1. IB CHEMISTRY
Topic 8 Acids and bases
Higher level

2. 8.1 Theories of acids and bases
OBJECTIVES
• A Brønsted–Lowry acid is a proton/H+ donor and a Brønsted–Lowry base is a proton/H+ acceptor. • Amphiprotic species can act as both Brønsted–Lowry acids and bases. • A pair of species differing by a single proton is called a conjugate acid-base pair. • Deduction of the Brønsted–Lowry acid and base in a chemical reaction. • Deduction of the conjugate acid or conjugate base in a chemical reaction.

3. NOS: Acid-Base Hypothesis
Greeks classifying substances based on taste – sweet, salty, sour....
Alkalis were known to neutralize the sour taste
1777 Antoine Lavoiser burnt phosphorus and sulfur with oxygen. These compounds were acidic when dissolved in water. Therefore acids contain oxygen. Created the word oxygen: Oxys = sour, genes = born in Greek.
1809 Humphry Davy questioned this as HCl, H2S, H2Te were acidic with no oxygen.

4. Key – Best hypothesis is the one with the most predictive power
Arrhenius proposed that salts form ions in solution even without an electric current. Hence acids were possibly substances that released H+ ions, and bases released OH- ions. But this did not explain outside of aqueous solutions or for bases that did not have an –OH group.
Bronsted and Lowry – acid is a H+ donor, base is a H+ acceptor. NH3 can now be explained as a base.
Lewis using his dot structures defined acids as an electron pair acceptor, and a base as an electron pair donor. BF3 + NH3 can now be explained.

5. Learning check - match Phosphorus and sulfur oxides make acids Lewis
HCl does not contain oxygen
Acids give off H+ ions
Acid is a H+ donor, base is a H+ acceptor
Acids are electron pair acceptors and and bases are electron pair donors
Lewis
Lavosier
Davy
Bronsted and Lowry
Arrhenius

6. 1 of 3. Arrhenius Acids and Bases
Arrhenius Acid – produce H+ ions (or hydronium ions H3O+) eg. HCl
Arrhenius Base – produce OH- ions eg. NaOH
(Problem: some bases don’t have hydroxide ions!)

7. 2 of 3. Bronsted-Lowry Acids and Bases
Bonsted-Lowry Acid – proton donor eg. H2O
Bronsted-Lowry Base – proton acceptor eg. NH3
A “proton” is really just a hydrogen atom that has lost it’s electron!

8. A Brønsted-Lowry acid is a proton donor
A Brønsted-Lowry base is a proton acceptor
Base – conjugate acid pair
conjugate acid
conjugate base
base
acid
Acid – conjugate base pair

9. The Brønsted definition means NH3 is a BASE in water — and water is itself an ACID

10. Conjugate Pairs

11. Amphoteric – can act as either an acid or a base
eg. Al2O3 as a base: Al2O3 + 6 HCl  —>   2 AlCl3 + 3 H2O
Al2O3 as an acid: Al2O3 + 2NaOH + 3 H2O   —>  NaAl(OH)4
Amphiprotic – a type of amphoteric substance that specifically can act as either an acid or base by accepting or donating hydrogen ions
eg. water, diprotic acids

12. Learning Check!
Label the acid, base, conjugate acid, and conjugate base in each reaction:
HCl + OH-    Cl- + H2O
H2O + H2SO4    HSO4- + H3O+

13. 8.2 Properties of acids and bases
OBJECTIVES
• Most acids have observable characteristic chemical reactions with reactive metals, metal oxides, metal hydroxides, hydrogen carbonates and carbonates. • Salt and water are produced in exothermic neutralization reactions. • Balancing chemical equations for the reaction of acids. • Identification of the acid and base needed to make different salts. • Candidates should have experience of acid-base titrations with different indicators.

