1. Christopher G. Hamaker, Illinois State University, Normal IL
Introductory Chemistry: Concepts & Connections 4th Edition by Charles H. Corwin
Chapter 15
Acids and Bases
Christopher G. Hamaker, Illinois State University, Normal IL
© 2005, Prentice Hall

2. Properties of Acids
An acid is any substance that releases hydrogen ions, H+, into water.
Blue litmus paper turns red in the presence of hydrogen ions. Blue litmus is used to test for acids.
Acids have a sour taste; lemons, limes, and vinegar are acidic.
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3. Properties of Bases
A base is a substance that releases hydroxide ions, OH –, into water.
Red litmus paper turns blue in the presence of hydroxide ions. Red litmus is used to test for bases.
Bases have a slippery, soapy feel.
Bases also have a bitter taste; milk of magnesia is a base.
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4. Acid/Base Neutralization
Recall that an acid and a base react with each other in a neutralization reaction.
When an acid and a base react, water and a salt are produced.
For example, nitric acid reacts with sodium hydroxide to produce sodium nitrate and water:
HNO3(aq) + NaOH(aq) → NaNO3(aq) + H2O(l)
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5. The pH Scale
A pH value expresses the acidity or basicity of a solution.
Most solutions have a pH between 0 and 14.
Acidic solutions have a pH less than 7.
As a solution becomes more acidic, the pH decreases.
Basic solutions have a pH greater than 7.
As a solution becomes more basic, the pH increases.
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6. Acid/Base Classifications of Solutions
A solution can be classified according to its pH:
Strongly acidic solutions have a pH less than 2.
Weakly acidic solutions have a pH between 2 and 7.
Weakly basic solutions have a pH between 7 and 12.
Strongly basic solutions have a pH greater than 12.
Neutral solutions have a pH of 7.

7. Buffers
A buffer is a solution that resists changes in pH when an acid or a base is added.
A buffer is a solution of a weak acid and one of its salts:
Citric acid and sodium citrate make a buffer solution
When acid is added to the buffer, the citrate reacts with the acid to neutralize it.
When base is added to the buffer, the citric acid reacts with the base to neutralize it.
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8. Arrhenius Acids and Bases
Svante Arrhenius proposed the following definitions for acids and bases in 1884:
An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions.
An Arrhenius base is a substance that ionizes in water to release hydroxide ions.
For example, HCl is an Arrhenius acid and NaOH is an Arrhenius base.
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9. Strengths of Acids Acids have varying strengths.
The strength of an Arrhenius acid is measured by the degree of ionization in solution.
Ionization is the process where polar compounds separate into cations and anions in solution.
The acid HCl ionizes into H+ and Cl– ions in solution.
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10. Strengths of Bases Bases also have varying strengths.
The strength of an Arrhenius base is measured by the degree of dissociation in solution.
Dissociation is the process where cations and anions in an ionic compound separate in solution.
A formula unit of NaOH dissociates into Na+ and OH– ions in solution.
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11. Strong & Weak Arrhenius Acids
Strong acids ionize extensively to release hydrogen ions into solution.
HCl is a strong acid and ionizes nearly 100%
Weak acids only ionize slightly in solution.
HF is a weak acid and ionizes only about 1%
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12. Arrhenius Acids in Solution
All Arrhenius acids have a hydrogen atom bonded to the rest of the molecule by a polar bond. This bond is broken when the acid ionizes.
Polar water molecules help ionize the acid by pulling the hydrogen atom away:
HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq) (~100%)
HC2H3O2(aq) + H2O(l) → H3O+(aq) + C2H3O2–(aq) (~1%)
The hydronium ion, H3O+, is formed when the aqueous hydrogen ion attaches to a water molecule.
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13. Strong & Weak Arrhenius Bases
Strong bases dissociate extensively to release hydroxide ions into solution.
NaOH is a strong base and dissociates nearly 100%
Weak bases only ionize slightly in solution.
NH4OH is a weak base and only partially dissociates
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14. Arrhenius Bases in Solution
When we dissolve Arrhenius bases in solution, they dissociate giving a cation and a hydroxide anion.
Strong bases dissociate almost fully and weak bases dissociate very little:
NaOH(aq) → Na+(aq) + OH–(aq) (~100%)
NH4OH(aq) → NH4+(aq) + OH–(aq) (~1%)
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15. Neutralization Reactions
Recall, an acid neutralizes a base to produce a salt and water.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
The reaction produces the aqueous salt NaCl.
If we have an acid with two hydrogens (sulfuric acid, H2SO4), we need two hydroxide ions to neutralize it.
