1. Acids and Bases
Ch 19

2. Intro to Acids and bases
19.1
Intro to Acids and bases

3. 19.1 Introduction to Acids and Bases
Identify the physical and chemical properties of acids and bases.
Classify solutions as acidic, basic, or neutral.
Compare the Arrhenius, Brønsted-Lowry, and Lewis models of acids and bases.
Lewis structure: a model that uses electron-dot structures to show how electrons are arranged in molecules

4. 19.1 Introduction to Acids and Bases (cont.)
acidic solution
basic solution
Arrhenius model
Brønsted-Lowry model
conjugate acid
conjugate base
conjugate acid-base pair
amphoteric
Lewis model
Different models help describe the behavior of acids and bases.

5. 19.1 Properties of Acids and Bases
Acids taste sour. Bases taste bitter and feel slippery.
Acids and bases are conductors of electricity.

6. 19.1 Properties of Acids and Bases
Acids and bases can be identified by their reactions with some metals and metal carbonates.
Acids react with metals to form H2 gas.
They react with carbonates and bicarbonates to produce a salt, water, and carbon dioxide gas
Geologists identify limestone because it produces bubbles of carbon dioxide when exposed to hydrochloric acid.

7. 19.1 Properties of Acids and Bases (cont.)
Acids turn blue litmus red.
Bases turn red litmus blue.

8. 19.1 Properties of Acids and Bases (cont.)
All water solutions contain hydrogen ions (H+) and hydroxide ions (OH–).
An acidic solution contains more hydrogen ions than hydroxide ions.
A basic solution contains more hydroxide ions than hydrogen ions.

9. 19.1 Properties of Acids and Bases (cont.)
The usual solvent for acids and bases is water—water produces equal numbers of hydrogen and hydroxide ions in a process called self-ionization.
H2O(l) + H2O(l) ↔ H3O+(aq) + OH–(aq)
The hydronium ion is H3O+.

10. 19.1 The Arrhenius Model
The Arrhenius model states that an acid is a substance that contains hydrogen and ionizes to produce hydrogen ions in aqueous solution, and a base is a substance that contains a hydroxide group and dissociates to produce a hydroxide ion in solution.
acid--donates a hydrogen ion (H+) in water
base--donates a hydroxide ion in water (OH-)
This theory was limited to substances with those "parts";
ammonia (NH3+) is a MAJOR exception!

11. 19.1 The Arrhenius Model (cont.)
Arrhenius acids and bases
HCl ionizes to produce H+ ions.
HCl(g) → H+(aq) + Cl–(aq)
NaOH dissociates to produce OH– ions.
NaOH(s) → Na+(aq) + OH–(aq)
Some solutions produce hydroxide ions even though they do not contain a hydroxide group.

12. 19.1 The Brønsted-Lowry Model
The Brønsted-Lowry Model of acids and bases states that an acid is a hydrogen ion donor, and a base is a hydrogen ion acceptor.
The Brønsted-Lowry Model is a more inclusive model of acids and bases.
acid--donates a proton in water
base--accepts a proton in water
This theory is better; it explains ammonia as a base! This is the main theory that we will use for our acid/base discussion.

14. 19.1 The Brønsted-Lowry Model (cont.)
A conjugate acid is the species produced when a base accepts a hydrogen ion.
A conjugate base is the species produced when an acid donates a hydrogen ion.
A conjugate acid-base pair consists of two substances related to each other by donating and accepting a single hydrogen ion.

15. 19.1 The Brønsted-Lowry Model (cont.)
Hydrogen fluoride—a Brønsted-Lowry acid
HF(aq) + H2O(l) ↔ H3O+(aq) + F–(aq)
HF = acid,
H2O = base,
H3O+ = conjugate acid,
F– = conjugate base

16. 19.1 The Brønsted-Lowry Model (cont.)
Ammonia— Brønsted-Lowry base
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH–(aq)
NH3 = base,
H2O(l) = acid,
NH4+ = conjugate acid,
OH– = conjugate base

17. 19.1 The Brønsted-Lowry Model (cont.)
Water and other substances that can act as acids or bases are called amphoteric.

