• 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.

• 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.

• 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.

• 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) + H
2
O(l)
→
H
3
O + (aq) + Cl
–
(aq) (~100%)
HC
2
H
3
O
2
(aq) + H
2
O(l)
→
H
3
O + (aq) + C
2
H
3
O
2
–
(aq) (~1%)
• The hydronium ion , H
3
O + , is formed when the aqueous hydrogen ion attaches to a water molecule.

• 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%)
NH
4
OH(aq)
→
NH
4
+ (aq) + OH
–
(aq) (~1%)

• 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 .

• Lets look at two acid/base reactions:
– HCl(aq) + NaOH(aq) →
NaCl(aq) + H
2
O(l)
– HCl(aq) + NH
3
(aq)
→
NH
4
Cl(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, NH
3 is the Brønsted-Lowry base.
accepts a proton and

• 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.

• Weak acids and bases and insoluble ionic compounds are weak electrolytes.
• Strong acids and bases and soluble ionic compounds are strong electrolytes.

• 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.

• 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.

• 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%

• 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.
– NH
4
OH is a weak base and only partially dissociates

• A substance that is capable of both donating and accepting a proton is an amphiprotic compound.
• NaHCO
3 is an example:
– HCl(aq) + NaHCO
3
(aq)
→
NaCl(aq) + H
2
CO
3
(aq)
– NaOH(aq) + NaHCO
3
(aq)
→
Na
2
CO
3
(aq) + H
2
O(l)
• NaHCO
3 accepts a proton from HCl in the first reaction and donates a proton to NaOH in the second reaction.

• 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.

• 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.
18

• Recall, an acid neutralizes a base to produce a salt and water.
– HCl(aq) + NaOH(aq) →
NaCl(aq) + H
2
O(l)
• The reaction produces the aqueous salt NaCl.
• If we have an acid with two hydrogens (sulfuric acid, H
2
SO
4
), we need two hydroxide ions to neutralize it.
– H
2
SO
4
(aq) + 2 NaOH(aq)
→
Na
2
SO
4
(aq) + 2 H
2
O(l)

• We can identify the Arrhenius acid and base that react in a neutralization reaction to produce a given salt such as calcium sulfate, CaSO
4
.
• The calcium must be from calcium hydroxide,
Ca(OH)
2
, and the sulfate must be from sulfuric acid, H
2
SO
4
.
– H
2
SO
4
(aq) + Ca(OH)
2
(aq)
→
CaSO
4
(aq) + 2 H
2
O(l)

• 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.
Chapter 15 21

• Shown below are the three indicators at different pH values
Methyl Red Bromothymol Blue Phenolphthalein
Chapter 15 22

• 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.

• Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH. What is the concentration of the acetic acid?
HC
2
H
3
O
2
(aq) + NaOH(aq)
→
NaC
2
H
3
O
2
(aq) + H
2
O(l)
• We want concentration acetic acid, we have concentration sodium hydroxide.
conc NaOH
 mol NaOH
 mol HC
2
H
3
O
2
 conc HC
2
H
3
O
2

• The molarity of NaOH can be written as the unit factor 0.233 mol NaOH / 1000 mL solution.
0.233 mol NaOH
37.55 mL solution ×
1000 mL solution
×
1 mol HC
2
H
3
O
2
1 mol NaOH
= 0.00837 mol HC
2
H
3
O
2
0.00837 mol HC
2
H
3
O
2
10.0 mL solution
×
1000 mL solution
1 L solution
= 0.837 M HC
2
H
3
O
2

• A 10.0 mL sample of 0.555
M H
2
SO
4 is titrated with 0.233 M NaOH. What volume of NaOH is required for the titration?
• We want mL of NaOH, we have 10.0 mL of
H
2
SO
4
.
• Use 0.555 mol H
2
NaOH/1000 mL.
SO
4
/1000 mL and 0.233 mol

H
2
SO
4
(aq) + 2 NaOH(aq)
→
Na
2
SO
4
(aq) + H
2
O(l)
10.0 mL H
2
SO
4
×
0.555 mol H
2
SO
4 ×
1000 mL H
2
SO
4
1 mol H
2
SO
2 mol NaOH
4 ×
0.233 mol NaOH
1000 mL NaOH
= 49.8 mL NaOH
• 49.8 mL of 0.233
M NaOH is required to neutralize 10.0 mL of 0.555 M H
2
SO
4
.

• 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 25.50 mL are required to neutralize
0.375 g Na
2
CO
3
?
2 HCl(aq) + Na
2
CO
3
(aq)
→
2 NaCl(aq) + H
2
O(l) + CO
2
(g)

0.375 g Na
2
CO
3
×
1 mol Na
2
CO
3
105.99 g Na
2
CO
3
×
2 mol HCl
1 mol Na
2
CO
3
= 0.00708 mol HCl
0.00708 mol HCl
×
1000 mL solution
25.50 mL solution 1 L solution
= 0.277 M HCl

• Water undergoes an autoionization reaction . Two water molecules react to produce a hydronium ion and a hydroxide ion:
– H
2
O(l) + H
2
O(l)
→
H
3
O + (aq) + OH (aq) or
– H
2
O(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.

• 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:
– H
2
O(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,
K w
.

+
-
• At 25 
C, [H + ][OH ] = 1.0 × 10 -14 . 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

• 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

• What is the pH if the hydrogen ion concentration in a vinegar solution is 0.001 M ?
• pH = –log[H + ]
• pH = –log(0.001)
• pH = – ( –3) = 3
• The pH of the vinegar is 3, so the vinegar is acidic.

+
• 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
= 0.000001 M
• [H + ] = 1 × 10
–6
M .

• 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
= 0. 0018 M
– [H + ] = 2.75 × 10
–3
M

• 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.

• 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

• In an aqueous solution, [H + ][OH ] = 1.0 × 10 -14 .
This is the ionization constant of water, K w
.
• pH = –log[H + ]
• [H + ] = 10
–pH
• Strong acids and bases are strong electrolytes.
They are completely dissociated in solution.
• Weak acids and bases are weak electrolytes. They are slightly dissociated in solution.

• Consider the titration of acetic acid with sodium hydroxide. A 10.0 mL sample of acetic acid requires 37.55 mL of 0.223 M NaOH.
What is the concentration of the acetic acid?
• A 10.0 mL sample of 0.555
M H
2
SO
4 is titrated with 0.233 M NaOH. What volume of NaOH is required for the titration?
• What is the molarity of a hydrochloric acid solution if 25.50 mL are required to neutralize
0.375 g Na
2
CO
3
?