1. Acids, Bases, and Salts
By: Chris Squires

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

3. Acids have a pH less than 7

4. OPERATIONAL ACIDS Reacts with metals to form H2 gas.
HCl(aq) + Mg(s) → MgCl2(aq) + H2(g)

5. Taste Sour

6. Conducts electricity.

7. Neutralizes BASES

9. Acids Turn Litmus Red
Blue litmus paper turns red in contact with an acid

10. Sulfuric Acid = H2SO4 Highest volume of any chemical in the U.S.
60 billion pounds/year

11. 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

12. OPERATIONAL BASES

13. BASES ARE SLIPPERY

14. BASES are Corrosive

15. BASE

16. BASES

17. BASES OPERATE

18. Bases have a pH greater than 7

19. Bases Turn Litmus Blue
Red litmus paper turns blue in contact with a base (and blue paper stays blue).

20. Properties of Bases (metallic hydroxides)
React with acids to form water and a salt.
Taste bitter & Feel slippery
Electrolytes in aqueous solution
Change the color of indicators (bases go Blue , with litmus).

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

22. Acid Theories Base H+ H+

23. Acid-Base Theories a) Arrhenius, b) Brønsted-Lowry
c) Lewis (not covered).

24. Svante Arrhenius ( )

25. Theoretical Definitions
Arrhenius acids and bases
Acid: when dissolved in water increases H+.
Base: Substance when dissolved in water, increases hydroxide ions OH-

26. 1. Arrhenius Definition - 1887
Acids produce hydrogen ions (H+) in aqueous solution (HCl → H+ + Cl-)
Bases produce OH- 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 OH- produced.

27. 3 flaws with Arrhenius
1. A proton (H+) does not exist ALONE in water, rather it combines with waters lone pairs to form H3O+(aq)
2. NH3 (g) is a base when dissolved in water; but does not contain the OH-(aq) Amphiprotics like HCO3-(aq) appear to contain the H+ ion BUT turns red litmus blue and ARE bases

28. Johannes Brønsted Thomas Lowry (1879-1947) (1874-1936) Denmark England

29. 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.

30. B/L Definition Brønsted–Lowry: must have both 1. an Acid: Proton donor
and
2. a Base: Proton acceptor

31. A Brønsted–Lowry acid…
…must have a removable (acidic) proton.
HCl, H2O, H2SO4
A Brønsted–Lowry base…
…must accept this H+
NH3, H2O, CO3 2- and usually has negative charge

32. Why Ammonia is a Base
Ammonia can be explained as a base by using Brønsted-Lowry:
NH3(aq) + H2O(l) ↔ NH4+(aq) + OH-(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

33. Acids and bases come in pairs
A “conjugate base” is the remainder of the original acid, after it donates it’s hydrogen ion
A “conjugate acid” is the particle 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.

34. 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

35. Conjugate Acids and Bases:
From the Latin word conjugare, meaning “to join together.”
Reactions between acids and bases always yield their conjugate bases and acids.

36. If can also be BOTH Acid and Base…
...it is amphiprotic.
HCO3–
HSO4 –
H2O

37. Three THEORETICAL definitions
Who
Theory:
Acid
When
Arrhenius
increases H+
1880’s
Brønsted
proton donor
1923
Lowry
ditto
same1923

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

39. 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.

40. How Do We Measure pH? Litmus paper turns blue above ~pH = 7
turns red below ~pH = 7
An indicator
Compound that changes color in solution.

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

42. Other “p” Scales
The “p” in pH tells us to take the negative log of the quantity
Some similar examples are
pOH –log [OH-]
pKw –log Kw

43. The pH concept – from 0 to 14 definition: pH = -log[H+]
pH = “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)
in base, pH > 7 (7 to 14 is base range)

44. Calculating pOH pOH = -log [OH-] [H+] x [OH-] = 1 x 10-14 M2
pH + pOH = 14

46. pH and Significant Figures
For pH calculations, the hydrogen ion concentration is usually expressed in scientific notation
[H+] = M = 1.0 x 10-3 M, and has 2 significant figures
the pH = 3.00, with the two numbers to the right of the decimal corresponding to the 2 sig figs

47. STRONG V WEAK

48. STRONG VS WEAK

49. Strength
In any acid-base reaction, the equilibrium favors the reaction that moves the proton to the stronger base.
HCl(aq) + H2O(l)  H3O+(aq) + Cl–(aq)
Cl– equilibrium lies so far to the right K is not measured (K>>1), so is not a base, it is neutral

50. Strength Strong acids completely dissociate in water.
Their conjugate bases are actually so weak as not to be able to act as a base at all.
Weak acids partially ionize in water.

51. Strength Strength is very different from Concentration
Acids and Bases are classified acording 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
Strength is very different from Concentration

52. 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).
But, not much dissolves- so it is not concentrated

53. Strong Acid Dissociation (makes 100 % ions)

55. Weak Acid Dissociation (only partially ionizes)

56. STRONG V WEAK ACIDS

57. Organic Acids (those with carbon)
Organic acids all contain the carboxyl group, (-COOH), sometimes several of them. CH3COOH – of the 4 hydrogen, only 1 ionizable
(due to being bonded to the highly electronegative Oxygen)
The carboxyl group is a poor proton donor, so ALL organic acids are weak acids.

