1. “Acids, Bases, and Salts”
Honors Chemistry

2. Acid-Base Theories OBJECTIVES:
Define the properties of acids and bases.

3. Acid-Base Theories
OBJECTIVES:
Compare and contrast acids and bases as defined by the theories of:
a) Arrhenius,
b) Brønsted-Lowry, and
c) Lewis.

4. Acid-Base in the Real World
Acids
How many foods can you think of that are sour?
That sour taste is due to the acid contained
in the food.

5. Examples Sour Milk –lactic acid Vinegar- acetic acid
Soda – phosphoric acid
Lemons and oranges – citric acid
Grape juice – tartaric acid

6. Acid-Base in the Real World
Bases
Many everyday items in the household
contain bases.
Cleaning products and simple medicines.

7. Examples Household cleaners – ammonia Drain and oven cleaners- NaOH
Antacids – magnesium hydroxide,
aluminum hydroxide and
sodium bicarbonate

8. Properties of Acids They taste sour. They can conduct electricity.
Can be strong or weak electrolytes in aqueous solution
React with metals to form H2 gas.
Change the color of indicators (for example: blue litmus turns to red).
React with bases to form water and a salt.

9. Properties of Acids
They have a pH of less than 7 (more on this concept of pH in a later lesson)
How do you know if a chemical is an acid?
It usually starts with Hydrogen.
HCl, H2SO4, HNO3, etc. (but not water!)

10. Acid Nomenclature A binary acid is an acid that contains
only two different elements: hydrogen
and one of the more electronegative
elements.
HF – hydrofluoric acid
HCl – hydrochloric acid
HBr – hydrobromic acid
HI – hydroiodic acid

11. Acid Nomenclature A oxyacid is an acid that contains
Hydrogen, oxygen and a third element, usually a nonmetal.
CH3COOH – acidic acid
HNO3 – nitric acid
H2SO4 – sulfuric acid
H3PO4 – phosphoric acid

12. Acids Affect Indicators, by changing their color
Blue litmus paper turns red in contact with an acid (and red paper stays red).

13. Acids have a pH less than 7

14. Acids React with Active Metals
Acids react with metals to form salts and hydrogen gas:
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)
This is a single-replacement reaction

15. Acids React with Carbonates and Bicarbonates
HCl + NaHCO3
Hydrochloric acid + sodium bicarbonate
NaCl + H2O + CO2
salt + water + carbon dioxide
An old-time home remedy for relieving an upset stomach

16. Acids Neutralize Bases
HCl + NaOH → NaCl + H2O
-Neutralization reactions ALWAYS produce a salt (which is an ionic compound) and water.

17. Effects of Acid Rain on Marble (marble is calcium carbonate)
George Washington:
BEFORE acid rain
George Washington:
AFTER acid rain

18. Properties of Bases React with acids to form water and a salt.
Taste bitter.
Feel slippery (don’t try this either).
Can be strong or weak electrolytes in aqueous solution
Change the color of indicators (red litmus turns blue).

19. Examples of Bases Sodium hydroxide, NaOH (lye for drain cleaner; soap)
Potassium hydroxide, KOH (alkaline batteries)
Magnesium hydroxide, Mg(OH)2 (Milk of Magnesia)
Calcium hydroxide, Ca(OH)2 (lime; masonry)

20. Bases Affect Indicators
Red litmus paper turns blue in contact with a base (and blue paper stays blue).

21. Bases have a pH greater than 7

22. Bases Neutralize Acids
Milk of Magnesia contains magnesium hydroxide, Mg(OH)2, which neutralizes stomach acid, HCl.
2 HCl + Mg(OH)2
Magnesium salts can cause diarrhea (thus they are used as a laxative) and may also cause kidney stones.
MgCl2 + 2 H2O

23. Acid-Base Theories

24. Svante Arrhenius
He was a Swedish chemist ( ), and a Nobel prize winner in chemistry (1903)
One of the first chemists to explain the chemical theory of the behavior of acids and bases

25. 1. Arrhenius Definition - 1887
Acids produce hydrogen ions (H1+) in water (HCl → H1+ + Cl1-)
Bases produce hydroxide ions (OH1-) when dissolved in water.
(NaOH → Na1+ + OH1-)
Limited to aqueous solutions.
Only one kind of base (hydroxides)
NH3 (ammonia) could not be an Arrhenius base: no OH1- produced.

26. Acids Also, not all the hydrogen in an acid may be released as ions
only those that have very polar bonds are ionizable - this is when the hydrogen is joined to a very electronegative element
(O, F, Br)

27. Arrhenius examples... Consider HCl = it is an acid!
What about CH4 (methane)?
CH3COOH (acetic acid) - it has 4 hydrogens just like methane but one hydrogen is bound to an oxygen atom.