18. 19.1 Monoprotic and Polyprotic Acids
An acid that can donate only one hydrogen ion is a monoprotic acid.
Only ionizable hydrogen atoms can be donated.

19. 19.1 Monoprotic and Polyprotic Acids (cont.)
Acids that can donate more than one hydrogen ion are polyprotic acids.

20. 19.1 The Lewis Model
According to the Lewis model, a Lewis acid is an electron-pair acceptor and a Lewis base is an electron pair donor.
The Lewis model includes all the substances classified as Brønsted-Lowry acids and bases and many more.

21. 19.1 Check A B C D A Lewis acid is a(n) ____. A. electron pair donor
B. hydrogen ion donor
C. electron pair acceptor
D. substance that contains an hydroxide group A B C D

22. 19.1 Check A B C D A conjugate acid is formed when:
A. a base accepts a hydrogen ion
B. an acid accepts a hydrogen ion
C. an acid donates a hydrogen ion
D. a base donates a hydrogen ion A B C D

23. Strengths of Acids and bases
19.2
Strengths of Acids and bases

24. 19.2 Strengths of Acids and Bases
Relate the strength of an acid or base to its degree of ionization.
electrolyte: an ionic compound whose aqueous solution conducts an electric current
Compare the strength of a weak acid with the strength of its conjugate base.
Explain the relationship between the strengths of acids and bases and the values of their ionization constants.

25. 19.2 Strengths of Acids and Bases (cont.)
strong acid
weak acid
strong base
weak base
In solution, strong acids and bases ionize completely, but weak acids and bases ionize only partially.

26. Review – Naming Acids and Bases
Ion
Name
Example
-ide (no oxygens)
Hydro -
+ -ic
HCl
-ite (less oxygens)
-ous
H2SO3
-ate (more oxygens)
-ic
H2SO4
Hydroxide (OH-)
hydroxide
NaOH
Hydrochloric acid
Sulfurous acid
Sulfuric acid
Sodium hydroxide

27. Review – Naming Acids and Bases
The most common oxyacids that are helpful to memorize are the following: H2SO4 : ______________________
HNO3: ______________________ H3PO4: ______________________
HClO3 : ______________________
Sulfuric acid
Nitric acid
Phosphoric acid
Chloric acid

28. Relative Strengths of Acids and Bases
Strength is determined by the amount of "dissociation”.
How much separates into ions in water?
Do Not
confuse concentration
with strength!

29. 19.2 Strengths of Acids Acids that ionize completely are strong acids.
With a strong acid, the conjugate base is a weak base.
Because they produce the maximum number of hydrogen ions, strong acids are good conductors of electricity.

30. Strong Acids hydrohalic acids: HCl, HBr, HI (except HF) nitric: HNO3
sulfuric: H2SO4
perchloric HClO4
chloric: HClO3

31. 19.2 Strengths of Acids (cont.)
Acids that ionize only partially in dilute aqueous solutions are called weak acids.

32. Strong
Weak

33. 19.2 Strengths of Acids (cont.)

34. 19.2 Strengths of Bases
A base that dissociates completely into metal ions and hydroxide ions is known as a strong base.
A weak base ionizes only partially in dilute aqueous solution.

35. oxides of IA and IIA metals
Strong Bases
Hydroxides OR
oxides of IA and IIA metals

36. 19.2 Strengths of Bases (cont.)

37. 19.3 Hydrogen Ions and pH Explain pH and pOH.
Le Châtelier’s principle: states that if a stress is applied to a system at equilibrium, the system shifts in the direction that relieves the stress
Relate pH and pOH to the ion product constant for water.
Calculate the pH and pOH of aqueous solutions.
ion product constant for water pH
pOH
pH and pOH are logarithmic scales that express the concentrations of hydrogen ions and hydroxide ions in aqueous solutions.