H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + 2 H2O(l)
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16. Predicting Neutralization Reactions
We can identify the Arrhenius acid and base that react in a neutralization reaction to produce a given salt such as calcium sulfate, CaSO4.
The calcium must be from calcium hydroxide, Ca(OH)2, and the sulfate must be from sulfuric acid, H2SO4.
H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2 H2O(l)
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17. Brønsted-Lowry Acids & Bases
The Brønsted-Lowry definitions of acids and bases are broader than the Arrhenius definitions.
A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor.
A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor.
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18. Brønsted-Lowry Acids & Bases
Lets look at two acid/base reactions:
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
HCl(aq) + NH3(aq) → NH4Cl(aq)
HCl donates a proton in both reactions and is a Brønsted-Lowry acid.
In the first reaction, the NaOH accepts a proton and is the Brønsted-Lowry base.
In the second reaction, NH3 accepts a proton and is the Brønsted-Lowry base.
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19. Amphiprotic Compounds
A substance that is capable of both donating and accepting a proton is an amphiprotic compound.
NaHCO3 is an example:
HCl(aq) + NaHCO3(aq) → NaCl(aq) + H2CO3(aq)
NaOH(aq) + NaHCO3(aq) → Na2CO3 (aq) + H2O(l)
NaHCO3 accepts a proton from HCl in the first reaction and donates a proton to NaOH in the second reaction.
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20. Acid-Base Indicators
A solution that changes color as the pH changes is an acid-base indicator.
Three common indicators are methyl red, bromothymol blue, and phenolphthalein.
Each has a different color above and below a certain pH.
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21. Acid-Base Indicators
Shown below are the three indicators at different pH values
Methyl Red Bromothymol Blue Phenolphthalein
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22. Acid-Base Titrations
A titration is used to analyze an acid solution using a solution of a base.
A measured volume of base is added to the acid solution. When all of the acid has been neutralized, the pH is 7. One extra drop of base solution after the endpoint increases the pH dramatically.
When the pH increases above 7, phenolphthalein changes from colorless to pink indicating the endpoint of the titration.
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23. Titration Problem
Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires mL of M NaOH. What is the concentration of the acetic acid?
HC2H3O2(aq) + NaOH(aq) → NaC2H3O2(aq) + H2O(l)
We want concentration acetic acid, we have concentration sodium hydroxide.
conc NaOH  mol NaOH 
mol HC2H3O2  conc HC2H3O2
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24. Titration Problem Continued
The molarity of NaOH can be written as the unit factor mol NaOH / 1000 mL solution.
= mol HC2H3O2
37.55 mL solution ×
1 mol HC2H3O2
1 mol NaOH
0.233 mol NaOH
1000 mL solution ×
1000 mL solution
10.0 mL solution
mol HC2H3O2
1 L solution ×
= M HC2H3O2
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25. Another Titration Problem
A 10.0 mL sample of M H2SO4 is titrated with M NaOH. What volume of NaOH is required for the titration?
We want mL of NaOH, we have 10.0 mL of H2SO4.
Use mol H2SO4/1000 mL and mol NaOH/1000 mL.
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26. H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l)
Problem Continued
H2SO4(aq) + 2 NaOH(aq) → Na2SO4(aq) + H2O(l)
= 49.8 mL NaOH
10.0 mL H2SO4 ×
1 mol H2SO4
2 mol NaOH
0.555 mol H2SO4
1000 mL H2SO4 ×
0.233 mol NaOH
1000 mL NaOH
49.8 mL of M NaOH is required to neutralize 10.0 mL of M H2SO4.
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27. Acid-Base Standardization
A standard solution is a solution where the concentration is precisely known.
Acid solutions are standardized by neutralizing a weighed quantity of a solid base.
What is the molarity of a hydrochloric acid solution if mL are required to neutralize g Na2CO3?
2 HCl(aq) + Na2CO3(aq) → 2 NaCl(aq) + H2O(l) + CO2(g)
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28. Standardization Continued
= mol HCl
0.375 g Na2CO3 × ×
1 mol Na2CO3
g Na2CO3
2 mol HCl
= M HCl ×
25.50 mL solution
mol HCl
1000 mL solution
1 L solution
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29. Ionization of Water
Water undergoes an autoionization reaction. Two water molecules react to produce a hydronium ion and a hydroxide ion:
H2O(l) + H2O(l) → H3O+(aq) + OH-(aq) or
H2O(l) → H+(aq) + OH-(aq)
Only about 1 in 5 million water molecules is present as ions so water is a weak conductor.