14. Properties of Acids
Often produce H+ in water (also written as the hydronium ion H3O+ which is a hydrogen ion attached to a water molecule)
Taste sour
Corrode metals
Electrolytes
React with bases (metal hydroxides) to form a salt, water
React with carbonates to form salt, water, CO2
React with metals/metal oxides to form salt, H2
pH is less than 7
Turns blue litmus paper red, phenolphthalein colourless, methyl orange red

15. Key reactions
1. Acid + base  salt + water Eg. HCl + NaOH  NaCl + H2O *Base can be a metal hydroxide, or a metal oxide. 2. Acid + metal  salt + hydrogen Eg. HCl + Mg  MgCl2 + H2 (unbalanced) 3. Acid + metal carbonate  salt + water + carbon dioxide Eg. HCl + Na2CO3  NaCl + H2O + CO2 (unbalanced)

16. Acid Nomenclature Flowchart

17. Learning check – nomenclature:
HBr (aq)
H2CO3
H2SO3
 hydrobromic acid
 carbonic acid
 sulfurous acid

18. Properties of Bases Often produce OH- ions in water Taste bitter
Are electrolytes
Feel soapy, slippery
React with acids to form salts and water
pH greater than 7
Turns red litmus paper blue, phenolphthalein pink, methyl orange yellow
An alkali is a base that dissolves in water

19. 8.3 The pH scale
OBJECTIVES
• pH = − log[H+(aq)] and [H+] = 10−pH. • A change of one pH unit represents a 10-fold change in the hydrogen ion concentration [H+]. • pH values distinguish between acidic, neutral and alkaline solutions. • The ionic product constant, Kw = [H+] [OH-] = 10−14 at 298 K. • Solving problems involving pH, [H+] and [OH-]. • Students should be familiar with the use of a pH meter and universal indicator.

20. The pH scale is a way of expressing the strength of acids and bases
The pH scale is a way of expressing the strength of acids and bases. Instead of using very small numbers, we just use the NEGATIVE power of 10 on the Molarity of the H+ (or OH-) ion. pH 1 is 10 times stronger than pH 2 pH 1 is 100 times stronger than pH 3 etc… Under 7 = acid = neutral Over 7 = base

21. Calculating pH pH = - log10 [H+(aq)] [H+] = 10-pH
where [H+] is the concentration of hydrogen ions in mol dm-3
Working backwards:
[H+ (aq)] = antilog (-pH) Or
[H+] = 10-pH

22. Ionic product constant of water - Kw
Water undergoes very slight ionisation ... H2O(l) ⇌ H+(aq) + OH¯(aq)
Applying the equilibrium law: Kc = [H+(aq)] [OH¯(aq)] [H2O(l)]
As the dissociation is small, the water concentration is very large compared with the dissociated ions and any changes to its value are insignificant; its concentration can be regarded as constant.
This “constant” is combined with (Kc) to get a new constant (Kw). At 25°C:
Kw = [H+(aq)] [OH¯(aq)] mol2 dm-6 = 1 x mol2 dm-6
Because the constant is based on an equilibrium, Kw VARIES WITH TEMPERATURE
Also note: pKw = 14 = pH + pOH

23. Kw variation with temperature
• Kw gets larger as the temperature increases
• this means the concentration of H+ and OH¯ ions gets greater
• this means the equilibrium has moved to the right
• if the concentration of H+ increases then the pH decreases
• pH decreases as the temperature increases
Because the equation moves to the right as the temperature goes up, it must be an ENDOTHERMIC process

24. Calculate the pH of 0.02M HCl?
Problem 1:
Calculate the pH of 0.02M HCl?
HCl  H+ + Cl¯
[H+] = 0.02M = 2x10-2 mol dm-3
pH = - log [H+] = 1.7

25. Calculate the pH of 0.1M NaOH?
Problem 2:
Calculate the pH of 0.1M NaOH?
NaOH  Na+ + OH¯
[OH¯] = 0.1M = 1x 10-1 mol dm-3
Method1:
Kw = [H+][OH¯] = 1 x mol2dm-6
[H+] = Kw / [OH¯] = 1 x moldm-3
pH = - log[H+] = 13
Method 2:
pOH = 1
pOH + pH = 14
1 + pH = 14
pH = 13

26. 8.4 Strong and weak acids and bases
OBJECTIVES
• Strong and weak acids and bases differ in the extent of ionization. • Strong acids and bases of equal concentrations have higher conductivities than weak acids and bases. • A strong acid is a good proton donor and has a weak conjugate base. • A strong base is a good proton acceptor and has a weak conjugate acid. • Distinction between strong and weak acids and bases in terms of the rates of their reactions with metals, metal oxides, metal hydroxides, metal hydrogen carbonates and metal carbonates and their electrical conductivities for solutions of equal concentrations.

