Unit # 9 Acids, Bases, and Salts
I. Acid-Base Theories:
A. Arrhenius Acids and Bases:
 Arrhenius acid = an acid that gives _________________________________________
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monoprotic acid = an acid containing ________________________________,
example = ______________
-
diprotic acid = an acid containing ___________________________________,
example = ______________
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triprotic acid = an acid containing __________________ ionizable hydrogens,
example = ______________
*** A hydrogen atom is only ionizable if it is _____________________________
______________________________, otherwise it will not be released in solution.
Question: How many of the hydrogen atoms in acetic acid are ionizable?
H O
H—C—C—O—H
H
 Arrhenius base = a base that gives _________________________________________
Properties of Acids and Bases
Acid
Base
Taste
pH range
Color change with
universal indicator
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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B. Bronsted-Lowry acids and Bases (a broader definition):
 Bronsted-Lowry acid = __________________________________________________
 Bronsted-Lowry base = ________________________________________________ –
__________________ definition than an Arrhenius base
Example: NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
 Conjugate Acids and Bases:
Example: NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
- consider the reverse reaction of ammonium and hydroxide ion, which substance
is the acid? __________ which is the base? __________ … so NH4+ is called a
________________________ and OH- is called a ________________________
- A conjugate acid-base pair is made up of the ____________________________
_______________________________________________________________OR
the base and its corresponding conjugate acid in the products – consists of ______
__________________________________________________________________
Additional Example: HCl(g) + H2O(l)
H3O+(aq) + Cl-(aq)
acid = ___________
base = ___________
conjugate acid = ___________
conjugate base = ___________
 amphoteric substances = substances that can _________________________________
_________________________________________________ – like _________________
C. Lewis Acids and Bases:
 Lewis acid = __________________________________________________________
 Lewis base = __________________________________________________________
Example: NH3(aq) + H2O(l)
NH4+(aq) + OH-(aq)
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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II. Hydrogen Ions and Acidity:
A. Hydrogen Ions:
 self-ionization = _______________________________________________________
H2O(l)  H+(aq) + OH-(aq)
note: H+(aq) = H3O+
- the ionization of water is _____________________ and Le Chatelier’s principle
applies
- the product of the hydrogen-ion concentration and the hydroxide-ion
concentration = __________________ @ 25°C = _________________________
_____________________________________________________________ = Kw
Kw = [H+] * [OH-] = 1.0 * 10-14
 acidic solution = _______________________________________________________
 basic solution (alkaline solution) = _________________________________________
_____________________________________________________________________
Example problem:
If the [H+] in a solution is 1.0 * 10-5 M, is the solution acidic, basic, or neutral?
What is the [OH-] of this solution?
B. pH:
 pH = ________________________________________________________________
_____________________________________________________________________
pH = -log[H+]
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in pure water [H+] = 1*10-7 M and the pH is 7.
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Review table 19.5 on page 656.
Example problem:
What is the pH of a solution with a hydrogen-ion concentration of 4.2 * 10-10 M?
Example problem:
The pH of an unknown solution is 6.35. What is the hydrogen-ion concentration?
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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Example problem:
What is the pH of a solution if [OH-] = 4.0 * 10-11 M?
C. Measuring pH:
 acid-base indicators or ________________________ can be used to measure pH.
 acid-base indicators = ___________________________________________________
_____________________________________________________________________
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review figure 19.8 on page 660.
 pH meter = ___________________________________________________________
_____________________________________________________________________
III. Strong and Weak Acids and Bases
 Acids and bases are classified as strong or weak depending on ______________________
___________________________________________________________________________.
A. Acids
strong acid = an acid that ________________________________ in aqueous solution
Examples: hydrochloric acid, sulfuric acid, nitric acid
HCl(g) + H2O(l)  H3O+(aq) + Cl-(aq)
*** notice the completion arrow (rxn goes to completion)
weak acid = an acid that _________________________________ in aqueous solution
Examples: acetic acid, carbonic acid
CH3COOH(aq) + H2O(l)
H3O+(aq) + CH3COO-(aq)
*** <1% ioniztion
Review table 19.6 on page 664.
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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B. Acid Dissociation Constant – Ka
Write the Keq for the dissociation of CH3COOH(aq) above.
*** for dilute aqueous solutions the concentration of water is a constant so,
Ka = Keq * [H2O] =
[H3O+] * [CH3COO-]
[CH3COOH]
 the acid dissociation constant (Ka) measures the extent to which an acid
dissociates in aqueous solution, and so can be used to represent the ___________
of an acid. The stronger the acid, the _____________________________ value.
 Example Calculation:
At equilibrium a solution of acetic acid has the following concentrations:
[CH3COOH] = 0.0987 M, and [H3O+] = [CH3COO-] = 1.34 * 10-3 M. Find Ka.
 Additional Example Calculation:
In a 0.200 M solution of a monoprotic weak acid, [H+] = 9.86 * 10-4 M. What is
the Ka for this acid?
