1. Ch. 17 Buffers and Acid/Base Titration Lecture 3 – Titrations

2. Definitions: Amphoteric
Amphoteric means the substance can react either as an acid or a base.

3. Definitions: Amphoteric
Example: HCO3-
HCO3- + H2O  CO32- + H3O+
acid
HCO3- + H2O  H2CO3 + OH-
base

4. Definitions: Monoprotic and Polyprotic Acids
Monoprotic acids have only 1 hydrogen to donate.
Polyprotic acids have more than 1 hydrogen to donate.
These acids ionize in steps (1 hydrogen at a time)
Removing the first H automatically makes the second H harder to remove.
H2SO4 is a stronger acid than HSO4-

5. Definitions: Monoprotic and Polyprotic Acids
Write the dissociation of H2CO3 in steps:
H2CO3 → H+ + HCO3-
HCO3- → H+ + CO32-

6. Neutralization Reactions
Neutralization reactions (acid + base) ALWAYS produce a salt and water.
HCl + NaOH 
NaCl + H2O

7. Neutralization Reactions
Neutralization reactions (acid + base) ALWAYS produce a salt and water.
HF + KOH 
KF + H2O

8. Neutralization Reactions – Double Replacement Rxns
When an acid reacts with a base, it is called NEUTRALIZATION REACTION.
“Driving force” is the formation of water.
In a NEUTRALIZATION REACTION the products are water and a salt.
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)

9. Neutralization Reactions – Strong Acid & Strong Base
When a strong acid reacts with a base, the net ionic equation is…
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) 
Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq)  H2O (l)

10. Neutralization Reactions – Weak Acid & Strong Base
When a weak acid (i.e., acetic acid) reacts with a strong base, the net ionic equation is…
2 CH3COOH(aq) + Ba(OH)2 (aq) 
BaCH3COO (aq) + 2H2O (l)
Weak acids do not completely dissociate, so the complete ionic equation is:
2 CH3COOH(aq) + Ba2+ +2OH- (aq) 
Ba2+ + CH3COO- + H2O (l)
and the net ionic equation is:
2 CH3COOH(aq) + 2OH- (aq)  CH3COO- + H2O (l)

11. Neutralization Reactions – Strong Acid & Solid Base
When an acid reacts with a solid base, the net ionic equation is…
2 HCl (aq) + Mg(OH)2  MgCl H2O (l)
2 H+ (aq) + 2 Cl- (aq) + Mg(OH)2 (s) 
Mg2+ (aq) + 2 Cl- (aq) + 2 H2O (l)
2 H+ (aq) + Mg(OH)2 (s)  Mg2+ (aq) + 2 H2O (l)
(s)
(aq)

12. Titration
A common laboratory method used to determine the unknown concentration of a reactant.

13. Titration strong acid (HCl) with strong base (NaOH)
INITIAL: (No base added)
pH = - log [H+] (acid only)
Species present: H+ Cl**
(4)
**Species present only include additional species from pure water:
H2O  H+ + OH ●
(3)
(2)
(1)

14. Titration of a Strong Acid with a Strong Base
0.100 M NaOH added to 50.0 mL of M HCl.
1. The initial pH:
pH = -log [H+]
pH = -log(0.100M)
= 1.00

15. Titration strong acid (HCl) with strong base (NaOH)
“Buffering Region”:
pH slowly increasing due to buffering of common ion
Species present: H+ Cl Na+ (all added OH  consumed neutralizing acid)
(4) ●
(3)
(2)
(1)

16. Titration of a Strong Acid with a Strong Base
0.100 M NaOH added to 50.0 mL of M HCl.
2. As NaOH is added, pH slowly increases and increases more rapidly near the equivalence point

17. Titration strong acid (HCl) with strong base (NaOH)
Equivalence Point:
pH = [H+] = [OH ]
Species present: Cl Na+
(4) ●
(3)
(2)
(1)
**Species present only include additional species from pure water: H2O  H+ + OH

18. Titration strong acid (HCl) with strong base (NaOH)
DEFINITIONS:
Equivalence Point:
Moles of added H+ = moles of added OH-
End point: observable point indicated by a color change that occurs when the equivalence point is reached.
(4) ●
(3)
(2)
(1)
**Species present only include additional species from pure water: H2O  H+ + OH

19. Titration of a Strong Acid with a Strong Base
0.100 M NaOH added to 50.0 mL of M HCl.
3. At the equivalence point,
Moles added H+ = moles added OH-
pH = 7.00
Solution contains only H2O and NaCl.

20. Titration strong acid (HCl) with strong base (NaOH)
Past the equivalence point:
pH increases as more OH is added
Species present: Cl Na+ OH (all added H + is neutralized)
(4) ●
(3)
(2)
(1)
**Species present only include additional species from pure water: H2O  H+ + OH

21. Titration of a Strong Acid with a Strong Base
0.100 M NaOH added to 50.0 mL of M HCl.
4. As more base is added, the increase in pH again levels off.
The solution has excess OH-.

22. At the equivalence point:
Titration - Problems
A 15 mL solution of HCl is titrated with M NaOH. What was the concentration of the HCl?
At the equivalence point:
Moles H+ = Moles OH–
MaVa = MbVb
Where:
M = concentration of acid [H+] or base [OH]
V = volume of acid or base
Ma x 15 mL = M OH x 10 mL
Ma = 0.25 M H+

23. Titration - Problems
A 35 mL solution of HCl is titrated with M Ca(OH)2. What was the concentration of the HCl?
Moles H+ = Moles OH-
MaVa = MbVb
Where:
M = concentration of acid [H+] or base [OH]
V = volume of acid or base
Ma x 35 mL = (2 x M OH) x 10 mL
Ma = M H+

24. Titration weak acid (HA) with strong base (NaOH)
Review:
What is the difference between strong and weak acids (or bases)?
Weak acids and bases do not completely dissociate in water.

