1. Chapter 15 Acids and Bases
Lecture Presentation
Chapter 15 Acids and Bases
Sherril Soman
Grand Valley State University

2. Stomach Acid and Heartburn
The cells that line your stomach produce hydrochloric acid.
To kill unwanted bacteria
To help break down food
To activate enzymes that break down food
If the stomach acid backs up into your esophagus, it irritates those tissues, resulting in heartburn.
Acid reflux

3. Curing Heartburn
Mild cases of heartburn can be cured by neutralizing the acid in the esophagus.
Swallowing saliva, which contains bicarbonate ion
Taking antacids that contain hydroxide ions and/or carbonate ions

4. GERD Chronic heartburn is a problem for some people.
GERD (gastroesophageal reflux disease) is chronic leaking of stomach acid into the esophagus.
In people with GERD, the muscles separating the stomach from the esophagus do not close tightly, allowing stomach acid to leak into the esophagus.
Physicians diagnose GERD by attaching a pH sensor to the esophagus to measure the acidity levels of the fluids over time.

5. Properties of Acids Sour taste Ability to dissolve many metals
Ability to neutralize bases
Change blue litmus paper to red

6. Common Acids

7. Structures of Acids
Binary acids have acid hydrogens attached to a nonmetal atom.
HCl, HF

8. Structure of Acids
Oxyacids have acid hydrogens attached to an oxygen atom.
H2SO4, HNO3

9. Structure of Acids Carboxylic acids have COOH group.
HC2H3O2, H3C6H5O7
Only the first H in the formula is acidic.
The H is on the COOH.

10. Properties of Bases Taste bitter Feel slippery
alkaloids = plant product that is alkaline
often poisonous
Feel slippery
Ability to turn red litmus paper blue
Ability to neutralize acids

11. Common Bases

12. Indicators
Indicators are chemicals that change color depending on the solution’s acidity or basicity.
Many vegetable dyes are indicators.
Anthocyanins
Litmus
From Spanish moss
Red in acid, blue in base
Phenolphthalein
Found in laxatives
Red in base, colorless in acid

13. Definitions of Acids and Bases
Arrhenius definition
Based on H+ and OH-
Brønsted–Lowry definition
Based on reactions in which H+ is transferred
Lewis definition

14. Arrhenius Theory Acids: produce H+ ions in aqueous solution.
HCl(aq) → H+(aq) + Cl−(aq)

15. Hydronium Ion
The H+ ions produced by the acid are so reactive they cannot exist in water.
H+ ions are protons!
Instead, they react with water molecules to produce complex ions, mainly hydronium ion, H3O+.
H+ + H2O  H3O+
There are also minor amounts of H+ with multiple water molecules, H(H2O)n+.

16. Arrhenius Theory Bases: produce OH− ions in aqueous solution.
NaOH(aq) → Na+(aq) + OH(aq)

17. Arrhenius Acid–Base Reactions
The H+ from the acid combines with the OH− from the base to make a molecule of H2O.
It is often helpful to think of H2O as H—OH.
The cation from the base combines with the anion from the acid to make a salt.
acid + base → salt + water
HCl(aq) + NaOH(aq) → NaCl(aq) + H2O(l)

18. Problems with Arrhenius Theory
It does not explain why molecular substances, such as NH3, dissolve in water to form basic solutions, even though they do not contain OH– ions.
It does not explain how some ionic compounds, such as Na2CO3 or Na2O, dissolve in water to form basic solutions, even though they do not contain OH– ions.
It does not explain why molecular substances, such as CO2, dissolve in water to form acidic solutions, even though they do not contain H+ ions.
It does not explain acid–base reactions that take place outside aqueous solution.

19. Brønsted–Lowry Acid–Base Theory
It defines acids and bases based on what happens in a reaction.
Any reaction involving H+ (proton) that transfers from one molecule to another is an acid–base reaction, regardless of whether it occurs in aqueous solution or if there is OH− present.
All reactions that fit the Arrhenius definition also fit the Brønsted–Lowry definition.

