1. Acids and Bases pH and Titrations

2. Overview Acid – Base Concepts Acid and Base Strengths
Arrhenius
Brønsted – Lowry
Lewis
Acid and Base Strengths
Relative Strengths of Acids and Bases
Molecular Structure and Acid Strength
Self – Ionization of Water and pH
Self – Ionization of Water
Solutions of a Strong Acid or Base
The pH of a Solution

3. II. Theories of acid and bases
A. Arrhenius theory
1. Arrhenius Acid as known as a traditional acid
a. a chemical compound that contains
hydrogens and ionizes in aqueous solution
to form hydrogen ions
2. Arrhenius base
a. any chemical that produces hydroxide ions

4. B. Bronsted -Lowry theory 1923
(worked independently)
1. Bronsted Acid
a. an ion or molecule that is a proton donor
2. Bronsted base
a. anything which will accept a proton

5. a. Strong acids have weak conjugate bases
3. Conjugate acid - base pairs
HF HOH ===> H3O F-
Acid1 Base2 Acid Base1
a. Strong acids have weak conjugate bases
b. Strong bases and weak conjugate acids
4. Autoionization - self-ionization
HOH <===> H+ + OH -

6. Conjugate Acid-Base Pairs
Fig. 17-1, p. 507

7. Conjugate Acid-Base Pairs

8. General Equation Brønsted-Lowry Theory of Acids & Bases
Conjugate Acid-Base Pairs
General Equation
p. 507

9. Brønsted-Lowry Theory of Acids & Bases
p. 504

10. 5. Amphiprotic ability to act as an acid or base
C. Lewis theory
1. Lewis acid
a. anything which can accept a pair of
electrons
2. Lewis Base
a. anything which will donate a pair of
electrons

11. Lewis Theory of Acids & Bases
p. 506

12. Objectives Properties of acids and bases The pH scale
Distinguish between strong and weak acids and list the clinical uses of these acids
Distinguish between strong and weak bases and list the clinical uses of these acids
Understand neutralisation and the clinical applications of neutralisation

13. On page 355 Do questions 4,6,7,10,14,
17,18,22,23,25,26,28

14. Some acids react with active metals to release
Properties of Acids
Aqueous solutions of acids have a sour taste.
Acids change the color of acid-base indicators.
Some acids react with active metals to release
hydrogen gas.
Acids react with bases to produce salts and water.
Some acids conduct electric current.

15. PROPERTIES OF ACIDS & BASES
Produce hydrogen ions (H+) in H2O
Taste sour
Act as electrolytes in solution
Neutralise solutions containing hydroxide
ions (OH -)
React with several metals
releasing H2(g)  corrosion
React with carbonates releasing CO2(g)
Destroy body tissue

16. How many foods can you think of that are sour?
Chances are, almost all of the foods that you
associate with being sour, owe their sour taste
to an acid:
Lemons – citric acid
Grapefruit – citric acid
Apples – malic acid
Sour milk – lactic acid
Vinegar – acetic acid
Grapes – tartaric acid

17. Strength of acids A strong acid is one that ionizes completely
in aqueous solutions.
The strength of an acid depends on the polarity of
the bond between hydrogen and the element to
which it is bonded and the ease with which that
bond can be broken.
Acid strength increases with increasing polarity
and decreasing bond energy.
Acids that are weak electrolytes are weak acids.

19. Strong acids are:
Strong electrolytes
~ 100% ionisation  good conductors
Severe burns to body tissue
*** Stomach lining protected against HCl by mucus

20. STRENGTHS OF ACIDS Strong Acids (very few) Eg HCl Hydrochloric Acid
Marieb, Fig 26.11
Strong Acids (very few) Eg
HCl Hydrochloric Acid
~ Stomach acid
HNO3 Nitric Acid
~ May be used to cauterise warts
~Drugs, explosives, fertilisers, dyes
H2SO4 Sulphuric Acid
~  conc. to treat stomach hypoacidity
~ Fertilisers, dyes, glues

21. Factors Affecting Acid Strength
Binary acids:
Bond strength is directly related to the acid
strength (bond size).
HI and HBr have larger bonds lengths and are
more acidic than HF and HCl, even though
fluorine is most electronegative.
For bonds of similar size the acid strength is
related to electronegativity difference.