25. Titration weak acid (HA) with strong base (NaOH)
Review:
Write the balanced equation and net ionic equation for the neutralization reaction between a weak acid (HA) and a strong base (NaOH):
HA(aq) + NaOH(aq)  H2O(aq) + NaA(aq)
HA(aq) + OH  H2O(aq) + A

26. Titration weak acid (HA) with strong base (NaOH)
(4)
(1) INITIAL:
(No base added)
pH calculated using ICE box
Species present:
HA H+ A**
(3)
(2)
(1)
**Species present only include additional species from pure water:
H2O  H+ + OH

27. Titration weak acid (HA) with strong base (NaOH)
(4)
(2) Buffering Region:
pH slowly increasing due to buffering of common ion
At ½ Veq, [HA] = [A] and pH = pKa
Species present:
HA H+ A Na+
(all added OH consumed neutralizing acid)
(3)
pKa
(2)
(1)
**Species present only include additional species from pure water:
H2O  H+ + OH
Veq

28. Titration Curve Titration of a weak acid by a strong base
HALF EQUIVALENCE POINT: at ½ Veq
[HA] = [A-]
pH = pKa

29. Titration Curve Titration of a weak base (B) by a strong acid (HA)
HALF EQUIVALENCE POINT: at ½ Veq
pH = pKa (pKa for weak acid, BH+)

30. Titration weak acid (HA) with strong base (NaOH)
(4)
(3) Equivalence Point:
pH > 7
Since HA is a weak acid, A reacts with water to form hydroxide ions:
A + H2O  HA + OH 
Species present:
A Na+ OH 
and very small quantities of HA
(3)
(2)
(1)
**Species present only include additional species from pure water:
H2O  H+ + OH

31. Titration weak acid (HA) with strong base (NaOH)
(4)
Past neutralization:
pH increases as more OH is added
Species present:
A Na+ OH 
and very small quantities of HA
(3)
(2)
(1)
**Species present only include additional species from pure water:
H2O  H+ + OH

32. Titration (acid with strong base)

33. Titration (base with strong acid)

34. Comparison of Strong-Strong v. Weak-Strong TitrationspH
Strong Acid-Strong Base
Weak Acid – Strong Base
Initial
At Equivalence Point

35. Comparison of Strong-Strong v. Weak-Strong TitrationspH
Strong Acid-Strong Base
Weak Acid – Strong Base
Initial
Lower pH
Higher pH
At Equivalence Point
pH = 7
Weak acid-strong base: pH >7
(A- + H2O  OH- + HA)
Weak base-strong acid:
pH<7
(BH+ + H2O  H3O+ + B)
(A = anion in acid
B = base)

36. Monoprotic v. Polyprotic Acids
Monoprotic acids:
Acid that can donate one (1) hydrogen ion (proton) (Ex. HCl, HNO3)
Polyprotic acids:
Acid that can donate more than one hydrogen ion (proton) (Ex. H2SO4, H3PO4)

37. Polyprotic Acid Titration of H2SO4
HSO4SO42-
H2SO4HSO4
1st eq. pt.
2nd eq. pt.

38. Selection of Indicators
Pick an indicator that changes color very close to the pH at the equivalence point for the titration (usually a little before).

39. Indicator Transitions
Low pH color
Transition pH range
High pH color
Gentian violet (Methyl violet 10B)
yellow
0.0–2.0
blue-violet
Leucomalachite green (first transition)
green
Leucomalachite green (second transition)
11.6–14
colorless
Thymol blue (first transition)
red
1.2–2.8
Thymol blue (second transition)
8.0–9.6
blue
Methyl yellow
2.9–4.0
Bromophenol blue
3.0–4.6
purple
Congo red
3.0–5.0
Methyl orange
3.1–4.4
orange
Bromocresol green
3.8–5.4
Methyl red
4.4–6.2
4.5–5.2
Azolitmin
4.5–8.3
Bromocresol purple
5.2–6.8
Bromothymol blue
6.0–7.6
Phenol red
6.8–8.4
Neutral red
6.8–8.0
Naphtholphthalein
colorless to reddish
7.3–8.7
greenish to blue
Cresol Red
7.2–8.8
reddish-purple
Phenolphthalein
8.3–10.0
fuchsia
Thymolphthalein
9.3–10.5
Alizarine Yellow R
10.2–12.0
Litmus
Source: Wikipedia

40. Measuring pH
Indicators change color at a different pH. INTERPRETING GRAPHS a. Identify Which indicator changes color in a solution with a pH of 2? b. Compare and Contrast What do you notice about the range over which each indicator changes color? c. Apply Concepts Which indicator would you choose to show that a solution has changed from pH 3 to pH 5?

41. Different indicators have different end points.
Measuring pH
Different indicators have different end points.
Choose the indicator with the end pt. closest to the titration equivalence point.
Indicators change color at a different pH. INTERPRETING GRAPHS a. Identify Which indicator changes color in a solution with a pH of 2? b. Compare and Contrast What do you notice about the range over which each indicator changes color? c. Apply Concepts Which indicator would you choose to show that a solution has changed from pH 3 to pH 5?