20. Brønsted–Lowry Theory
The acid is an H+ donor.
The base is an H+ acceptor.
Base structure must contain an atom with an unshared pair of electrons.
In a Brønsted–Lowry acid–base reaction, the acid molecule donates an H+ to the base molecule.
H–A + :B  :A– + H–B+

21. Brønsted–Lowry Acids Brønsted–Lowry acids are H+ donors.
Any material that has H can potentially be a Brønsted–Lowry acid.
Because of the molecular structure, often one H in the molecule is easier to transfer than others.
When HCl dissolves in water, the HCl is the acid because HCl transfers an H+ to H2O, forming H3O+ ions.
Water acts as base, accepting H+.
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
acid base

22. Brønsted–Lowry Bases Brønsted–Lowry bases are H+ acceptors.
Any material that has atoms with lone pairs can potentially be a Brønsted–Lowry base.
Because of the molecular structure, often one atom in the molecule is more willing to accept H+ transfer than others.
When NH3 dissolves in water, the NH3(aq) is the base because NH3 accepts an H+ from H2O, forming OH–(aq).
Water acts as acid, donating H+.
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)
base acid

23. Amphoteric Substances
Amphoteric substances can act as either an acid or a base because they have both a transferable H and an atom with lone pair electrons.
Water acts as base, accepting H+ from HCl.
HCl(aq) + H2O(l) → Cl–(aq) + H3O+(aq)
Water acts as acid, donating H+ to NH3.
NH3(aq) + H2O(l)  NH4+(aq) + OH–(aq)

24. Brønsted–Lowry Acid–Base Reactions
One of the advantages of Brønsted–Lowry theory is that it illustrates reversible reactions to be as follows:
H–A + :B  :A– + H–B+
The original base has an extra H+ after the reaction, so it will act as an acid in the reverse process.
And the original acid has a lone pair of electrons after the reaction, so it will act as a base in the reverse process:
:A– + H–B+  H–A + :B

25. Conjugate Acid–Base Pairs
In a Brønsted-–Lowry acid–base reaction,
the original base becomes an acid in the reverse reaction.
the original acid becomes a base in the reverse process.
Each reactant and the product it becomes is called a conjugate pair.

26. Conjugate Pairs A base accepts a proton and becomes a conjugate acid.
An acid donates a proton and becomes a conjugate base.

27. Arrow Conventions  in these notes
Chemists commonly use two kinds of arrows in reactions to indicate the degree of completion of the reactions.
A single arrow indicates that all the reactant molecules are converted to product molecules at the end.
A double arrow indicates the reaction stops when only some of the reactant molecules have been converted into products.
 in these notes

28. Strong or Weak A strong acid is a strong electrolyte.
Practically all the acid molecules ionize.
A strong base is a strong electrolyte.
Practically all the base molecules form OH– ions, either through dissociation or reaction with water.
A weak acid is a weak electrolyte.
Only a small percentage of the molecules ionize, .
A weak base is a weak electrolyte
only a small percentage of the base molecules form OH– ions, either through dissociation or reaction with water, .

29. Strong Acids Strong acids donate practically all their H’s.
100% ionized in water
Strong electrolyte
[H3O+] = [strong acid]
[X] means the molarity of X

30. Weak Acids Weak acids donate a small fraction of their H’s.
Most of the weak acid molecules do not donate H to water.
Much less than 1% ionized in water
[H3O+] << [weak acid]

31. Examples of Strong Acids

32. Examples of Weak Acids

33. Strengths of Acids and Bases
Commonly, acid or base strength is measured by determining the equilibrium constant of a substance’s reaction with water.
HAcid + H2O  Acid− + H3O+
Base: + H2O  HBase+ + OH−
The farther the equilibrium position lies toward the products, the stronger the acid or base.
The position of equilibrium depends on the strength of attraction between the base form and the H+.
Stronger attraction means stronger base or weaker acid.