22. Oxyacids are acidic substances that contain
oxygen and some other nonmetal, e.g. HNO3
An increase in the electronegativity of an
atom bound to oxygen increases in polarity
of the bond and makes it more acidic.
More oxygen = more polar.

23. Weak Acids (most acids in nature)
CH3COOH Acetic Acid
~ Vaginal jellies, antimicrobial solution  ears, plastics, dyes, insecticides
H2CO3 Carbonic Acid
~Bicarbonate buffer system, carbonated drinks
H3PO4 Phosphoric Acid
~ Drugs, fertilisers, soaps, detergents, animal feed

24. Bases
Produce or cause an increase in
hydroxide ions (OH-) in H2O
Taste bitter
Turn red litmus  blue
Act as electrolytes in solution
Neutralise solutions containing
hydrogen ions (H +)
Have a slippery, ‘soapy’ feel
Destroy body tissue/ dissolve fatty (lipid) material

25. How many bases can you think of?
Ammonia
Sodium hydroxide – lye – drain cleaner
Milk of magnesia – Mg(OH)2 – antacid
Aluminum hydroxide – antacid
Baking soda – sodium hydrogen carbonate

26. 4. STRENGTHS OF BASES Strong Bases NaOH Sodium Hydroxide
~ Removes grease – drains, ovens
Mg(OH)2 Magnesium hydroxide
~ Antacid ~ Laxative
Al(OH)3 Aluminium hydroxide
~ Antacid
~ Absorbs toxins, gases,
~ Causes constipation

27. Strength of Bases
The strength of a base depends on the extent to which the base dissociates.
Strong bases are strong electrolytes.
Strong bases: calcium hydroxide, barium hydroxide, sodium hydroxide, etc
Weak bases: ammonia, aniline
Table 15-4 page 461

28. Relative Strengths of Acids and Bases and Extent of Reaction
The table of relative strengths of acids and conjugate bases can be used to predict if a reaction will produce product.
E.g. Which will produce product?
HNO3 + CN or HCN +

29. Strong bases are:
Strong electrolytes
~ 100% dissociation in water  good conductors
Severe damage to skin & eyes
(Group 1A elements)

30. Weak Bases Eg
NH3 Ammonia
~ Waste product of protein break down in body.
CO3 2- In antacids
HCO3 – In antacids, buffers
HPO4 2- In buffers

31. Reaction with Water : Weak bases NH3(g) + H2O NH4 + (aq) + OH – (aq)
Weak bases are:
Weak electrolytes
Do not contain OH – but react with H2O  small numbers of OH –
Reaction with Water : Weak bases
NH3(g) + H2O NH4 + (aq) + OH – (aq)
HCO3 – (aq) + H2O H2CO3 (aq) OH-(aq)

32. Acids & Bases STRONG vs WEAK H2SO4 NaOH HI KOH HBr Ca(OH)2 HCl Sr(OH)2
_ completely ionized _ partially ionized
_ strong electrolyte _ weak electrolyte
_ ionic/very polar bonds _ some covalent bonds
Strong Acids: Strong Bases:
HClO4 LiOH
H2SO4 NaOH
HI KOH
HBr Ca(OH)2
HCl Sr(OH)2
HNO3 Ba(OH)2

33. Dissociation in Water : Strong bases Metal hydroxides  ions
H2O
NaOH(s) Na+(aq) OH-(aq)
H2O
Mg(OH)2(s) Mg 2 + (aq) OH- (aq)
H2O
Al(OH)3(s) Al 3+(aq) OH- (aq)

34. 5. ACID-BASE NEUTRALISATION
Neutralisation Reaction
Acid + Base  Salt + Water
HCl + NaOH  NaCl + H2O
H+ + OH –  H2O
Neutralise each other
Must be equal concentrations

35. Antacids – clinical applications (Check for side effects!!)
~ Neutralise excess stomach acid
~ Raise stomach pH > 4
Pepsin inactive
~ Assist with ulcer treatment
~  solubility in H2O but still produce high % of ions

36. CaCO 3
2HCl + CaCO 3  CaCl H2O + CO 2 (g)
~Also a Ca 2 + supplement
Long term overuse  Ca 2 + levels
 risk kidney stones (renal calculi)

