1. Unit F: Applications of Acid-Base Reactions

2. A. Hydrolysis of Salts
A salt is a compound with an ionic bond, but not an acid or base.
Salts can be formed by
Acid/base neutralization
KOH + HNO3 →KNO3 + H2O
Synthesis
2 Ag + Cl2 → 2 AgCl
Acid/metal reaction
2 HCl + Zn → ZnCl2 + H2

3. Hydrolysis
hydrolysis is the reaction of a salt with water to give an acidic or basic solution.
The regular [H+] and [OH-]are changed

4. Spectator ions
spectator ions
Are ions which do not hydrolyze(react with water)
Are the conjugates of strong acids and strong bases.
Examples
HCl → H+ + Cl Cl- never hydrolyzes.
NaOH→ Na+ + OH- Na+ never hydrolyzes.

5. Spectator ions are the ions of Groups I and II and
the first 5 anions at the top of the Ka table.
Anions: ClO4-, I-,Br-,Cl-,NO3-
Cations: LI+, Na+, K+, Rb+, Cs+, Fr+,
Be2+, Mg2+, Ca2+ , Sr2+ , Ba2+ , Ra2+

6. Salt Solution: Acidic or Basic?
1. Figure out which ions are produced by the salt in solution.
2. Ignore the spectator ions.
3. If the remaining ion is on the left side of the
Ka table, it will act as an acid.
If the remaining ion is on the right side of the
Ka table, it will act as a base.

7. CH3COO- + H2O ⇌ CH3COOH +OH-
Examples of Hydrolysis of salts
1. NaCl
NaCl→ Na+ + Cl-
Both ions are spectator. Solution is neutral.
2. NaCH3COO
NaCH3COO → Na+ + CH3COO-
Na+ is a spectator
CH3COO- + H2O ⇌ CH3COOH +OH-
solution is basic.

8. Examples of Hydrolysis of salts
3. NaCN
NaCN→ Na+ + CN-
Na+ is a spectator
CN- + H2O ⇌ HCN + OH-
Solution is basic.
4. NH4Cl
NH4Cl → NH4+ + CI­-
CI- is a spectator.
NH4+ + H2O ⇌ H3O+ + NH3
Solution is acidic.

9. CH3COO- + H2O ⇌ CH3COOH +OH-
Examples of Hydrolysis of salts
5. NH4CH3COO
NH4CH3COO → NH4+ + CH3COO-
Both ions will hydrolyze. Who wins?
NH4+ + H2O ⇌ H3O+ + NH3
Ka=____________________
CH3COO- + H2O ⇌ CH3COOH +OH-
Kb=___________________

10. HCO3- + H2O ⇌ H2CO3 +OH- Examples of Hydrolysis of salts 6. NaHCO3
NaHCO3 → Na+ + HCO3-
Na+ is a spectator. But HCO3- is amphiprotic and can hydrolyze two different ways! Who wins?
HCO3- + H2O ⇌ H3O+ + CO32-
Ka=____________________
HCO3- + H2O ⇌ H2CO3 +OH-
Kb=___________________

11. Assignment
Hebden # 69abcdgh, 70abcdgh, 71, 72 and 73.
Heath I and A p. 614 Q. 16, 21

12. Anhydrides
1. Acid Anhydride: an oxygen-containing compound that reacts with water to produce an acidic solution.
Examples:
SO ____________________
CO _____________________
N2O5 ____________________
Find the acid anhydride of H2SO3:

13. Anhydrides
2. Basic Anhydride: an oxygen-containing compound that reacts with water to produce a basic solution.
Examples:
K2O _____________________
Find the basic anhydride of Ca(OH)2

14. Anhydrides
Acid Anhydride + Basic Anhydride→ SALT
______________________________ a salt
If mix the anhydrides with water first:
_____________________ base
______________________ acid
Then mix those two solutions:
______________________________ salt and water!

15. Anydrides
Basic Amphiprotic Acidic
Periodic Table of Elements

16. Assignment
Hebden #144 abcdef, 145 abcdef

17. Buffers
Buffer: an equilibrium system that maintains a relatively constant pH when small amounts of base or acid are added.
Buffers are commonly present in blood, soap, baby shampoo.