28. 2. Brønsted-Lowry - 1923 A broader definition than Arrhenius
Acid is hydrogen-ion donor (H+ or proton); base is hydrogen-ion acceptor.
Acids and bases always come in pairs.
HCl is an acid.
When it dissolves in water, it gives it’s proton to water.
HCl(g) + H2O(l) ↔ H3O+(aq) + Cl-(aq)
Water is a base; makes hydronium ion.

29. Why Ammonia is a Base
Ammonia can be explained as a base by using Brønsted-Lowry:
NH3(aq) + H2O(l) ↔ NH41+(aq) + OH1-(aq)
Ammonia is the hydrogen ion acceptor (base), and water is the hydrogen ion donor (acid).
This causes the OH1- concentration to be greater than in pure water, and the ammonia solution is basic

30. Acids and bases come in pairs
A “conjugate base” is the remainder of the original acid, after it donates its hydrogen ion
A “conjugate acid” is formed when the original base gains a hydrogen ion
Thus, a conjugate acid-base pair is related by the loss or gain of a single hydrogen ion.

31. Acids and bases come in pairs
General equation is:
HA(aq) + H2O(l) ↔ H3O+(aq) + A-(aq)
Acid + Base ↔ Conjugate acid + Conjugate base
NH3 + H2O ↔ NH41+ + OH1-
base acid c.a c.b.
HCl + H2O ↔ H3O1+ + Cl1-
acid base c.a c.b.
Amphoteric – a substance that can act as both an acid and base- as water shows

32. 3. Lewis Acids and Bases
Gilbert Lewis focused on the donation or acceptance of a pair of electrons during a reaction
Lewis Acid - electron pair acceptor
Lewis Base - electron pair donor
Most general of all 3 definitions; acids don’t even need hydrogen!

33. Hydrogen Ions and Acidity
OBJECTIVES:
Describe how [H1+] and [OH1-] are related in an aqueous solution.

34. Hydrogen Ions and Acidity
OBJECTIVES:
Classify a solution as neutral, acidic, or basic given the hydrogen-ion or hydroxide-ion concentration.

35. Hydrogen Ions and Acidity
OBJECTIVES:
Convert hydrogen-ion [H+] concentrations into pH values and hydroxide-ion concentrations into pOH values.

36. Hydrogen Ions and Acidity
OBJECTIVES:
Describe the purpose of an acid-base indicator.

37. Hydrogen Ions from Water
Water ionizes, or falls apart into ions:
2H2O ↔ 2H+ + 2OH-
Called the “self ionization” of water
Occurs to a very small extent:
[H1+ ] = [OH1-] = 1 x 10-7 M
Since they are equal, a neutral solution results from water
Kw = [H1+ ] x [OH1-] = 1 x M2
Kw is called the “ionization constant” for water

38. Ion Product Constant 2H2O ↔ 2H+ + 2OH-
Kw is constant in every aqueous solution: [H+] x [OH-] = 1 x M
If [H+] > 10-7 then [OH-] < 10-7
If [H+] < 10-7 then [OH-] > 10-7
If we know one, other can be determined
If [H+] > 10-7 , it is acidic and [OH-] < 10-7
If [H+] < 10-7 , it is basic and [OH-] > 10-7
Basic solutions also called “alkaline”

39. Strong acids and bases are completely ionized in aqueous solutions:
HCl H Cl-
NaOH Na OH-

40. This is a basic solution
Therefore:
NaOH Na OH-
1 mole mole
1 x 10-2 M x 10-2 M
This is a basic solution

41. pH

42. The pH Scale – from 0 to 14 definition: pH = -log[H+]
pH = pouvoir hydrogene (Fr.) “hydrogen power”
definition: pH = -log[H+]
in neutral pH = -log(1 x 10-7) = 7
in acidic solution [H+] > 10-7
pH < -log(10-7)
pH < 7 (from 0 to 7 is the acid range)
pH > 7 (7 to 14 is base range)

43. Calculating pOH pOH = -log [OH-] [H+] x [OH-] = 1 x 10-14 M2
pH + pOH = 14
Not greatly used like pH is.

44. Problem 1 What is the pH of a solution if the [H3O+] is 1 x 10-6 M?
pH = -log[H3O+]
pH = -log(1 x 10-6) = 6
The solution is acidic

45. Problem 2 What is the pH of a solution if the [OH-] is 1 x 10-2 M?
[H+] x [OH-] = 1 x M2
[H+] x [1 x 10-2] = 1 x M2
[H+] = 1 x M
pH = -log[H3O+]

46. Problem 2 What is the pH of a solution if the [OH-] is 1 x 10-2 M?
pH = -log[H3O+]
pH = -log(1 x 10-12) = 12
The solution is basic

47. Problem 3 What is the pH of a solution if the [H3O+] is 3.4 x 10-5 M?
Hint: since 3.4 x 10-5 M is between 10-4 and 10-5 the pH of the solution must be between 4 and 5.