37. Eg
Mg(OH) 2 Milk of Magnesia & in Mylanta
2HCl + Mg(OH) 2  MgCl H2O
Al(OH) 3 In Mylanta
3HCl + Al(OH) 3  AlCl H2O

38. NaHCO 3 Baking Soda
Not recommended!!
HCl + NaHCO  NaCl + H2O + CO2 (g)
~ Elderly tend to OD
 Stomach can ‘explode’

39. Dissociation in Water : Weak acids Polar covalent molecules
Weak acids are:
Weak electrolytes
Small % ionisation  weak conductors
Dissociation in Water : Weak acids
Polar covalent molecules
 Mainly stay as molecules

40. Dissociation in Water : Strong acids Polar covalent molecules  ions
Eg.
HCl(l) H+(aq) + Cl-(aq)
HNO3(l) H+ (aq) + NO3- (aq)
H2SO4(l) H+ (aq) + SO42- (aq)
H2O
H2O
H2O

41. Dissociation in water : Weak acids (cont)
CH3COOH (l) H+ (aq) + CH3COO- (aq)
H2O
H2CO3 (l) H+ (aq) + HCO3-(aq)
H2O
H2O
H3PO4 (l) H+ (aq) + H2PO4- (aq)

42. 2. THE pH SCALE Ion Product of Water Pure H2O at 25°C
Some molecules ionise
H2O  H+ + OH-
[H+ ] = 1 x 10-7 M = [OH- ]

43. On the pH scale, values below 7 are acidic, a value of 7 is neutral, and values above 7 are basic.

44. The pH scale pH = - log [H+] 1 2 3 4 5 6 7 8 9 10 11 12 13 14
The pH scale ranges from 1x 100 to 1 x mol/L or from 1 to 14.
pH = - log [H+]
acid neutral base

45. Stomach juice: pH = 1.0 – 3.0
Human blood: pH = 7.3 – 7.5
Lemon juice: pH = 2.2 – 2.4
Seawater: pH = 7.8 – 8.3
Vinegar: pH = 2.4 – 3.4
Ammonia: pH = 10.5 – 11.5
Carbonated drinks: pH = 2.0 – 4.0
0.1M Na2CO3: pH = 11.7
Orange juice: pH = 3.0 – 4.0
1.0M NaOH: pH = 14.0

46. pH describes [H+ ] & [OH- ]
Indicates if a fluid is :
0 Acidic [H+ ] = [OH- ] =10-14
7 Neutral [H+ ] = [OH- ] =10-7
14 Basic [H+ ] = [OH- ] = 100

47. Using the pH scale
Exponential values for [H+ ] & [OH- ] inconvenient in a clinical workplace
Simplify  pH scale
 acid-base concentration
p  potential or Power
H  Hydrogen

48. Acidic solution
[H+ ] > [OH- ]
Neutral solution
[H+ ] = [OH- ]
Basic solution
[H+ ] < [OH- ]

49. Equilibrium constant for water
Water Equilibrium
Kw = [H+] [OH-] = 1.0 x 10-14
Equilibrium constant for water
Water or water solutions in which [H+] = [OH-] = 10-7 M are neutral solutions.
A solution in which [H+] > [OH-] is acidic
A solution in which [H+] < [OH-] is basic

50. Autoionization of Water
Water can act as both an acid and a base equilibrium is:
H2O + H2O  H3O+ + OH.
Kw = [OH][H3O+] = 1.00 x 1014M2.
Since [OH] = [H3O+] [H3O+] = 1 x 107 M
(called a neutral solution)

51.  Ion Product of H2O: * Add exponents = 1 x 10-14
[H+ ] x [OH- ] = [1 x 10-7 ] x [1 x 10-7 ]
* Add exponents
= 1 x 10-14

52. Acidic [H3O+] > 1.00x107M
Neutral [H3O+] = 1.00x107M
Basic [H3O+] < 1.00x107M
All acids/bases dissolved in water must obey
equation for the ionization of water.
They either add H3O+ or OH to water.

53. Most of the acids in this chapter will be stronger
than water and add significantly to the
hydronium ion concentration.
E.g. the hydronium ion concentration of an acidic
solution was 1.00x105 M. What was the [OH]?
E.g. what is the hydronium ion concentration
if the hydroxide concentration was 2.50x103 M?