34. Ionic Attraction and Acid Strength

35. General Trends in Acidity
The stronger an acid is at donating H, the weaker the conjugate base is at accepting H.
Higher oxidation number = stronger oxyacid
H2SO4 > H2SO3; HNO3 > HNO2
Cation stronger acid than neutral molecule; neutral stronger acid than anion
H3O+ > H2O > OH−; NH4+ > NH3 > NH2−
Trend in base strength opposite

36. Autoionization of Water
Water is amphoteric; it can act either as an acid or a base.
Therefore, there must be a few ions present.
About 2 out of every 1 billion water molecules form ions through a process called autoionization.
H2O Û H+ + OH–
H2O + H2O Û H3O+ + OH–
All aqueous solutions contain both H3O+ and OH–.
The concentration of H3O+ and OH– are equal in water.
[H3O+] = [OH–] = 10−7M at 25 °C

37. Ion Product of Water
The product of the H3O+ and OH– concentrations is always the same number.
The number is called the ion product of water and has the symbol Kw.
Also know as the dissociation constant of water
[H3O+] × [OH–] = Kw = 1.00 × 10−14 at 25 °C
If you measure one of the concentrations, you can calculate the other.
As [H3O+] increases the [OH–] must decrease so the product stays constant.
Inversely proportional

38. Acid Ionization Constant, Ka
Acid strength is measured by the size of the equilibrium constant when it reacts with H2O.
The equilibrium constant for this reaction is called the acid ionization constant, Ka.
larger Ka = stronger acid

39. Table 15.5

40. Acidic and Basic Solutions
All aqueous solutions contain both H3O+ and OH– ions.
Neutral solutions have equal [H3O+] and [OH–].
[H3O+] = [OH–] = 1.00 × 10−7
Acidic solutions have a larger [H3O+] than [OH–].
[H3O+] > 1.00 × 10−7; [OH–] < 1.00 × 10−7
Basic solutions have a larger [OH–] than [H3O+].
[H3O+] < 1.00 × 10−7; [OH–] > 1.00 × 10−7

41. Measuring Acidity: pH
The acidity or basicity of a solution is often expressed as pH.
pH = −log[H3O+]
Exponent on 10 with a positive sign
pHwater = −log[10−7] = 7
Need to know the [H3O+] concentration to find pH
pH < 7 is acidic; pH > 7 is basic. pH = 7 is neutral.
[H3O+] = 10−pH

42. Sig., Figs., and Logs
When you take the log of a number written in scientific notation, the digits before the decimal point come from the exponent on 10, and the digits after the decimal point come from the decimal part of the number.
log(2.0 x 106) = log(106) + log(2.0)
= … =
Because the part of the scientific notation number that determines the significant figures is the decimal part, the sig. figs. are the digits after the decimal point in the log.
log(2.0 × 106) = 6.30

43. What Does the pH Number Imply?
The lower the pH, the more acidic the solution; the higher the pH, the more basic the solution.
1 pH unit corresponds to a factor of 10 difference in acidity.
Normal range of pH is 0 to 14.
pH 0 is [H3O+] = 1 M; pH 14 is [OH–] = 1 M.
pH can be negative (very acidic) or larger than 14 (very alkaline).

44. pOH
Another way of expressing the acidity/basicity of a solution is pOH.
pOH = −log[OH], [OH] = 10−pOH
pOHwater = −log[10−7] = 7
Need to know the [OH] concentration to find pOH
pOH < 7 is basic; pOH > 7 is acidic; pOH = 7 is neutral.
pH + pOH = 14.0

45. Relationship between pH and pOH
pH + pOH = at 25 °C.
You can use pOH to find the pH of a solution.

46. pK A way of expressing the strength of an acid or base is pK.
pKa = −log(Ka), Ka = 10−pKa
pKb = −log(Kb), Kb = 10−pKb
The stronger the acid, the smaller the pKa.
Larger Ka = smaller pKa
Because it is the –log
The stronger the base, the smaller the pKb.
Larger Kb = smaller pKb

47. [H3O+] and [OH−] in a Strong Acid or Strong Base Solution
There are two sources of H3O+ in an aqueous solution of a strong acid—the acid and the water.
There are two sources of OH− in an aqueous solution of a strong acid—the base and the water.
For a strong acid or base, the contribution of the water to the total [H3O+] or [OH−] is negligible.
The [H3O+]acid shifts the Kw equilibrium so far that [H3O+]water is too small to be significant.
Except in very dilute solutions, generally < 1 × 10−4 M