38. 19.3 Ion Product Constant for Water
Pure water contains equal concentrations of H+ and OH– ions.
The ion production of water, Kw = [H+][OH–].
The ion product constant for water is the value of the equilibrium constant expression for the self-ionization of water.

39. 19.3 Ion Product Constant for Water (cont)
With pure water at 298 K, both [H+] and [OH–] are equal to 1.0 × 10–7M.
Kw at 298 K = 1.0 × 10–14
Kw and LeChâtelier’s Principle proves [H+] × [OH–] must equal 1.0 × 10–14 at 298 K, and as [H+] goes up, [OH–] must go down.

40. Ion-Product constant for water Kw
Kw  =  [H+] [OH-]   =   1.0 X  10-14  mol/L
acid solution [H+]  >   [OH-]
basic solution  [OH‑]   >   [H+]
Square bracket indicates molar concentrations

41. 19.3 pH and pOH
Concentrations of H+ ions are often small numbers expressed in exponential notation.
pH is the negative logarithm of the hydrogen ion concentration of a solution. pH = –log [H+]

42. 19.3 pH and pOH (cont.)
pOH of a solution is the negative logarithm of the hydroxide ion concentration.
pOH = –log [OH–]
The sum of pH and pOH equals 14.

43. 19.3 pH and pOH (cont.) pH scale ranges from 0-14
pH-negative logarithm of the hydrogen ion concentration
pH   =  -log [H+]
pOH = -log [OH-]
pH 7  =  neutral
[H+]   =   1.0  X   10-7 mol/L
[OH-]  =  1.0  X   10-7 mol /L
acid pH       0   7         
[H+]   >   1.0   X  10-7  mol/L
base pH       7   14    
[OH-]  <   1.0  X  10-7 mol/L 7
base
acid 14
[H+] = [OH-]

44. 19.3 pH and pOH (cont.) pH + pOH= 14
You can find either if you have the other by subtracting from 14.

45. 19.3 pH and pOH (cont.)
For all strong monoprotic acids, the concentration of the acid is the concentration of H+ ions.
For all strong bases, the concentration of the OH– ions available is the concentration of OH–.
Weak acids and weak bases only partially ionize.
(equilibrium constant values must be used to evaluate the concentration and we will not do these specific calculations.)

46. 19.3 pH and pOH (cont.)
Litmus paper, pH paper, and a pH meter with electrodes can determine the pH of a solution.

47. Calculating [H3O+] and [OH-]
[H3O+]= 1.0 x 10-pH
Ex: pH=7.52
1.0 x =[H3O+]= 3.0 X 10-8
[OH-]= 1.0 x 10-pOH
Ex: pOH=6.3
1.0 x =[OH-]=5.0 x 10-7

48. Finding (H3O+) or (OH-) (knowing kw= [H3O+][OH-] )
If you want H3O+ and you have OH-,
1.0 x 10-14= [H3O+ ]
[OH-]
If you want OH- and you have H3O+
1.0 x 10-14= [OH-]
[H3O+]

49. 19.3 Check A B C D In dilute aqueous solution, as [H+] increases:
A. pH decreases
B. pOH increases
C. [OH–] decreases
D. all of the above A B C D

50. 19.3 Check
What is the pH of a neutral solution such as pure water?
A. 0
B. 7
C. 14
D. 1.0 × 10–14 A B C D

51. 19.4 Neutralization
Write chemical equations for neutralization reactions.
stoichiometry: the study of quantitative relationships between the amounts of reactants used and products formed by a chemical reaction; is based on the law of conservation of mass
Explain how neutralization reactions are used in acid-base titrations.
Compare the properties of buffered and unbuffered solutions.

52. 19.4 Neutralization (cont.)
neutralization reaction
salt
titration
titrant
equivalence point
acid-base indicator
end point
salt hydrolysis
buffer
buffer capacity
In a neutralization reaction, an acid reacts with a base to produce a salt and water.

