1. Principles of Neutralization Titration
Version 2012 Updated on Copyright © All rights reserved
Dong-Sun Lee, Prof., Ph.D. Chemistry, Seoul Women’s University
Chapter 12
Principles of Neutralization Titration

2. Standard solutions for Acid-Base Titrations
Analyte Titrant vs Standard solution
N V = N’V’
The standard reagents used in acid-base titrations are always strong acids or strong bases, most commonly HCl, HClO4, H2SO4, NaOH, KOH. Weak acids and bases are never used as standard reagents because they react incompletely with analyte.
N= unknown
V= measure
N’= known
V’= known

3. Primary standards for standardizing
Acids Potassium acid phthalate
Sulfamic acid ( H2NSO3H)
HCl
Potassium hydrogen iodate
Bases TRIS(hydroxymethylaminomethane)
Sodium carbonate
Borax (= sodium tetraborate )
HgO 

4. General goal and procedure of acid-base titration 
1) Standardization
2) Determination
3) Titration curve
4) Interpret titration curve ,
understand what is
happening during titration
Primary standard solution
known normality(N)
known volume(V)
Acid or base
standard solution
unknown normality(N’)
unknown volume(V’)
known normality(N’)
Analyte Sample
solution
unknown Ns
known Vs
Plots of pH versus volume of titrant

5. Finding the end point 1) Indicator
2) Potentiometry : Titration curve ( pH or mV vs Va )
1st derivative titration curve
2nd derivative titration curve
Gran plot
3) Conductometry
4) Spectrometry

6. Indicator Ex. Thymol blue H2In  HIn–  In2 – Transition range
An acid-base indicator is itself an acid or base whose different protonated species have different colors.
Ex. Thymol blue
H2In  HIn–  In2 –
Red Yellow Blue
H2In  H+ + HIn– pKa1 = 1.7
pH = pKa1 + log [HIn– ]/[H2In]
if [HIn– ]/[H2In]  1: pH = Red color
[HIn– ]/[H2In] = 1: pH = pKa1 = Orange color
[HIn– ]/[H2In]  10: pH = Yellow color
HIn–  H In2 – pKa2 = 8.9
Transition
range
pKa11
The approximate pH transition range of most acid-type indicator is roughly pKa 1

7. Choosing an indicator PP 8.0~9.0 MR 4.8~6.0 MO 3.1~4.4 Phenolphthalein
Volume of M HCl (ml) 5 10 15 20 25 30 35
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00
PP 8.0~9.0
MR ~6.0
MO ~4.4 pH
Phenolphthalein
acid form :
colorless
transition range(pH):
8.0~9.0
Phenolphthalein
base form :
pink
The calculated pH titration curve for the
titration of 10.00ml of M Na2CO3
with M HCl.

8. Calculated titration curve for the reaction of 100 mL of 0
Calculated titration curve for the reaction of 100 mL of M base (pKb = 5.00) with M HCl.

9. Indicator color as a function of pH (pKa=5.0)

12. Experiments. Standardization of 0.1000N NaOH
Primary standard KHP
204.22g/1000ml=1.0000N
0.35g/25ml = xN
x = N
Titration of
25.00ml of
KHP with
NaOH
End point=18.50ml
N×25.00ml= x N×18.50ml
x= N

13. Calculation of concentration (Finding of NaOH concentration)
HOOCC6H4COOK (mw )  NaOH (mw 40.01)  1 Eq wt.
g  1 N × 1000 mL
mg  0.1 N × 1 mL
a g  x N’ × Ve mL
 x N’ = (a g × 1000 mL) / ( g × Ve mL)

16. Gran plot
HA = H+ + A–
Ka = [H+] fH+ [A–]fA– / [HA]fHA
HA vs NaOH
[A–] = (moles of OH– delivered) / (total volume)
= VbFb / (Vb + Va)
[HA] = (original moles of HA – moles of OH –) / (total volume)
= (VaFa – VbFb ) / (Vb + Va)
Ka = [H+] f H+ VbFb fA– / (VaFa – VbFb ) fHA
[H+] fH+ Vb = (fHA / fA– ) Ka{(VaFa – VbFb )/ Fb}
10–pH Vb = (fHA / fA– ) Ka(Ve – Vb)
= –(fHA / fA– ) Ka Vb + (fHA / fA– ) KaVe
Slope =–(fHA / fA– ) Ka
10–pH Vb Ve
Vb(l)

17. Gran plot for the first equivalence point
Gran plot for the first equivalence point. This plot gives an estimate of Ve that differs from that in Figure 12.7 by 0.2μL (88.4 versus 88.2μL). The last 10-20% of volume prior to Ve is normally used form a Gran plot.

18. Conductometric end-point detection

19. Titration of strong acid with strong baseV a
(mL) 2 4 6 8 10 12 14 16 18 20 B C D pH
Neutralization :
HCl + NaOH  NaCl + H2O
Ve mL×0.1000M = mL × M
Ve = mL M
NaOH
50.00 mL
0.1000M
HCl
Calculated titration curve for the reaction of
50.00 mL of M NaOH with M HCl.