27. In a nutshell:
A strong acid or base is one that completely dissociates in water. Eg. HCl  H+ + Cl- <1% >99%

28. Strong acids/bases  Note the arrow indicates COMPLETE DISSOCIATION
Strong acids completely dissociate (split up) into ions in aqueous solution
Eg. HCl  H+(aq) Cl¯(aq) MONOPROTIC HNO3  H+(aq) NO3¯(aq)
H2SO4  2H+(aq) SO42-(aq) DIPROTIC
Strong bases completely dissociate into ions in aqueous solution
Eg. NaOH  Na+(aq) + OH¯(aq)
 Note the arrow indicates COMPLETE DISSOCIATION

29. Weak acids/bases ⇌ Note the arrow indicates PARTIAL DISSOCIATION
Weak acids partially dissociate into ions in aqueous solution
Eg. ethanoic acid CH3COOH(aq) ⇌ CH3COO¯(aq) + H+(aq)
Can also be written as:
HA(aq) + H2O(l) ⇌ A¯(aq) + H3O+(aq)
or HA(aq) ⇌ A¯(aq) + H+(aq)
Weak bases partially dissociate into ions in aqueous solution
Eg. NH3(aq) + H2O(l) ⇌ NH4+(aq) + OH¯(aq)
⇌ Note the arrow indicates PARTIAL DISSOCIATION

30. Determining weak/strong, acids/bases experimentally
A weak acid has a higher pH than a strong acid of equal concentration
A weak acid does not conduct electricity as well as a strong acid
A weak acid reacts more slowly than a strong acid eg. acid + metal  H2

31. Examples to be learnt - weak/strong, acids/bases
Strong acids: hydrochloric acid, nitric acid, sulfuric acid
Weak acids: carboxylic acids (eg. ethanoic acid is vinegar), carbonic acid
Strong bases: group 1 hydroxides (eg. NaOH), barium hydroxide
Weak bases: ammonia and amines
18M glacial acetic acid pH of %

32. Learning check: Weak or strong, acid or base?
Weak base
Strong acid
Strong base
Ammonia NH4+ Hydrochloric acid HCl Potassium hydroxide KOH Ethanoic acid CH3COOH

33. 8.5 Acid deposition
OBJECTIVES
• Rain is naturally acidic because of dissolved CO2 and has a pH of 5.6. Acid deposition has a lower pH, usually below 5.0. • Acid deposition is formed when nitrogen or sulfur oxides dissolve in water to form HNO3, HNO2, H2SO4 and H2SO3. • Sources of the oxides of sulfur and nitrogen and the effects of acid deposition should be covered. • Balancing the equations that describe the combustion of sulfur and nitrogen to their oxides and the subsequent formation of HNO3, HNO2, H2SO4 and H2SO3. • Distinction between the pre-combustion and post-combustion methods of reducing sulfur oxides emissions. • Deduction of acid deposition equations for acid deposition with reactive metals and carbonates.

34. Acid deposition
Carbonic acid from CO2 in the atmosphere makes rain a little acidic naturally: CO2(g) + H2O(l) ⇌ H2CO3(aq) H2CO3(aq) ⇌ H+(aq) + HCO3-(aq) HCO3-(aq) ⇌ H+(aq) + CO32-(aq) Precipitation that is below pH 5.6 is considered acid deposition.

35. Acid deposition – nitrogen
At high temperatures such as in engines: N2(g) + O2(g)  2NO(g) (primary reaction) This further reacts with oxygen: 2NO(g) + O2(g)  2NO2(g) This then reacts with water: 2NO2(g) + H2O(l)  HNO3(aq) + HNO2(aq)
1. Keep adding oxygen to the radicals, 2. add water to form acid ions

36. Acid deposition – sulphur
Sulphur dioxide is released by burning coal (up to 3% sulphur): S(s) + O2(g)  SO2(g) (primary reaction) This reacts with water to form sulfurous acid: SO2(g) + H2O(l) ⇌ H2SO3(aq) This can also react further with air, then water to form sulphuric acid: 2SO2(g) + O2(g) ⇌ 2SO3(g) SO3(g) + H2O(l)  H2SO4(aq)