54. The pH loop
Fig. 17-3, p. 516

55. The pH of a Solution E.g. determine the pH of a solution in which
pH = log[H3O+] and [H3O+] = 10pH
Acidic pH < 7.00
Neutral pH = 7.00
Basic pH > 7.00
E.g. determine the pH of a solution in which
[H3O+] = 5.40x106 M
E.g.2 determine the pH of a solution in which the
[OH] = 3.33x103 M

56. E.g.3 determine the pOH of a solution in which
the [OH] = 3.33x103 M
E.g.4 Determine the [H3O+] if the pH of the
solution is The term pX is defined in exactly
the same way as pH
Eg.5 What is the pCa if [Ca2+] = 6.44x10-4

57. E.g. 6 what is pH and [OH] of 0.125 M Ba(OH)2.
[H3O+] of water is small compared to added [H3O+]
from the acid and ignored in the calculation.

58. pH scale
The scale for measuring the hydronium ion concentration [H+] in any solution must be able to cover a large range. A logarithmic scale covers factors of 10. A solution with a pH of 1 is 10 times stronger than a solution with a pH of 2
A solution with a pH of 1 has [H+] of 0.1 mol/L
or 10-1
A solution with a pH of 3 has [H+] of
mol/L or 10-3
A solution with a pH of 7 has [H+] of
mol/L or 10-7

59. pH + pOH = 14 ; the entire pH range!
Manipulating pH
Algebraic manipulation of:
pH = - log [H+]
allows for:
[H+] = 10-pH
If pH is a measure of the hydronium ion concentration then the same equations could be used to describe the hydroxide (base) concentration.
[OH-] = 10-pOH pOH = - log [OH-]
thus:
pH + pOH = 14 ; the entire pH range!

60. Table 17-3, p. 518

61. Methods of Measuring pH
pH paper is used that has compounds in it
which are change to different colors for different
pH ranges.
An colored indicator can be placed in the solution
and its color correlated with pH.
HIn(aq) + H2O(l)  H3O+(aq) + In(aq).
E.g. phenolphthalein is colorless in acid form but
pink in basic form.

62. The pH at which they change color depends on
their equilibrium constant.
More accurate and precise
measurements are made with a pH meter.
A combination of voltmeter and electrodes.

63. Indicators and pH meters
Acid-Base indicators are compounds whose color is sensitive to pH.
The color of an indicator changes as the pH of a solution changes.
Indicators come in many colors.
The pH range over which an indicator changes color also varies. This pH range is also called the transition interval.

64. Indicators and pH meters con’t
A universal indicator is made by mixing several different indicators.
If an exact value for the pH of a solution is needed, a pH meter should be used.
A pH meter determines the pH of a solution be measuring the voltage between the two electrodes that are placed in the solution.

65. Sample Problems Q1: Calculate the pH of a solution if
[H+] = 2.7 x 10-4 M
pH = -log[H+] pH = -log(2.7 x 10-4) = 3.57
Q2: Find the hydrogen ion concentration of a
solution if its pH is
[H+] = 10-pH [H+] = = 2.4 x 10-12M

66. Q3: Find the pOH and the pH of a solution if its
hydroxide ion concentration is 7.9 x 10-5M
pOH = -log[OH-]
pOH = -log(7.9 x 10-5) = 4.10
pH + pOH = 14
pH =
pH = 9.9

67. A brief introduction to equilibrium constants and Ionization constants
Equilibrium is defined as a state where concentrations of
products and reactants are unchanging
The equilibrium constant is a mathematical expression of the
ratio of the products compared to the reactants
K = { products}
The { } are read as concentration
{ reactants}
All species are considered to be at equilibrium

68. Equilibrium constant and Ionization constants
can be determine for many reactions
For acids
Ka = { H3O +} { A-}
{HA} {H2O}
Since the concentration of H2O is roughly 55 moles /liter
it can be deleted from the above equation therefore the
equation becomes
Ka = { H3O +} { A-}
{HA}

69. Titration A neutralization reaction occurs between an acid and a base.
Because acids and bases react, the progressive addition of an acid to a base or vice versa can be used to compare the concentrations of the acid and the base.

70. Titrations con’t
Titration is the controlled addition and measurement of the amount of a solution of known concentration required to react completely with a measure amount of a solution of unknown concentration
Titration provides a sensitive means of determining the chemically equivalent volumes of acidic and basic solutions.