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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C. Bases
strong base = a base that ________________________________ into ____________
___________ and _________________________________ in aqueous solution
Examples: calcium hydroxide, magnesium hydroxide, sodium hydroxide
weak base = a base that _________________________________________________
__________________________________________________________ – amount of
dissociation is relatively __________________
Examples: acetic acid, carbonic acid
NH4+(aq) + OH-(aq)
NH3(aq) + H2O(l)
*** this equilibrium favors the reactants
D. Base Dissociation Constant – Kb
Write the Keq for the dissociation of NH3(aq) above.
*** for dilute aqueous solutions the concentration of water is a constant so,
Kb = Keq * [H2O] =
Or in general form:
Kb =
[conjugate acid] * [OH-]
[base]
 the base dissociation constant (Kb) measures the extent to which a base
dissociates in aqueous solution, and so can be used to represent the ___________
of a base. The stronger the base, the _____________________________ value.
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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E. Differentiating Concentration and Strength: Strength refers to the extent of
_________________________________________ of an acid or base, while
concentration indicates how much of a substance is _______________________. 
Recall that substances __________________________________________________
(sugar).
Acidic Solution
Molar
Concentration
Hydrochloric acid 12 M HCl
Relative
Concentration
Concentrated
Relative
Strength
Strong
Gastric Juice
(stomach acid)
Acetic acid
0.08 M HCl
Concentrated
Weak
Vinegar
0.2 M CH3COOH
17 M CH3COOH
*** Strength is _________________________ of concentration and vice versa.
IV. Neutralization Reactions
A. Acid-Base Reactions = Neutralization Reactions
 In general acids and bases react to produce a _____________ (ionic substance)
and __________________.
 The complete reaction of a strong acid and a strong base produces a neutral
solution – _________________________________________.
Example: HCl(aq) + NaOH(aq) 
 Mole Ratios and Neutralization Reactions (example problem)
How many moles of sulfuric acid are needed to neutralize 0.50 moles of sodium
hydroxide?
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B. Titration
Titration = the process of ________________________________________________
_____________________________________________________________________
– using a neutralization reaction to determine concentration
Steps:
1. a measured volume of an acid solution of _____________________________
is added to a flask
2. an ___________________________ is added
3. measured volumes of a _________________________________________ are
mixed into the acid until the indicator just barely changes color
 The solution of known concentration is called the ______________________.
 Neutralization occurs (titration is complete) when _______________________
_________________________________________________________________.
= the equivalence point = end point
 Determining the concentration by titration mathematically:
Example: A 25 mL solution of H2SO4 is neutralized by 18 mL of 1.0 M
NaOH. What is the concentration of the H2SO4 solution? The equation
for the reaction is:
H2SO4(aq) + 2NaOH(aq)  Na2SO4(aq) + 2H2O(l)
Conversion plan:
Additional Example:
How many milliliters of 0.45 M HCl will neutralize 25.0 mL of 1.00 M
KOH?
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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V. Salts in Solution
A. Salt Hydrolysis
Recall that some salts can hydrolyze water – _____________________________
________________________________________________________ (Hydrolysis
of Salts Lab)  salt hydrolysis.
 salts that produce acidic solutions have _________________________
____________________________________________________________
 salts that produce basic solutions have __________________________
____________________________________________________________
Example: ammonium chloride
 Ammonium chloride completely ionizes in water:
NH4Cl(aq) 
 the ammonium ion is a strong enough acid to donate a hydrogen ion to
a water molecule:
NH+4(aq) + H2O(l) 
 the resulting H3O+ ions make the solution somewhat _______________
Strong acid + Strong base  neutral solution
Strong acid + Weak base  acidic solution (salt’s cation releases hydrogens to water)
Weak acid + Strong base  basic solution (salt’s anion attracts hydrogens from water)
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B. Buffers
Buffer = a solution that ___________________________________________ – a
solution of a weak acid and one of its salts or a solution of a weak base and one of
its salts
 a buffer contains ___________________________________________
_________________________ and _______________________________
_____________________________________________, therefore adding
hydrogen ions and/or hydroxide ions has little effect on the pH of the
solution
 Example: A solution of carbonic acid and sodium bicarbonate forms a
buffer:
 If you add H+ ions:
HCO3- + H+
 If you add OH- ions
H2CO3 + OH Buffer Capacity = the _______________________________________
____________________________________________________________
_____________________________________ before a significant change
in pH occurs (before your H+ ion acceptors and/or H+ ion donors run out)
Important Buffer Systems
Buffer name
Ethanoic acid-ethanoate ion
Dihydrogen phosphate ion-hydrogen
phosphate ion
Formulas
Buffer pH*
CH3COOH/CH3COO−
4.76
H2PO4−/HPO42−
7.20
H2CO3/HCO3−
6.46
NH4+/NH3
9.25
Carbonic acid-hydrogen carbonate ion
(solution saturated with CO2)
Ammonium ion-ammonia
* Components
have concentrations of 0.1M.
 Example Problem:
A buffer consists of methanoic acid (HCOOH) and methanoate ion
(HCOO-) Write an equation to show what happens when an acid is added
to this buffer. And a base.
Unit # 9 (Book Chapter 19) Acids, Bases, and Salts
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