18. A salt of that weak acid or base
Buffers
Buffer Recipe
Mix
A weak acid or base
With
A salt of that weak acid or base

19. Buffers
Example 1: add KCH3COO to a solution of CH3COOH.
________________________________________
How It works:
If you add a base • If you add an acid:

20. Buffers
Example 2: How could you make a buffer using H2CO3?

21. CH3COOH ⇌ CH3COO- + H+ Buffers
Can a weak acid alone (like acetic acid) buffer against the addition of a base ? ( _________)
Equilibrium:
CH3COOH ⇌ CH3COO- + H+
Add OH-
Equilibrium shifts right to replace H+ removed by OH-
And it can keep shifting right because [CH3COOH] is high because it is a weak acid and most remains undissociated.
So [H+] remains relatively constant.

22. CH3COOH ⇌ CH3COO- + H+ Buffers
Can a weak acid alone (like acetic acid) buffer against the addition of a acid ? ( _________)
Equilibrium:
CH3COOH ⇌ CH3COO- + H+
Add H+
Equilibrium shifts left to use up extra H+
But can not shift so long, because [CH3COO-] is very small because it is a weak acid and most remains undissociated.
So, [H+] increase. It is not truly buffered.

23. Assignment
Read Heath p. 603 and 604, then do Review# 1- 3.
Hebden # 132, 135, 136, 138, 142

24. Volumetric Analysis
Volumetric Analysis is used to find unknown concentrations by chemical reactions with known solutions.
The process usually involves titrations.

25. Hind + H2O ⇌ Ind- + H3O+ Definitions Indicator
a substance which changes colour at a specific pH.
usually the indicator itself is a weak acid or base in equilibrium
Hind + H2O ⇌ Ind- + H3O+
Phenolphthalein:
Phenol red:
When OH- ions are added:
they react with H3O+ ions to form water.
The equilibrium shifts right.
The solution turns pink for phenolphthalein; red for phenol red.

26. Definitions
Standard Solution
a solution of known concentration
can be prepared by:
Direct mass measurement.
Add a specific mass to a specific volume of solution
This creates a primary standard.
Not useful for gases or deliquescent substances, like NaOH.
Standardization against a known primary standard.

27. Definitions
Equivalence Point or Stoichiometric Point
when there are equivalent amounts of acid and base combined.
the ratio of H+ to OH- is 1:1.

28. Definitions
Endpoint or Transition Point
when the indicator changes colour.
titrations work best if equivalence point is near the
end-point.

29. Definitions
Neutralization
reaction of an acid and a base to form salt and water.

30. Assignment
Hebden #121,122

31. Sample Titration Calculations
Example 1: A 0. I 00 M standard solution of HCl is titrated against 20.0 mL KOH. The initial buret reading of HCI is
6.2 mL, and the final reading is 58.2 mL. Calculate the original [KOH].

32. Sample Titration Calculations
Example 2: HCl is standardized using Na2CO g
Na2 CO3 is dissolved in 250 mL water. The initial reading on the buret is 1.5mL of HCl. The final reading at equivalence is 31.5 mL. Find [HCl].

33. Assignment
Hebden #94,95,96,97
Lab 20C or 20G
Heath p.595#1-5 and p.615 challenge #2

34. Titration curves: strong Acid ∕ strong base
Consider the addition of 1.0 M NaOH to HCl (titration of HCl with NaOH).
Formula equation:
NaOH +HCl →NaCl + H2O
Ionic equation:
Na+ + OH- + H+ + Cl- → Na+ + Cl- + H2O
Net ionic equation:
OH-+ H+ →H2O

35. Titration Curves
Strong Acid/Strong Base

36. Calculating pH during titration
Just calculate the excess H+ or OH- by subtracting moles OH- added from moles H+ at the start.

37. Calculating pH during titration
Example 1: find the pH of the solution produced when 20.0mL of 0.150M HCl is mixed with 25.0mL of 0.100M NaOH.
Always start by writing the equation first.
Find moles H+ from acid
Find moles OH- from base
Excess moles remaining after reaction:
[H+]= pH

38. Calculating pH during titration
Example 2: During a titration mL of M HNO3 is added to mL of M KOH. Find the pH of the final mixture.
Equation:
Excess= (moles H+) - (moles OH-)

39. Titration Curves: Strong Base/Weak Acid
Consider the addition of NaOH to acetic acid.
Equation: ___________________________
At the equivalence point, the added NaOH has completely consumed the CH3COOH. Only NaCH3COO is present.
But here's the catch: the CH3COO-ion hydrolyses!
So, at the equivalence point [OH-] ˃ [H+] and the pH at equivalence is ˃7.0!
Indicator choice: one that changes colour at a pH˃7.0, like____________________

40. Titration Curves: Weak Base/Strong Acid
Consider the addition of ammonium hydroxide to 0.1M HCl
Equation: ___________________________
At the equivalence point, the HCl has been consumed by NH4OH. Only NH4Cl is present.
But here's the catch: the NH4 + ion hydrolyses:_______
So, at the equivalence point [OH-] ˂ [H+] and the pH at equivalence is ˂ 7.0!
Indicator choice: one that changes colour at a pH˂7.0, like____________________

41. Titration Curves: Weak Base/Weak Acid
Consider the addition of ammonium hydroxide to acetic acid
Equation: ___________________________
There is no apparent equivalence point.
hydrolysis occurs both above and below pH 7
the graph has no useful purpose.