37. Prevention of acid deposition
Pre-combustion methods involve crushing and washing coal with water to remove up to 90% of solid sulphur. Post-combustion methods such as calcium oxide (lime) can remove gaseous sulphur (and other pollutants): CaO(s) + SO2(g)  CaSO3(s)

38. 18.1 Lewis acids and bases
Higher level
OBJECTIVES
• A Lewis acid is a lone pair acceptor and a Lewis base is a lone pair donor. • When a Lewis base reacts with a Lewis acid a coordinate bond is formed. • A nucleophile is a Lewis base and an electrophile is a Lewis acid. • Application of Lewis’ acid–base theory to inorganic and organic chemistry to identify the role of the reacting species.

39. 3 of 3. Lewis Acids and Bases
Higher level
Lewis acid - a substance that accepts an electron pair eg. BF3
Lewis base - a substance that donates an electron pair eg. NH3

40. Formation of hydronium ion is also an excellent example.
Higher level
Formation of hydronium ion is also an excellent example.
Electron pair of the new O-H bond originates on the Lewis base.
Bond formed is a coordinate bond
Water acts as the nucleophile
H+ acts as the electrophile

41. Higher level
Coordinate
bond
Coordinate
bond

42. Lewis Acid-Base interactions in Biology
Higher level
Lewis Acid-Base interactions in Biology
The heme group in hemoglobin can interact with O2 and CO.
The Fe ion in hemoglobin is a Lewis acid
O2 and CO can act as Lewis bases
Heme group

43. 18.2 Calculations involving acids and bases
Higher level
OBJECTIVES
• The expression for the dissociation constant of a weak acid (Ka) and a weak base (Kb). • For a conjugate acid base pair, Ka × Kb = Kw. • The relationship between Ka and pK a is (pKa = -log Ka), and between Kb and pKb is (pKb = -log Kb). • Solution of problems involving [H+(aq)], [OH–(aq)], pH, pOH, Ka, pKa, Kb and pKb. • Discussion of the relative strengths of acids and bases using values of Ka, pKa, Kb and pKb.

44. Acid dissociation constant Ka and pH
Higher level
A weak acid, HA, only partially dissociates:
HA(aq) ⇌ H+(aq)+ A¯(aq)
Applying the Equilibrium Law:
𝐾 𝑎 = 𝐴 − [ 𝐻 + ] [𝐻𝐴]
𝐾 𝑎 = 𝐻 + [ 𝐻 + ] [𝐻𝐴] (as A- and H+ are in equal amounts)
Therefore [H+] = Ka[HA]
From this pH can be determined.

45. Other calculations
Higher level
CH3CHOOH ⇌ CH3COO- + H+ 𝐾 𝑎 = 𝐶𝐻3𝐶𝑂𝑂 − [ 𝐻 + ] 𝐶𝐻3𝐶𝑂𝑂𝐻 CH3COO- + H20 ⇌ CH3CHOOH + OH- 𝐾 𝑏 = 𝐶𝐻3𝐶𝑂𝑂𝐻 [ 𝑂𝐻 − ] 𝐶𝐻3𝐶𝑂𝑂 − Therefore KaKb = [H+][OH-] = Kw By mathematical rule: pKa + pKb = pKw
Addition

46. KEY formula
Higher level
pH = -log[H+] [H+] = 10-pH K = [products]/[reactants] Kw = Ka x Kb = pKw = pH + pOH = 14

47. Problem 1:
Calculate the pH of a weak acid HX of concentration 0.1M (Ka = 4x10-5 mol dm-3).
HA(aq) ⇌ H+(aq)+ A¯(aq)
𝐾 𝑎 = 𝐴 − [ 𝐻 + ] [𝐻𝐴] (H+ and A- are in equal amounts)
[H+] = [HA]Ka
Key assumption: Equilibrium and original concentration of weak acid is the same due to little dissociation:
[H+(aq)] = x 4x10-5 mol dm-3
= x10-6 mol dm-3
= x10-3 mol dm-3
pH = - log [H+(aq)] = 2.699

