1. Lemons and limes are examples of foods that contain acidic solutions.
15 Acids, Bases, and Salts
Lemons and limes are examples of foods
that contain acidic solutions.
Foundations of College Chemistry, 15th Ed.
Morris Hein, Susan Arena, and Cary Willard
Copyright © 2016 John Wiley & Sons, Inc. All rights reserved.

2. Copyright © 2016 John Wiley & Sons, Inc. All rights reserved.
Chapter Outline
15.1 Acids and Bases
15.2 Salts
15.3 Electrolytes and Nonelectrolytes
15.4 Introduction to pH
15.5 Neutralization
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3. Arrhenius Acid: An acid solution contains
Arrhenius Acids
Arrhenius Acid: An acid solution contains
an excess of H+ ions.
Common Properties of Acids
1. Sour taste
2. Turns litmus paper pink
3. Reacts with:
Metals to produce H2 gas
Bases to yield water and a salt
Carbonates to give carbon dioxide
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4. Arrhenius Bases: A basic solution contains
an excess of OH– ions.
Common Properties of Bases
1. Bitter/caustic taste
2. Turns litmus paper blue
3. Slippery, soapy texture
4. Neutralizes acids
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5. Summary of the Acid/Base Theories
The theory that best explains the reaction
of interest is used.
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6. Reactions of Acids 1. Reactivity with Metals
Acids react with any metals above
hydrogen in the activity series.
2 HCl (aq) + Mg (s) MgCl2 (s) + H2 (g)
In general:
acid + metal salt + hydrogen
2. Reactivity with Bases
Also called a neutralization reaction.
2 HCl (aq) + Ca(OH)2 (aq) CaCl2 (aq) + 2 H2O (l)
In general:
acid + base salt + water
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7. Reactions of Acids 3. Reactivity with Metal Oxides
2 HCl (aq) + Na2O (s) NaCl (aq) + H2O (l)
In general:
acid + metal oxide salt + water
(base)
4. Reactivity with Metal Carbonates
2 HCl (aq) + Na2CO3 (aq) NaCl (aq) + H2O (l) + CO2 (g)
In general:
acid + carbonate salt + water + carbon dioxide
(base)
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8. Base Reactions Bases can be amphoteric As a base:
Zn(OH)2 (aq) + 2 HBr (aq) ZnBr2 (aq) + 2 H2O (l)
As an acid:
Zn(OH)2 (aq) + 2 NaOH (aq) Na2Zn(OH)4 (aq)
NaOH and KOH can also react with metals.
2 NaOH (aq) + 2 Al (s) + 6 H2O (l)
2 NaAl(OH)4 (aq) + 3 H2 (g)
In general:
base + metal + water salt + hydrogen
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9. Salts Salts: products from acid-base reactions.
HCl (aq) + NaOH (aq) NaCl (aq) + H2O (l)
Sodium chloride
(table salt)
Salts are ionic compounds.
Salts contain a cation (a metal or ammonium ion)
derived from the base and an anion
(excluding oxide or hydroxide ions) derived from the acid.
Salts are generally crystalline compounds with
high melting and boiling points.
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10. Salts Practice Which salt forms in the reaction of
aluminum oxide and hydrobromic acid?
3 Al2O3 (s) + 18 HBr (aq) AlBr3 (aq) + 9 H2O (l)
a. AlBr
b. AlBr3
c. Al2Br
d. Al2Br3
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11. Electrolytes and Nonelectrolytes
Electrolytes: compounds that conduct electricity
when dissolved in water.
Nonelectrolytes: substances that do not conduct
electricity when dissolved in water.
Comparing Solution Conductivity
(Distilled water)
(Sugar solution)
(Salt solution)
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12. Electrolytes and Nonelectrolytes
Ion movement causes conduction of electricity in water.
3 classes of compounds, acids, bases, and salts
are electrolytes because they produce ions
in water when they dissolve.
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13. Electrolytes and Nonelectrolytes Practice
Which compound will not dissociate in water?
a. HCl
b. KBr
c. NaOH
d. CH3OH
a. HCl (acid, dissociates)
b. KBr (salt, dissociates)
c. NaOH (base, dissociates)
d. CH3OH (organic, does not dissociate)
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14. Dissociation of Electrolytes
Salts dissociate into their respective cations
and anions when dissolved in water.
NaCl (s) Na+ (aq) + Cl- (aq)
Hydrated sodium (purple) and chloride (green) ions
The negative end of
the water dipole is
attracted to the
positive Na+ ion.
The positive end of
the water dipole is
attracted to the
positive Cl- ion.
When NaCl dissolves in water,
each ion is surrounded by several water molecules.
The permanent dipoles in the water molecules cause
specific alignment around the ions.
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15. Electrolyte Ionization
Ionization: process of ion formation in solution.
Ionization results from the chemical reaction
between a compound and water.
Acids ionize in water, producing the
hydronium ion (H3O+) and a counter anion.
HCl (g) + H2O (l) H3O+ (aq) + Cl- (aq)
H3PO4 (aq) + H2O (l) H2PO4- (aq) + H3O+ (aq)
Bases ionize in water, producing the
hydroxide ion (OH-) and a counter cation.