48. Finding pH of a Strong Acid or Strong Base Solution
For a monoprotic strong acid [H3O+] = [HAcid].
For polyprotic acids, the other ionizations can generally be ignored.
For H2SO4, the second ionization cannot be ignored.
0.10 M HCl has [H3O+] = 0.10 M and pH = 1.00
For a strong ionic base, [OH−] = (number OH−) × [Base].
For molecular bases with multiple lone pairs available, only one lone pair accepts an H; the other reactions can generally be ignored.
0.10 M Ca(OH)2 has [OH−] = 0.20 M and pH =

49. Finding the pH of a Weak Acid
There are also two sources of H3O+ in an aqueous solution of a weak acid—the acid and the water.
However, finding the [H3O+] is complicated by the fact that the acid only undergoes partial ionization.
Calculating the [H3O+] requires solving an equilibrium problem for the reaction that defines the acidity of the acid.
HAcid + H2O  Acid + H3O+

50. Percent Ionization
Another way to measure the strength of an acid is to determine the percentage of acid molecules that ionize when dissolved in water; this is called the percent ionization.
The higher the percent ionization, the stronger the acid.
Because [ionized acid]equil = [H3O+]equil

51. Relationship between [H3O+]equilibrium and [HA]initial
Increasing the initial concentration of acid results in increased [H3O+] at equilibrium.
Increasing the initial concentration of acid results in decreased percent ionization.
This means that the increase in [H3O+] concentration is slower than the increase in acid concentration.

52. Why Doesn’t the Increase in H3O+ Keep Up with the Increase in HA?
The reaction for ionization of a weak acid is as follows:
HA(aq) + H2O(l)  A−(aq) + H3O+(aq)
According to Le Châtelier’s principle, if we reduce the concentrations of all the (aq) components, the equilibrium should shift to the right to increase the total number of dissolved particles.
We can reduce the (aq) concentrations by using a more dilute initial acid concentration.
The result will be a larger [H3O+] in the dilute solution compared to the initial acid concentration.
This will result in a larger percent ionization.

53. Finding the pH of Mixtures of Acids
Generally, you can ignore the contribution of the weaker acid to the [H3O+]equil.
For a mixture of a strong acid with a weak acid, the complete ionization of the strong acid provides more than enough [H3O+] to shift the weak acid equilibrium to the left so far that the weak acid’s added [H3O+] is negligible.
For mixtures of weak acids, you generally only need to consider the stronger for the same reasons, as long as one is significantly stronger than the other and their concentrations are similar.

54. Strong Bases
The stronger the base, the more willing it is to accept H.
Use water as the standard acid.
For ionic bases, practically all units are dissociated into OH– or accept H’s.
Strong electrolyte
Multi-OH strong bases completely dissociated

55. Weak Bases
In weak bases, only a small fraction of molecules accept H’s.
Weak electrolyte
Most of the weak base molecules do not take H from water.
Much less than 1% ionization in water
[HO–] << [weak base]
Finding the pH of a weak base solution is similar to finding the pH of a weak acid.

56. Base Ionization Constant, Kb
Base strength is measured by the size of the equilibrium constant when it reacts with H2O
:Base + H2O  OH− + H:Base+
The equilibrium constant is called the base ionization constant, Kb.
Larger Kb = stronger base 56

57. Common Weak Bases
Table 15.8 page 721

58. Structure of Amines

59. Acid–Base Properties of Ions and Salts
Salts are water-soluble ionic compounds.
Salts that contain the cation of a strong base and an anion that is the conjugate base of a weak acid are basic.
NaHCO3 solutions are basic.
Na+ is the cation of the strong base NaOH.
HCO3− is the conjugate base of the weak acid H2CO3.
Salts that contain cations that are the conjugate acid of a weak base and an anion of a strong acid are acidic.
NH4Cl solutions are acidic.
NH4+ is the conjugate acid of the weak base NH3.
Cl− is the anion of the strong acid HCl.

60. Anions as Weak Bases
Every anion can be thought of as the conjugate base of an acid.
Therefore, every anion can potentially be a base.
A−(aq) + H2O(l)  HA(aq) + OH−(aq)
The stronger the acid, the weaker the conjugate base.
An anion that is the conjugate base of a strong acid is pH neutral.
Cl−(aq) + H2O(l)  HCl(aq) + OH−(aq)
An anion that is the conjugate base of a weak acid is basic.
F−(aq) + H2O(l)  HF(aq) + OH−(aq)

61. Relationship between Ka of an Acid and Kb of Its Conjugate Base
Many reference books only give tables of Ka values because Kb values can be found from them.
When you add equations, you multiply the K’s.

62. NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
Cations as Weak Acids
Some cations can be thought of as the conjugate acid of a weak base.
Others are the counterions of a strong base.
Therefore, some cations can potentially be acidic.
MH+(aq) + H2O(l)  MOH(aq) + H3O+(aq)
The stronger the base, the weaker the conjugate acid.
A cation that is the counterion of a strong base is pH neutral.
A cation that is the conjugate acid of a weak base is acidic.
NH4+(aq) + H2O(l)  NH3(aq) + H3O+(aq)
Because NH3 is a weak base, the position of this equilibrium favors the right.

63. Metal Cations as Weak Acids
Cations of small, highly charged metals are weakly acidic.
Alkali metal cations and alkali earth metal cations are pH neutral.
Cations are hydrated.
Al(H2O)63+(aq) + H2O(l)  Al(H2O)5(OH)2+ (aq) + H3O+(aq)

64. Classifying Salt Solutions as Acidic, Basic, or Neutral
If the salt cation is the counterion of a strong base and the anion is the conjugate base of a strong acid, it will form a neutral solution.
NaCl Ca(NO3)2 KBr
If the salt cation is the counterion of a strong base and the anion is the conjugate base of a weak acid, it will form a basic solution.
NaF Ca(C2H3O2) KNO2

65. Classifying Salt Solutions as Acidic, Basic, or Neutral
If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a strong acid, it will form an acidic solution.
NH4Cl
If the salt cation is a highly charged metal ion and the anion is the conjugate base of a strong acid, it will form an acidic solution.
Al(NO3)3

66. Classifying Salt Solutions as Acidic, Basic, or Neutral
If the salt cation is the conjugate acid of a weak base and the anion is the conjugate base of a weak acid, the pH of the solution depends on the relative strengths of the acid and base.
NH4F because HF is a stronger acid than NH4+, Ka of NH4+ is larger than Kb of the F−; therefore, the solution will be acidic.

67. Ionization in Polyprotic Acids
Because polyprotic acids ionize in steps, each H has a separate Ka.
Ka1 > Ka2 > Ka3
Generally, the difference in Ka values is great enough so that the second ionization does not happen to a large enough extent to affect the pH.
Most pH problems just do first ionization.
Except H2SO4  uses [H2SO4] as the [H3O+] for the second ionization.
[A2−] = Ka2 as long as the second ionization is negligible.

69. Ionization in H2SO4 The ionization constants for H2SO4 are as follows:
H2SO4 + H2O  HSO4 + H3O+ strong
HSO4 + H2O  SO42 + H3O+ Ka2 = 1.2 × 10−2
For most sulfuric acid solutions, the second ionization is significant and must be accounted for.
Because the first ionization is complete, use the given [H2SO4] = [HSO4−]initial = [H3O+]initial.

70. Strengths of Binary Acids
The more d+ H─X d− polarized the bond, the more acidic the bond.
The stronger the H─X bond, the weaker the acid.
Binary acid strength increases to the right across a period.
Acidity: H─C < H─N < H─O < H─F
Binary acid strength increases down the column.
Acidity: H─F < H─Cl < H─Br < H─I

71. Relationship between Bond Strength and Acidity
Bond Energy
kJ/mol
Type of Acid HF
565
weak
HCl
431
strong
HBr
364

72. Strengths of Oxyacids, H–O–Y
The more electronegative the Y atom, the stronger the oxyacid.
HClO > HIO
Acidity of oxyacids decreases down a group.
Same trend as binary acids
Helps weaken the H–O bond.
The larger the oxidation number of the central atom, the stronger the oxyacid.
H2CO3 > H3BO3
Acidity of oxyacids increases to the right across a period.
Opposite trend of binary acids
The more oxygens attached to Y, the stronger the oxyacid.
Further weakens and polarizes the H–O bond
HClO3 > HClO2

