Christopher G. Hamaker, Illinois State University, Normal IL© 2008, Prentice Hall
Chapter 15Acids and Bases
INTRODUCTORY CHEMISTRYINTRODUCTORY CHEMISTRYConcepts & Connections
Fifth Edition by Charles H. Corwin
Chapter 15 2
• 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.
Properties of Acids
Chapter 15 3
• 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.
Properties of Bases
Chapter 15 4
• 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)
Acid/Base Neutralization
Chapter 15 5
• 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.
The pH Scale
Chapter 15 6
• 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.
Acid/Base Classifications of Solutions
Chapter 15 7
• 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.
Buffers
Chapter 15 8
• 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.
Arrhenius Acids and Bases
Chapter 15 9
• 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.
Strengths of Acids
Chapter 15 10
• 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.
Strengths of Bases
Chapter 15 11
• 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 and Weak Arrhenius Acids
Chapter 15 12
• 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.
Arrhenius Acids in Solution
Chapter 15 13
• 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
Strong and Weak Arrhenius Bases
Chapter 15 14
• 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%)
Arrhenius Bases in Solution
Chapter 15 15
Chemistry Connection: Svante Arrhenius
• Svante Arrhenius noted that NaCl solutions conducted electricity; while sugar solutions did not.
• He also noticed that the freezing point of NaCl solutions were lowered twice as much as sugar solutions at the same concentration.
• He proposed that NaCl produces ions when dissolved, while sugar was in solution as molecules.
• It was nearly 20 years before his ideas were accepted by the scientific community.
Chapter 15 16
• 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)
Neutralization Reactions
Chapter 15 17
• 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; the sulfate must be from sulfuric acid, H2SO4.
– H2SO4(aq) + Ca(OH)2(aq) → CaSO4(aq) + 2 H2O(l)
Predicting Neutralization Reactions
Chapter 15 18
• 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.
Brønsted-Lowry Acids and Bases
Chapter 15 19
• Let’s 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.
Brønsted-Lowry Acids and Bases
Chapter 15 20
• 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.
Amphiprotic Compounds
Chapter 15 21
• A solution that changes color as the pH changes is an acid-base indicator.
• Shown below are the three indicators at different pH values.
Methyl Red Bromothymol Blue Phenolphthalein
Acid-Base Indicators
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.
Acid-Base Titrations
Chapter 15 23
• 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?
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
Titration Problem
Chapter 15 24
• The molarity of NaOH can be written as the unit factor 0.233 mol NaOH / 1000 mL solution.
= 0.00837 mol HC2H3O2
37.55 mL solution ×1 mol HC2H3O2
1 mol NaOH
0.233 mol NaOH
1000 mL solution×
= 0.837 M HC2H3O2
1000 mL solution
10.0 mL solution
0.00837 mol HC2H3O2
1 L solution×
Titration Problem, continued
Chapter 15 25
• A 10.0 mL sample of 0.555 M H2SO4 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 H2SO4.
• Use 0.555 mol H2SO4/1000 mL and 0.233 mol NaOH/1000 mL.
Another Titration Problem
Chapter 15 26
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 0.233 M NaOH is required to neutralize 10.0 mL of 0.555 M H2SO4.
Titration Problem, continued
Chapter 15 27
• 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 Na2CO3?
2 HCl(aq) + Na2CO3(aq) → 2 NaCl(aq) + H2O(l) + CO2(g)
Acid-Base Standardization
Chapter 15 28
= 0.277 M HCl×25.50 mL solution
0.00708 mol HCl 1000 mL solution
1 L solution
= 0.00708 mol HCl
0.375 g Na2CO3 × ×1 mol Na2CO3
105.99 g Na2CO3
2 mol HCl
1 mol Na2CO3
Standardization, continued
Chapter 15 29
• 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 an ion, 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.
Ionization of Water
Chapter 15 30
• 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.
Autoionization of Water
Chapter 15 31
• 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
[H+] and [OH-] Relationship
Chapter 15 32
• 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 10 to express the hydrogen ion concentration.
• Mathematically: pH = –log[H+]
– [H+] is the molar hydrogen ion concentration
The pH Concept
Chapter 15 33
• 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.
Calculating pH
Chapter 15 34
• 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.
Calculating [H+] from pH
Chapter 15 35
• What is the pH of blood with [H+] = 4.8 × 10–8 M?
– pH = –log[H+] = –log(4.8 × 10–8) = – (–3.82)
– pH = 3.82
• 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
Advanced pH Calculations
Chapter 15 36
Critical Thinking: Acid Rain• Nitrogen oxides and sulfur oxides, produced from
the combustion of fossil fuels, react with rainwater to produce nitric and sulfuric acids.
• These strong acids reduce the pH of rainwater.
• Acid rain refers to rain with a pH below 5.
• Acid rain can lower the pH of lakes and cause corrosion of metal and degradation of
limestone and marble statues.
Chapter 15 37
• 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.
Strong and Weak Electrolytes
Chapter 15 38
• Weak acids and bases are weak electrolytes.
• Strong acids and bases are strong electrolytes.
• Insoluble ionic compounds are weak electrolytes.
• Soluble ionic compounds are strong electrolytes.
Electrolyte Strength
Chapter 15 39
Strengths of Electrolytes
Chapter 15 40
• 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.
Total Ionic Equations
Chapter 15 41
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)
Net Ionic Equations
Chapter 15 42
• 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.
Writing Net Ionic Equations
Chapter 15 43
• 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:
Chapter Summary
Chapter 15 44
• 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.
Chapter Summary, continued
Chapter 15 45
• In an aqueous solution, [H+][OH-] = 1.0 × 10-14. 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.
Chapter Summary, continued