48. Problem 3 What is the pH of a solution if the [H3O+] is 3.4 x 10-5 M?
pH = -log[H3O+]
pH = -log[3.4 x 10-5 ]
pH = 4.47

49. Problems
1) What is the pH of a solution if the [H3O+] is 6.7 x 10-4 M?
2) What is the pH of a solution if the [H3O+] is 2.5 x 10-2 M?

50. Problems What is the pH of a solution if the [H3O+] is 6.7 x 10-4 M?
Answer: 3.17
2) What is the pH of a solution if the [H3O+] is 2.5 x 10-2 M?
Answer: 1.60

51. Problems
3) Determine the pH of a 2.0 x 10-2 M Sr(OH)2 solution.

52. Problems 3) Determine the pH of a 2.0 x 10-2 M Sr(OH)2 solution.
[OH-] = 4.0 x 10-2 M
[H+] = 2.5 x M
pH = -log[2.5 x 10-13] = 12.60

53. Measuring pH Why measure pH?
Everyday solutions we use - everything from swimming pools, soil conditions for plants, medical diagnosis, soaps and shampoos, etc.
Sometimes we can use indicators, other times we might need a pH meter

54. How to measure pH with wide-range paper
1. Moisten the pH indicator paper strip with a few drops of solution, by using a stirring rod.
2.Compare the color to the chart on the vial – then read the pH value.

55. Strengths of Acids and Bases
OBJECTIVES:
Define strong acids and weak acids.

56. Strength
Acids and Bases are classified according to the degree to which they ionize in water:
Strong are completely ionized in aqueous solution; this means they ionize 100 %
Weak ionize only slightly in aqueous solution

57. Strength
Strong – means it forms many ions when dissolved (100% ionization)
Mg(OH)2 is a strong base- it falls completely apart (nearly 100% when dissolved).

58. Acid-Base Indicators Compounds whose colors are sensitive to pH.
The color of the indicator changes as the pH of a solution changes.
Indicators are weak acids or bases
HIn H In (In- anion part)

59. Acid-Base Indicators HIn H+ + In-
The color that the indicator displays result from the fact that HIn and In- are different colors
Ex.
HIn H In-
Pink Blue

60. Some of the many pH Indicators and their pH range

61. Acid-Base Indicators
Although useful, there are limitations to indicators:
what if the solution already has a color, like paint?
the ability of the human eye to distinguish colors is limited

62. Acid-Base Indicators
A pH meter may give more definitive results

63. Neutralization Reactions
OBJECTIVES:
Define the products of an acid-base reaction.

64. Neutralization Reactions
OBJECTIVES:
Explain how acid-base titration is used to calculate the concentration of an acid or a base.

65. Acid-Base Reactions Acid + Base Water + Salt
Properties related to every day:
antacids depend on neutralization
farmers adjust the soil pH

66. Acid-Base Reactions
Neutralization Reaction - a reaction in which an acid and a base react in an aqueous solution to produce a salt and water:
HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
H2SO4(aq) + 2KOH(aq) K2SO4(aq) + 2 H2O(l)

67. Titration
Titration is the process of adding a known amount of solution of known concentration to determine the concentration of another solution
Remember - a balanced equation is a mole ratio
The equivalence point is when the moles of hydrogen ions is equal to the moles of hydroxide ions (= neutralized!)

68. Titration – Equivalence Point

69. Titration
The concentration of acid (or base) in solution can be determined by performing a neutralization reaction
An indicator is used to show when neutralization has occurred
Often we use phenolphthalein- because it is colorless in neutral and acid; turns pink in base

70. Steps - Neutralization reaction
#1. A measured volume of acid of unknown concentration is added to a flask
#2. Several drops of indicator added
#3. A base of known concentration is slowly added, until the indicator changes color; measure the volume

71. Neutralization Continue adding until the indicator changes color
added by using a buret
Continue adding until the indicator changes color
called the “end point” of the titration

72. Problem 1
Suppose 20.0 mL of 5.0 x 10-3M NaOH is required to reach the end point in the titration of 10.0 mL of HCl of unknown concentration. Determine the molarity of HCl.
M x V = M x V
Acid base

73. HCl(aq) + NaOH(aq) NaCl(aq) + H2O(l)
known: 20 mL of 5.0 x 10-3 M NaOH
10 mL of HCl
unknown: M of HCl
(0.005 M)(0.020 L) = (M HCl)(0.010 L)
M HCl = 0.01 M

74. Problem 2
In a titration, 27.4 mL of M Ba(OH)2 is added to a 20.0 mL of HCl of unknown concentration. Determine the molarity of HCl.

75. 2HCl + Ba(OH)2 BaCl2 + 2H2O known: 27.4 mL of 0.0154M Ba(OH)2
2 mole mole
known: mL of M Ba(OH)2
20 mL of HCl
unknown: M of HCl
(2)( M)( L) = (M HCl)(0.020 L)
M HCl = M

76. Sample Problems 1 and 2
page 503

77. End of Chapter