71. TITRATION nacid = nbase
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is NOT 1:1 then you must rely on the mole relationship and handle the problem like any other stoichiometry problem.
MOLES OF ACID = MOLES OF BASE
nacid = nbase

72. Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
TITRATION
MAVA = MBVB
1. Suppose mL of hydrochloric acid was required to neutralize mL of 0.52 M NaOH. What is the molarity of the acid?
HCl + NaOH  H2O + NaCl
Ma Va = Mb Vb rearranges to Ma = Mb Vb / Va
so Ma = (0.52 M) (22.50 mL) / (75.00 mL)
= 0.16 M
Now you try:
2. If mL of M HNO3 neutralized mL of KOH, what is the molarity of the base?
Mb = mol/L

73. TITRATION
1. If mL of M LiOH neutralized mL of H2SO4, what is the molarity of the acid?
2 LiOH + H2SO4  Li2SO4 + 2 H2O
First calculate the moles of base:
L LiOH (0.543 mol/1 L) = mol LiOH
Next calculate the moles of acid:
mol LiOH (1 mol H2SO4 / 2 mol LiOH)= mol H2SO4
Last calculate the Molarity:
Ma = n/V = mol H2SO4 / L = M
2. If mL of Ba(OH)2 solution was used to titrate mL of M HCl, what is the molarity of the barium hydroxide solution?
Mb = mol/L

74. TITRATION nacid = nbase MAVA = MBVB
Titration of a strong acid with a strong base
ENDPOINT = POINT OF NEUTRALIZATION = EQUIVALENCE POINT
At the end point for the titration of a strong acid with a strong base, the moles of acid (H+) equals the moles of base (OH-) to produce the neutral species water (H2O). If the mole ratio in the balanced chemical equation is 1:1 then the following equation can be used.
MOLES OF ACID = MOLES OF BASE
nacid = nbase
Since M=n/V
MAVA = MBVB

75. Equivalence Point (E.P.)
The point at which the two solutions used in a titrations are present in chemically equivalent amounts is the equivalence point (e.p.).
Indicators and pH meters cane be used to determine the e.p.
A pH meter will show a large voltage change occurring at the e.p.
An indicator changes color over a range that includes the pH of the e.p.

76. End point
The point in a titration at which an indicator changes color is called the end point.

77. Which indicator should you use?
Indicators that undergo transition at about pH 7 are used to determine the equivalence point of strong acid/ strong base titrations because the neutralization of strong acids with strong bases produces a salt solution with a pH of approximately 7.

78. Which indicator should you use?
Indicators that change color at pH lower than 7 are useful in determining the equivalence point of strong acid/ weak base titrations. The equivalence point of this titration is acidic because the salt formed is a weak acid. Thus the salt solution has a pH lower than 7.

79. Which indicator should you use?
Indicators that change color at pH higher than 7 are useful in determining the equivalence point of weak acid/ strong base titrations. These reactions produce salt solutions whose pH is greater than 7. This occurs because the salt solution formed is a weak base.

80. Which indicator should you use?
To determine the equivalence point for a weak acid/ weak base, we use no indicator. The pH of the equivalence point of weak acids and weak bases may be almost any value, depending on the relative strengths of the reactants. The color transition of an indicator helps very little in determining whether reactions between weak acids and bases are complete.

81. Molarity and titration
To calculate the molarity of a substance using titration:
Start with the balanced equation and determine the chemically equivalent amounts of the acid and base.
Determine the moles of acid or base from the known solution used in the titration.

82. Molarity and titration
Determine the moles of solute of the unknown solution used during the titration.
Determine the molarity of the unknown solution.

83. Predicting Acid-Base Reactions
Fig. 17-2, p. 511

84. Review questions (cont)
Distinguish between strong & weak bases; List clinical uses of these bases & write equations for their dissociation in water.
Complete simple equations for the neutralisation reaction of an acid & a base; Discuss clinical applications of acid-base neutralisation.

85. Review questions List the properties of acids & bases.
Discuss the pH scale: *Define the ion product of water & indicate how this is determines the pH scale * Use the pH scale to determine if a given solution is acidic, neutral or basic.
Distinguish between strong acids & weak acids: List clinical uses of these acids & write equations for their dissociation in water

86. Brønsted-Lowry Theory of Acids & Bases
Notice that water is both an acid & a base = amphoteric
Reversible reaction
p. 504

87. Water Equilibrium
p. 514