The concentration of hydrogen ions, [H+], in pure water is about 1 × 10-7 mol/L at 25C.
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30. Autoionization of Water
Since [H+] is 1 × 10-7 mol/L at 25C, the hydroxide ion concentration, [OH-], must also be 1 × 10-7 mol/L at 25C:
H2O(l) → H+(aq) + OH-(aq)
At 25C
[H+][OH-] = (1 × 10-7)(1 × 10-7) = 1.0 × 10-14
This is the ionization constant of water, Kw.
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31. [H+] and [OH-] Relationship
At 25C, [H+][OH-] = 1.0 × So, if we know the [H+], we can calculate [OH-].
What is the [OH-] if [H+] = 0.1 M ?
[H+][OH-] = 1.0 × 10-14
(0.1)[OH-] = 1.0 × 10-14
[OH-] = 1.0 × 10-13
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32. The pH Concept
Recall that pH is a measure of the acidity of a solution.
A neutral solution has a pH of 7, an acidic solution has a pH less than 7, and a basic solution has a pH greater than 7.
The pH scale uses powers of ten to express the hydrogen ion concentration.
Mathematically: pH = –log[H+]
[H+] is the molar hydrogen ion concentration
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33. Calculating pH
What is the pH if the hydrogen ion concentration in a vinegar solution is M?
pH = –log[H+]
pH = –log(0.001)
pH = – ( –3) = 3
The pH of the vinegar is 3, so the vinegar is acidic.
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34. Calculating [H+] from pH
If we rearrange the pH equation for [H+], we get:
[H+] = 10–pH
Milk has a pH of 6. What is the concentration of hydrogen ion in milk?
[H+] = 10–pH = 10–6 = M
[H+] = 1 × 10–6 M.
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35. Advanced pH Calculations
What is the pH of blood with [H+] = 4.8 × 10–8 M?
pH = –log[H+] = –log(4.8 × 10–8) = – (–7.32)
pH = 7.32
What is the [H+] in orange juice with a pH of 2.75?
[H+] = 10–pH = 10–2.75 = M
[H+] = 2.75 × 10–3 M
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36. Strong & Weak Electrolytes
An aqueous solution that is a good conductor of electricity is a strong electrolyte.
An aqueous solution that is a poor conductor of electricity is a weak electrolyte.
The greater the degree of ionization or dissociation, the greater the conductivity of the solution.
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37. Electrolyte Strength
Weak acids and bases and insoluble ionic compounds are weak electrolytes.
Strong acids and bases and soluble ionic compounds are strong electrolytes.
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38. H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
Total Ionic Equation
The concept of ionization allows us to portray ionic solutions more accurately by showing strong electrolytes in their ionized form.
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)
Write strong acids and bases and soluble ionic compounds as ions:
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
This is the total ionic equation. Each species is written as it predominantly exists in solution.
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39. Net Ionic Equation
H+(aq) + Cl-(aq) + Na+(aq) + OH-(aq) → Na+(aq) + Cl-(aq) + H2O(l)
Notice that Na+ and Cl- appear on both sides of the equation. They are spectator ions.
Spectator ions are in the solution, but do not participate in the overall reaction. We can cancel out the spectator ions to give the net ionic equation.
The net ionic equation is:
H+(aq) + OH-(aq) → H2O(l)
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40. Writing Net Ionic Equations
Complete and balance the non-ionized chemical equation.
Convert the non-ionized equation into the total ionic equation
Write strong electrolytes in the ionized form
Write weak electrolytes, water, and gases in the non-ionized form
Cancel all the spectator ions to obtain the net ionic equation.
If all species are eliminated, there is no reaction.

41. Conclusions
pH is a measure of the acidity of a solution. The typical range for pH is 0 to 14.
Neutral solutions have a pH of 7.
Below are some properties of acids and bases.
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42. Conclusions Continued
An Arrhenius acid is a substance that ionizes in water to produce hydrogen ions.
An Arrhenius base is a substance that ionizes in water to release hydroxide ions.
A Brønsted-Lowry acid is a substance that donates a hydrogen ion to any other substance. It is a proton donor.
A Brønsted-Lowry base is a substance that accepts a hydrogen ion. It is a proton acceptor
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43. Conclusions Continued
In an aqueous solution, [H+][OH-] = 1.0 × This is the ionization constant of water, Kw.
pH = –log[H+]
[H+] = 10–pH
Strong electrolytes are mostly dissociated in solution.
Weak electrolytes are slightly dissociated in solution.
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