NH3 (aq) + H2O (l) OH- (aq) + NH4+ (aq)
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16. Strong and Weak Electrolytes
Strong electrolytes: undergo complete ionization in water.
Example: HCl (strong acid)
Weak electrolytes: undergo incomplete ionization in water.
Example: CH3COOH (weak acid)
HCl (left) is 100% ionized.
CH3COOH exists primarily in the unionized form.
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17. Strong and Weak Electrolytes
Double arrows indicate incomplete ionization
(typically weak electrolytes).
HF (aq) + H2O (l) F- (aq) + H3O+ (aq)
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18. Salts Salts can dissociate into more than 2 ions,
depending upon the compound.
A 1 M solution of NaCl produces a total of 2 M of ions.
NaCl (s) Na+ (aq) + Cl- (aq) 1M 1M 1M
A 1 M solution of CaCl2 produces a total of 3 M of ions.
CaCl2 (s) Ca2+ (aq) + 2 Cl- (aq) 1M 1M 2M
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19. Salts Practice What is the concentration of bromide ion in a
1.5 M solution of magnesium bromide?
a. 0.75M
b. 1.5M
c. 3.0M
d. 4.5M
For every one mole of MgBr2, 2 moles of Br- ionize.
MgBr2 (s) Mg2+ (aq) Br- (aq)
1.5 M
1.5M
3.0M
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20. Autoionization of Water
Pure water auto(self) ionizes according to the reaction:
H2O (l) + H2O (l) H3O+ (aq) + OH– (aq)
Based on the reaction stoichiometry:
Concentration H3O+ = Concentration OH– = 1 x 10–7 M
[H3O+] x [OH–] = (1 x 10–7)2 = 1 x 10–14
When acid or base is present in water,
[H3O+] and [OH-] change.
In acidic solutions, [H3O+] > [OH–].
In basic solutions, [H3O+] < [OH–].
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21. In pure water, [H3O+] = 1 x 10-7 M, so
Introduction to pH
pH = - log[H3O+]
In pure water, [H3O+] = 1 x 10-7 M, so
pH = - log(1 x 10-7) = 7
Neutral
High H3O+
Low OH-
Increasing acidity
Increasing basicity
Low H3O+
High OH-
The pH scale
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22. Copyright © 2016 John Wiley & Sons, Inc. All rights reserved.pH
Copyright © 2016 John Wiley & Sons, Inc. All rights reserved.

23. pH Practice By what factor is a solution of pH = 3 more acidic
than a solution with a pH = 5?
a. 2
b. 20
c. 200
d. 100
pH = 5 means [H3O+] = 1 x 10-5 M
pH = 3 means [H3O+] = 1 x 10-3 M
1 x 10-3 M
Factor =
= 100
1 x 10-5 M
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24. Copyright © 2016 John Wiley & Sons, Inc. All rights reserved.
pH Calculations
pH = - log[H3O+]
[H3O+] = 10-pH
Generalizations
Exponent = pH
pH = 5
[H3O+] = 1 x 10-5 M
The pH is between
the exponent and
next lowest whole
number
pH = 4.7
If exactly 1
[H3O+] = 2 x 10-5 M
If a number between 1 and 10
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25. pH Calculations Practice
What is the pH of a M HCl solution? a
b. -2.0
c. 1.70
d. 1.7
pH = - log(0.020) = 1.7
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26. Titration Titration: experiment where the volume of one reagent
(titrant) required to react with a measured amount
of another reagent is measured.
Titrations allow the amount of an acid or base
present in a sample to be determined.
Indicators are used to signal the endpoint of a titration,
the point when enough titrant is added to react
with the acid/base present.
Burets deliver measured amounts of the titrant
into a solution of the unknown reagent.
Endpoint
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27. Titration Calculations
What is the molarity of a NaOH solution if mL of
base is required to titrate g of oxalic acid (H2C2O4)?
Reaction
H2C2O4 (aq) + 2 NaOH (aq) NaC2O4 (aq) + 2 H2O (l)
Knowns g H2C2O4
21.93 mL NaOH
Solving for Molarity of base
Calculate
1 mol H2C2O4
90.04 g H2C2O4
2 mol NaOH
0.243 g H2C2O4 × ×
= mol NaOH
1 mol H2C2O4
mol NaOH
L soln
Molarity NaOH =
= M NaOH
Always check reaction stoichiometry!
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28. Learning Objectives 15.1 Acids and Bases 15.2 Salts
Compare the definitions of acids and bases, including
Arrhenius
15.2 Salts
Explain how a salt is formed and predict the formula
of a salt given an acid and base precursor.
15.3 Electrolytes and Nonelectrolytes
Describe properties, ionization, dissociation, and strength
of electrolytes and compare them to nonelectrolytes.
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29. Calculate the pH of a solution from the hydrogen ion concentration.
Learning Objectives
15.4 Introduction to pH
Calculate the pH of a solution from the
hydrogen ion concentration.
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