.
Many hair shampoos are
pH-controlled.
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Typically, a buffer is composed of a conjugate acid-base pair:
a weak acid and its conjugate base
CH3CO2H/CH3CO2H2PO4-/HPO42H2CO3/HCO3- (buffer in human blood)
• Added base is absorbed by the _______
OH- + H2CO3 → HCO3- + H2O
• Added acid is absorbed by the conjugate _______
H3O+ + HCO3- → H2CO3 + H2O
H2CO3 is unstable and decomposes; excess CO2 is expelled through the lungs.
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Typical graphs depicting how pH is buffered.
(Note that Figure 10.14 in your text is in error; the labeling of the two
lines must be switched.)
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Ch 10.14 Electrolytes
Solutions in which ions are present are _______ conductors of
electricity (recall demo comparing conductivity of ionic and
covalent compounds in Ch 4).
Strong Electrolyte (NaCl)
Nonelectrolyte (sugar)
Figure 10.15
Weak electrolytes, such as weak
acids and bases, incompletely
ionize/dissociate into ions in
aqueous solutions.
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Electrolytes are important components of body fluids and
essential to the proper functioning of the human body.
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Drill Problem. Classify each of the following compounds as a
strong electrolyte, weak electrolyte, or nonelectrolyte:
H3PO4
HCl
Cl2
HF
KBr
CH3CH2-OH
CH3COOH
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Ch 10.15 Equivalents and Milliequivalents of Electrolytes
For solutions containing multiple types of ions, concentrations
are often given in equivalent units. An equivalent (Eq) of an ion
is the molar amount of that ion needed to supply one mole of
positive or negative charge.
1 mEq = ______ Eq
1 mole Na+ = 1 equivalent
1 mole Ca2+ = 2 equivalents
1 mole PO43- = ___ equivalents
Table 10.8 Concentrations of Major Electrolytes in Blood Plasma
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Sample Calculation. Human blood plasma contains 2.4 mg
Mg2+ per dL. How many Eq or mEq are in 1.0 L of plasma?
Strategy: dL → L
mg → g → mol → Eq → mEq
1.0 L x 2.4 mg Mg2+ x 10 dL x . 1 g . x 1 mol x _____ Mg2+
mol Mg2+
1 dL
1L
103 mg 24.3 g
= 2.0 x 10-3 Eq Mg2+ or 2.0 mEq Mg2+
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Ch 10.16 Acid-Base Titrations
An acid-base titration is a neutralization
reaction in which a measured volume of
acid or base of known concentration is
completely reacted with a measured
volume of a base or acid of unknown
concentration. The concentration of the
unknown solution is then calculated.
To detect the endpoint or neutralization
point, an acid-base indicator, such as
phenolphthalein, is added.
HInd + H2O
H3O+ + Ind-
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In an acid-base titration, 13.07 mL of 0.100 M
H3PO4 is needed to neutralize 25.0 mL of KOH of unknown
concentration. Calculate the molarity of the KOH.
Sample calculation.
H3PO4 (aq) + 3 KOH(aq) → K3PO4(aq) + 3 H2O(l)
13.07 mL x 0.100 mol H3PO4 x ___ mole KOH = _______ mole KOH
1 mol H3PO4
1000 mL
M = mol/L = ________ mole KOH/0.025 L = ______ M KOH