48. Size of Ka Kb pKa pKb values
Higher level
Size of Ka Kb pKa pKb values
Similarly to pH, a large [H+] ion concentration gives a small pH value.
A large Ka value means the acid is strong, and hence small pKa value also means the acid is strong.
A strong acid ‘completely dissociates’
HCl Ka = 1.3 x 106, H2SO4 Ka = 1.0 x 103, HNO3 Ka = 24
A weak acid ‘only partially dissociates’
Ethanoic acid 1.8 x 10-5, Carbonic acid 4.4 x 10-7
Addition

49. 18.3 pH curves
Higher level
OBJECTIVES
• The characteristics of the pH curves produced by the different combinations of strong and weak acids and bases. • An acid–base indicator is a weak acid or a weak base where the components of the conjugate acid–base pair have different colours. • The relationship between the pH range of an acid–base indicator, which is a weak acid, and its pKa value. • The buffer region on the pH curve represents the region where small additions of acid or base result in little or no change in pH. • The composition and action of a buffer solution. • The general shapes of graphs of pH against volume for titrations involving strong and weak acids and bases with an explanation of their important features. • Selection of an appropriate indicator for a titration, given the equivalence point of the titration and the end point of the indicator. • While the nature of the acid–base buffer always remains the same, buffer solutions can be prepared by either mixing a weak acid/base with a solution of a salt containing its conjugate, or by partial neutralization of a weak acid/base with a strong acid/base. • Prediction of the relative pH of aqueous salt solutions formed by the different combinations of strong and weak acid and base.

50. Salt hydrolysis Higher level
Salts of weak acids and bases react with water in such a way that the pH of these salt solution is not neutral.
Eg.
FeCl2 salt: Fe2+ + H2O ⇌ Fe(OH)2 + 2H+
Hence a salt of a weak base will make the solution slightly acidic.
CH3COONa salt: CH3COO- + H2O ⇌ CH3COOH + OH-
Hence a salt of a weak acid will make the solution slightly alkaline.
Na+ will not convert to NaOH, Cl- will not covert to HCl, as they are strong acids and bases. Think – weak acids and bases, really are strong! Strong at holding onto their hydrogen or hydroxide ions!

51. Buffers Higher level small quantities of acid or alkali are added.”
Definition “Solutions which resist changes in pH when
small quantities of acid or alkali are added.”
Acidic Buffer (pH < 7) made from a weak acid its sodium or potassium salt ethanoic acid sodium ethanoate
Alkaline Buffer (pH > 7) made from a weak base its chloride ammonia ammonium chloride
Uses Standardising pH meters
Buffering biological systems (eg in blood)
Maintaining the pH of shampoos

52. Buffer solutions - ideal concentration
Higher level
The concentration of a buffer solution is also important
If the concentration is too low, there won’t be enough CH3COOH and CH3COO¯
to cope with the ions added.
Summary
For an acidic buffer solution one needs ...
large [CH3COOH(aq)] - for dissociating into H+(aq) when alkali is added
large [CH3COO¯(aq)] - for removing H+(aq) as it is added
This situation can’t exist if only acid is present; a mixture of the acid and salt is used.
The weak acid provides the equilibrium and the large CH3COOH(aq) concentration.
The sodium salt provides the large CH3COO¯(aq) concentration.
One uses a WEAK ACID + its SODIUM OR POTASSIUM SALT

53. How buffers work CH3CHOOH ⇌ CH3COO- + H+ In equilibrium
Higher level
CH3CHOOH ⇌ CH3COO- + H+ In equilibrium
CH3CHOONa  CH3COO- + Na+ Complete dissociation
If an acid is added: HCl  H+ + Cl-
The excess H+ will cause equilibrium to be pushed to the left and so the pH is remains the same:
CH3COO- + H+  CH3CHOOH
If a base is added: NaOH  Na+ + OH-
The H+ will neutralize the OH-: H+ + OH-  H2O
The equilibrium will be pushed to the right due to loss of H+ and so the pH remains the same:
CH3CHOOH  CH3COO- + H+
(note: state symbols must also be written with these equations in the exam)

54. Indicator chemical reactions
Higher level
Indicator chemical reactions
Indicators are weak acids which have a different colour to their conjugate base
HIn H+ + In-
colour colour 2
low pH: equilibrium pushed left = colour 1
high pH: equilibrium pushed right = colour 2