53. 19.4 Reactions Between Acids and Bases
A neutralization reaction is a reaction in which an acid and a base in an aqueous solution react to produce a salt and water.
A salt is an ionic compound made up of a cation from a base and an anion from an acid.
Neutralization is a double-replacement reaction.

54. 19.4 Neutralization (cont.)
Reaction in which an acid and a base react in an aqueous solution to produce a salt and water.
*These are double-replacement reactions.
Salt- compound consisting of an anion from an acid and a cation from a base.
Ex: HCl + KOH → KCl + H2O
acid base salt water
HOH = H2O

55. 19.4 Reactions Between Acids and Bases (cont.)

56. 19.4 Reactions Between Acids and Bases (cont.)
Titration is a method for determining the concentration of a solution by reacting a known volume of that solution with a solution of known concentration.

57. 19.4 Reactions Between Acids and Bases (cont.)
In a titration procedure, a measured volume of an acid or base of unknown concentration is placed in a beaker, and initial pH recorded.
A buret is filled with the titrating solution of known concentration, called a titrant.

58. 19.4 Titration Setup
Titration Demo

59. Examples of Titrants and Analytes

60. Example Titration

61. 19.4 Reactions Between Acids and Bases (cont.)
Measured volumes of the standard solution are added slowly and mixed into the solution in the beaker, and the pH is read and recorded after each addition. The process continues until the reaction reaches the equivalence point, which is the point at which moles of H+ ion from the acid equals moles of OH– ion from the base.
An abrupt change in pH occurs at the equivalence point.

62. 19.4 Reactions Between Acids and Bases (cont.)

63. 19.4 Reactions Between Acids and Bases (cont.)
Chemical dyes whose color are affected by acidic and basic solutions are called acid-base indicators.

64. Section 18-4
An end point is the point at which an indicator used in a titration changes color.
An indicator will change color at the equivalence point.

65. 19.4 Salt Hydrolysis
In salt hydrolysis, the anions of the dissociated salt accept hydrogen ions from water or the cations of the dissociated salt donate hydrogen ions to water.

66. 19.4 Salt Hydrolysis (cont.)
Salts that produce basic solutions
KF is the salt of a strong base (KOH) and a weak acid (HF).
KF(s) → K+(aq) + F–(aq)

67. 19.4 Salt Hydrolysis (cont.)
Salts that produce acidic solutions
NH4Cl is the salt of a weak base (NH3) and strong acid (HCl).
When dissolved in water, the salt dissociates into ammonium ions and chloride ions.
NH4Cl(s) → NH4+(aq) + Cl–(aq)

68. 19.4 Salt Hydrolysis (cont.)
Salts that produce neutral solutions
NaNO3 is the salt of a strong acid (HNO3) and a strong base (NaOH).
Little or no salt hydrolysis occurs because neither Na+ nor NO3– react with water.

69. 19.4 Buffered Solutions
The pH of blood must be kept in within a narrow range.
Buffers are solutions that resist changes in pH when limited amounts of acid or base are added.

70. 19.4 Buffered Solutions (cont.)
Ions and molecules in a buffer solution resist changes in pH by reacting with any hydrogen ions of hydroxide ions added to the buffered solution.
HF(aq) ↔ H+(aq) + F–(aq)
When acid is added, the equilibrium shifts to the left.

71. 19.4 Buffered Solutions (cont.)
Additional H+ ions react with F– ions to form undissociated HF molecules but the pH changes little.
The amount of acid or base that a buffer solution can absorb without a significant change in pH is called the buffer capacity.

72. 19.4 Buffered Solutions (cont.)
A buffer is most effective when the concentrations of the conjugate acid-base pair are equal or nearly equal.

75. 19.4 Check
In a neutralization reaction, an acid and base react to form:
A. salt and oxygen gas
B. salt and ammonia
C. salt and water
D. precipitate and water A B C D

76. 19.4 Check
Solutions that resist changes in pH are called ____.
A. titrants
B. salts
C. conjugate pairs
D. buffers A B C D