73. Relationship between Electronegativity and Acidity
H─O─Y
Electronegativity of Y Ka
H─O─Cl
3.0
2.9 × 10−8
H─O─Br
2.8
2.0 × 10−9
H─O─I
2.5
2.3 × 10−11

74. Relationship between Number of Oxygens on the Central Atom and Acidity

75. Lewis Acid–Base Theory
Lewis acid–base theory focuses on transferring an electron pair.
Lone pair  bond
Bond  lone pair
Does NOT require H atoms
The electron donor is called the Lewis base.
Electron rich; therefore nucleophile
The electron acceptor is called the Lewis acid.
Electron deficient; therefore electrophile

76. Anions are better Lewis bases than neutral atoms or molecules.
The Lewis base has electrons it is willing to give away to or share with another atom.
The Lewis base must have a lone pair of electrons on it that it can donate.
Anions are better Lewis bases than neutral atoms or molecules.
N: < N:−
Generally, the more electronegative an atom, the less willing it is to be a Lewis base.
O: < S: 76

77. Lewis Acids
They are electron deficient, either from being attached to electronegative atom(s) or not having an octet.
They must have an empty orbital willing to accept the electron pair.
H+ has empty 1s orbital.
B in BF3 has empty 2p orbital and an incomplete octet.
Many small, highly charged metal cations have empty orbitals they can use to accept electrons.
Atoms that are attached to highly electronegative atoms and have multiple bonds can be Lewis acids.

78. Lewis Acid–Base Reactions
The base donates a pair of electrons to the acid.
It generally results in a covalent bond forming
H3N: + BF3  H3N─BF3
The product that forms is called an adduct.
Arrhenius and Brønsted–Lowry acid–base reactions are also Lewis.

79. Examples of Lewis Acid–Base Reactions

80. Examples of Lewis Acid–Base Reactions
Ag+(aq) :NH3(aq)  Ag(NH3)2+(aq)
Lewis
Acid
Lewis
Base
Adduct

81. CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
U.S. Fuel Consumption
Over 85% of the energy use in the United States comes from the combustion of fossil fuels.
Oil, natural gas, coal
Combustion of fossil fuels produces CO2.
CH4(g) + 2 O2(g) → CO2(g) + 2 H2O(g)
Natural fossil fuels also contain small amounts of S that burn to produce SO2(g).
S(s) + O2(g) → SO2(g)
The high temperatures of combustion allow N2(g) in the air to combine with O2(g) to form oxides of nitrogen.
N2(g) + 2 O2(g) → 2 NO2(g)

83. What Is Acid Rain? Natural rain water has a pH of 5.6.
Naturally slightly acidic due mainly to CO2
Rain water with a pH lower than 5.6 is called acid rain.
Acid rain is linked to damage in ecosystems and structures.

84. What Causes Acid Rain?
Many natural and pollutant gases dissolved in the air are nonmetal oxides.
CO2, SO2, NO2
Nonmetal oxides are acidic.
CO2(g) + H2O(l)  H2CO3(aq)
2 SO2(g) + O2(g) + 2 H2O(l)  2 H2SO4(aq)
4 NO2(g) + O2(g) + 2 H2O(l)  4 HNO3(aq)
Processes that produce nonmetal oxide gases as waste increase the acidity of the rain.
Natural – volcanoes and some bacterial action
Man made – combustion of fuel

85. pH of Rain in Different Regions
Figure pg 740

86. Weather Patterns
The prevailing winds in the United States travel west to east.
Weather patterns may cause rain to be acidic in regions other than where the nonmetal oxide is produced.
Much of the Northeast United States has rain of very low pH, even though it has very low sulfur emissions, due in part to the general weather patterns.

87. Damage from Acid Rain
Acids react with metals and materials that contain carbonates.
Acid rain damages bridges, cars, and other metallic structures.
Acid rain damages buildings and other structures made of limestone or cement.
Acidifying lakes affects aquatic life.
Soil acidity causes more dissolving of minerals and leaching more minerals from soil, making it difficult for trees.

88. Damage from Acid Rain

89. Acid Rain Legislation 1990 Clean Air Act attacks acid rain.
Forces utilities to reduce SO2
The result is acid rain in the Northeast stabilized and beginning to be reduced.