55. Higher level
Methyl orange
pH < 3.2
pH > 4.4

56. Higher level
acid
end-point
alkali

57. Higher level
Phenolphthalein
pH < 8.2
pH > 10.0

58. Higher level
phenolphthalein
acid
alkali

59. Relationship of pH to pKa value
Higher level
Relationship of pH to pKa value

60. Relationship of pH to pKa value
Higher level
Relationship of pH to pKa value
When pH = pKa the colour will be midway. If you increase [H+] it pushes the indicator to the acid colour. It takes about a 10 times difference in indicator colour species to see the colour change therefore most colour changes take about 2 pH units.

61. Indicators
Higher level
For an indicator to work in a titration, the pH range of the indicator’s colour change must be within the range of the pH change at the end-point.
Indicator
colour at
low pH
pH range of
colour change
high pH
methyl orange
red
3.2 – 4.4
orange
phenolphthalein
colourless
8.2 – 10.0
purple

62. pH curves weak acid (CH3COOH) vs strong alkali (NaOH)
Higher level
Types There are four types of acid-base titration; each has a characteristic curve.
strong acid (HCl) vs strong base (NaOH)
weak acid (CH3COOH) vs strong alkali (NaOH)
strong acid (HCl) vs weak base (NH3)
weak acid (CH3COOH) vs weak base (NH3)
In the following examples, alkali (0.1M) is added to 25cm3 of acid (0.1M)
End points need not be “neutral‘ due to the phenomenon of salt hydrolysis

63. Terminology Higher level
Equivalence point (stoichiometric point) – exact point at which the solution is neutralized for the given acid (solution may not actually be neutral), given as a volume
Endpoint – point at which an indicator changes colour, given as a pH
Point of inflection – place in the pH curve where a large change in pH occurs
Half-equivalence – point in the graph where half of the acid has been neutralized and turned into a salt
Buffer region – point on the graph that is most resistant to changes in pH

64. Strong acid (HCl) vs strong base (NaOH)
Higher level
Strong acid (HCl) vs strong base (NaOH)
Curve levels off at pH 13 due to excess 0.1M NaOH
(a strong alkali)
Very sharp change in pH over the addition of less than half a drop of NaOH
Very little pH change during the initial 20cm3
pH 1 at the start due to 0.1M HCl (strong monoprotic acid)

65. Higher level
Any of the indicators listed will be suitable - they all change in the ‘vertical’ portion
PHENOLPHTHALEIN
LITMUS
METHYL ORANGE

66. Strong acid (HCl) vs weak base (NH3)
Higher level
Curve levels off at pH 10 due to excess 0.1M NH3
(a weak alkali)
Sharp change in pH over the addition of less than half a drop of NH3
Very little pH change during the initial 20cm3
pH 1 at the start due to 0.1M HCl

67. Higher level
Only methyl orange is suitable - it is the only one to change in the ‘vertical’ portion
PHENOLPHTHALEIN
LITMUS
METHYL ORANGE

68. Weak acid (CH3COOH) vs strong base (NaOH)
Higher level
Curve levels off at pH 13 due to excess 0.1M NaOH
(a strong alkali)
Sharp change in pH over the addition of less than half a drop of NaOH
Steady pH change
pH 4 due to 0.1M CH3COOH (weak monoprotic acid)

69. Higher level PHENOLPHTHALEIN LITMUS METHYL ORANGE
Only phenolphthalein is suitable - it is the only one to change in the ‘vertical’ portion

70. Weak acid (CH3COOH) vs weak base (NH3)
Higher level
Curve levels off at pH 10 due to excess 0.1M NH3
(a weak alkali)
NO SHARP
CHANGE IN pH
Steady pH change
pH 4 due to 0.1M CH3COOH (weak monoprotic acid)

71. Higher level
There is no suitable indicator- none change in the ‘vertical’ portion.
The end point can be detected by plotting a curve using a pH meter.
PHENOLPHTHALEIN
LITMUS
METHYL ORANGE
NOTHING SUITABLE

72. SUMMARY for Phenolphthalein
Higher level
SUMMARY for Phenolphthalein

73. SUMMARY for methyl orange
Higher level
SUMMARY for methyl orange