59. Measuring strength Ionization is reversible for WEAK acids:
HA + H2O ↔ H+ + A-
This makes an equilibrium
Acid dissociation constant = Ka
Ka = [H+ ][A- ] [HA]
Stronger acid = more products (ions), thus a larger Ka
(Note that the arrow goes both directions.)
(Note that water is NOT shown, because its concentration is constant, and built into Ka)

60. What about bases?
Strong bases dissociate completely.
MOH + H2O ↔ M+ + OH- (M = a metal)
Base dissociation constant = Kb
Kb = [M+ ][OH-] [MOH]
Stronger base = more dissociated ions are produced, thus a
larger Kb and a small Ka.

61. Calculating pH from Ka
Calculate the pH of a 0.30 M solution of acetic acid, C2H3O2H, at 25°C.
Ka for acetic acid at 25°C is 1.8  10-5.

62. Calculating pH from Ka Use the ICE table: [C2H3O2], M [H3O+], M
Initial
0.30
Change –x +x
Equilibrium
0.30 – x x

63. Calculating pH from Ka Use the ICE table: [C2H3O2], M [H3O+], M
Initial
0.30
Change –x +x
Equilibrium
0.30 – x ≈ 0.30 x
Simplify: how big is x relative to 0.30?

64. Calculating pH from Ka (1.8  10-5) (0.30) = x2 5.4  10-6 = x2
Now,
(1.8  10-5) (0.30) = x2
5.4  10-6 = x2
2.3  10-3 = x
pH = = 2.64

65. Calculating Ka from the pH
The pH at 25°C is 2.38 for a 0.10 M solution of formic acid, HCOOH,. Calculate Ka
Ka, requires equilibrium concentrations. We can find [H3O+], which is the same as [HCOO−], from pH, so pH = X

66. Calculating Ka from the pH
pH = –log [H3O+]
– 2.38 = log [H3O+]
= 10log [H3O+] = [H3O+]
4.2  10-3 = [H3O+] = [HCOO–] = X

67. Calculating Ka from pH All values In Mol/L [HCOOH] [H3O+] [HCOO−]
Initially
0.10
Change
–4.2  10-3
+4.2  10-3
Equilibrium
0.10 – 4.2  10-3
= = 0.10
4.2  10-3
4.2 

68. Calculating Ka from pH [4.2  10-3] [4.2  10-3] Ka = [0.10]
= 1.8  10-4

69. Calculating Percent Ionization
[A-]eq = [H3O+]eq = 4.2  10-3 M
[A-]eq + [HCOOH]eq = [HCOOH]initial= 0.10 M
%rx = [4.2  10-3 / 0.10] (x 100)
= 4.2%

70. Weak Bases
The Kb for this reaction is

71. pH of Basic Solutions What is the pH of a 0.15M solution of NH3?
[NH4+] [OH−]
[NH3]
Kb =
= 1.8  10-5

72. pH of Basic Solutions [NH3] [NH4+] [OH−] Initial 0.15 Equil
Equil
x  0.15 x

73. pH of Basic Solutions (x)2 (0.15-X) 1.8  10-5 = Ignore X
pOH = –log (1.6  10-3)
pOH = 2.80
pH = 11.20

74. Strength vs. Concentration
The words concentrated and dilute tell how much of an acid or base is dissolved in solution - refers to the # of moles of acid or base
The words strong and weak refer to the extent of ionization
Is a concentrated, weak acid possible?

76. Acid and Base Strength
Acetate is a stronger base than H2O, LOWER on table,
The stronger base “wins” the proton.
HC2H3O2(aq) + H2O
H3O+(aq) + C2H3O2–(aq)

77. TITRATIONS

79. 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!)

80. 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

81. 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

82. Neutralization
The solution of known concentration is called the standard solution
added by using a buret
Continue adding until the indicator changes color
called the “end point” of the titration

83. Polyprotic Acids?
Some compounds have more than one ionizable hydrogen to release
HNO3 nitric acid - monoprotic
H2SO4 sulfuric acid - diprotic - 2 H+
H3PO4 - triprotic - 3 H+
IMPT for TITRATING BUT NOT FOR pH calculating…why?

84. Salt Hydrolysis A salt is an ionic compound that:
comes from the anion of an acid
comes from the cation of a base
is formed from a neutralization reaction
some neutral; others acidic or basic
“Salt hydrolysis” - a salt that reacts with water to produce an acid or base

85. Salt Hydrolysis
To see if the resulting salt is acidic or basic, check the by dissociating the salt
NaCl, a neutral salt
(NH4)2SO4, acidic salt
CH3COOK, basic salt

87. Buffers
Buffers are solutions in which the pH remains relatively constant, even when small amounts of acid or base are added
made from a pair of chemicals: a weak acid and one of it’s salts; or a weak base and one of it’s salts

88. Buffers
A buffer system is better able to resist changes in pH than pure water
Since it is a pair of chemicals:
one chemical neutralizes any acid added, while the other chemical would neutralize any additional base
AND, they produce each other in the process!!!

89. Buffers
Example: Ethanoic (acetic) acid and sodium ethanoate (also called sodium acetate)
The buffer capacity is the amount of acid or base that can be added before a significant change in pH

90. Buffers
The two buffers that are crucial to maintain the pH of human blood are:
1. carbonic acid (H2CO3) & hydrogen carbonate (HCO31-)
2. dihydrogen phosphate (H2PO41-) & monohydrogen phoshate (HPO42-)

91. Aspirin (which is a type of acid) sometimes causes stomach upset; thus by adding a “buffer”, it does not cause the acid irritation.
Bufferin is one brand of a buffered aspirin that is sold in stores.

92. End