42. Interpret the graph
Deduce as much information as you can from the shape and position of the curve.

43. Assignment
Heath p.614 problems 1,2,8,11,18

44. Hind + H2O ⇌ Ind- + H3O+ Indicators for Titrations
The transition point of the indicator should be near the equivalence point of the titration.
To determine the transition point; look at the Ka of the indicator.
Hind + H2O ⇌ Ind- + H3O+
Ka=
the endpoint = transition point occurs when
[Hind] = [Ind-] ) therefore transition occurs when the
[H3O+] = Ka!!

45. Select an indicator with a Ka = [H3O+] at the equivalence point

46. Indicators for Titrations
Example 1:
Find the Ka of litmus paper. The colour changes between
pH Assume endpoint is in the middle of the range, at
6.75.

47. Indicators for Titrations
Example 2:
for an indicator whose Ka = 2.6 x 10-5 , calculate the
pH at which it will change colour and identify the indicator.

48. Indicators for Titrations
Example 3:
Find a suitable indicator for
a. H2SO4 + NaOH
b. Ca(OH)2 +CH3COOH
c. HCI + NH3

49. Indicators for Titrations
Example 4:
Match the indicator to the titration:
a. C6H5COOH + NaOH A. Neutral red ( )
b. HNO3 +NaOH B. Bromcresol Green( )
c. NaHCO3 + HCl C. Phenolphthalein ( )

50. Assignment
Hebden # 108,109,110,111,114,116,117,120,122,125.

51. Real life acid- base chemistry
1.Buffers in blood (pH 7.4)
Relatively constant pH required for health (even life)
Blood contains a H2CO3/HCO3- buffer system
Equation:_____________________________
if H+ is added, the equilibrium shifts left.
If OH- is added, the equilibrium shifts right.
cells produce CO2 in a water solution:
CO2 + H2O → H2CO3
and H2CO3 ⇌ HCO3- + H+
thus, CO2 would cause a decrease in pH (and death!) were it not for the high [HCO3-] which resists the dissociation of H2CO3, so pH is maintained.

52. Real Life Acid-Base Chemistry
2. Buffers in food and household products
to lower acidity in jam or fruit juices manufacturers add ____________________
baby shampoo is buffered to pH of tears to keep eyes from stinging.

53. Real life acid- base chemistry
3. Acid rain
when coal or oil is burned oxides of S and N released. E.g.__________________
these are acid anhydrides, so they produce acidic solutions in the presence of water (rain)
SO2 + H2O ⇌H2SO3
SO3 + H2O→ H2SO4
SO2 is a big product of coal burning for electricity in the northeastern USA.

54. Real life acid- base chemistry
3. Acid rain
NO2 is produced mainly by lighting (97%) and by auto exhaust(3%)
2NO2 + H2O→ HNO3 + HNO 2
Normal rain has a pH of 6.4 (slight acidic because of dissolved CO2).
Acid rain has a pH of<5.6 (by definition).

55. Real Life Acid-Base Chemistry
Effects of Acid Rain
lakes
kills fish
leaching causes higher metal ion [ ].
e.g. Cu2+ , Al3+
forests
damage stomata causing water loss through leaves.
damage root hairs.
leaches away mineral in soil.
buildings
dissolves limestone and marble (CaCO3).
CaCO3 + 2 H+ → Ca2+ + H2O + CO2

56. Real Life Acid-Base Chemistry
Temporary Solutions:
i) spray lime (Ca(OH)2) or CaO into lakes to neutralize acid.
Ca(OH)2 + 2 H+ → Ca2+ + 2H2O
ii) coat buildings and statues with plastics, waxes, or paint.

57. Real Life Acid-Base Chemistry
Permanent Solutions:
Lower sulfur emissions
burn low sulfur fuels
Trap SO2 with CaO
CaO + SO2 → CaSO3
Dilute the SO2 by using a higher smokestack.
eg. Sudbury copper smelter has a 1250’ smokestack that cost $5.5 million (same height as Empire State Building).

58. Assignment
Hebden # 141,142,143
Heath p.614 I and A #9