20. A. Before the beginning of titration : Va = 0.00 mL. NaOH  Na+ + OH–,
M [OH–] = M
[H+] = 1.00×10–14/ 2.000×10–2
pH = 12.30
B. Before the equivalence point : 0 < Va < Ve
remaining NaOH
Initial NaOH amount =F × Vi
added HCl amount = used NaOH amount = F× Va
Remaining NaOH amount =
[OH–] = {(Ve – Va )/ Ve}FNaOH {Vi /(Vi + Va)}
V a Ve
Ex. Va = 3.00 mL
[OH–] = {(10.00 – 3.00)/10.00}( ){50.00 /( )} = M
[H+] = 1.00×10–14/ = 7.58×10– pH = 12.12

21. C. At the equivalence point : Va = Ve H2O = H+ + OH–,
[H+] = 1.00×10– pH = 7.00
D. After the equivalence point : Va > Ve
excess HCl
added HCl amount = FHCl× Va
initial NaOH amount = used HCl = F × Vi
Excess HCl amount = FHCl(Va–Vi)
= FHCl (Va–Ve)
[H+] = FHCl {(Va– Ve) /(Vi + Va)}
NaOH Vi
V a Ve
Ex. Va = mL
[H+] = (0.1000) { –10.00) /( )} = 8.26×10– pH = 3.08

22. Titration curves for NaOH with HCl.
50.0 mL of M NaOH with M HCl.
50.00 mL of M NaOH with M HCl.

23. Titration curves for HCl with NaOH.
A mL of M HCl
with M NaOH.
50.00 mL of M HCl
with M NaOH.

24. The effects of titrant and analyte concentrations on neutralization titration curves.
For the titration with diluted concentrations, the change in pH in equivalence point region is markedly less than concentrated solution V a
of NaOH (mL) 5 10 15 20 25 30 35
0.00
1.00
2.00
3.00
4.00
5.00
6.00
7.00
8.00
9.00
10.00
11.00
12.00
13.00
14.00 A B pH
Changes in pH during the totration of
strong acid with strong base.
A: 50 mL of M HCl vs M NaOH
B: 50 mL of M HCl vs M NaOH

25. Weak acid titrated with strong base
  
Weak acid titrated with strong base
1) Titration reaction :
CH3COOH + NaOH  CH3COO– Na+ + H2O
strong + weak  complete reaction
K = 1/Kb = 1.76 ×109
2) Calculation of equivalence point :
[CH3COOH] Vi mL = [NaOH] Ve mL
M×5.00 mL = M × Ve mL
Ve= 5.00 mL
3) Titration curve : V a
of NaOH(mL) 1 2 3 4 5 6 7 8
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 pH A B
Changes in pH during the titration of a weak acid with a strong base.
A: mL of M HOAC vs M NaOH
B: ml of M HOAC vs M NaOH

26. Titration curves for the titration of acetic acid with NaOH.
A M acetic acid
with M NaOH.
B M acetic acid
with M NaOH.

27. A. Before adding any titrant : Va = 0.00   
HAC = H+ + AC–
[H+] = KaF =  1.75 ×10–5 ×0.1000
= × 10–3
pH= 2.88
   V a
of M NaOH (mL) 1 2 3 4 5 6 7 pH
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 A
0.1000M
NaOH
0.1000M
HAC
25.00ml
Titration curve.
M HAC vs M NaOH.

28. B. Before equivalence point : 0 < Va < Ve   
HAC + NaOH  H2O + Na+AC–
[H+] = Ka[HAC] / [AC–]
initial HAC amount = F×Ve
– added NaOH amount
= used amount HAC= F×Va
Remaining HAC amount =F(Vi–Va)
[HAC] = {(F×Ve) – (F×Va)}/(Vi +Va)
[AC–] = (F×Va) / (Vi +Va)
pH = pKa + log [AC–] /[HAC]
at Va = Ve/ pH= pKa = 4.76
   V a
of M NaOH (mL) 1 2 3 4 5 6 7 pH
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 B
Ve/2 Vi
Titration curve.
M HAC vs M NaOH. Va Ve

29. C. Equivalence point Va = Ve
  
C. Equivalence point Va = Ve
HAC + NaOH  H2O + Na+AC–
initial
final
AC– = HAC + OH–
[OH–] = KbF’ =  KwF’ /Ka
F’=( F×Vi) / (Vi +Va) V a
of M NaOH (mL) 1 2 3 4 5 6 7 pH
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 C
Titration curve.
M HAC vs M NaOH.

30. D. After equivalence point Va> Ve
  
D. After equivalence point Va> Ve
added NaOH amount = F×Va
initial HAC amount
= used NaOH amount = F×Ve
Excess NaOH amount = F×(Va –Ve)
[OH–] = FNaOH ×(Va– Vi) / (Vi +Va) V a
of M NaOH (mL) 1 2 3 4 5 6 7 pH
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 D Vi
Titration curve.
M HAC vs M NaOH. Va Ve

31. The effect of acid strength (dissociation constant) on titration curves. Each curve represents the titration of mL of M acid with M base.

32. Weak base titrated with strong acid
1) Titration reaction :
NH3 + HCl  NH4+Cl–
2) Equivalence point :
NH HCl
M×20.00 mL = M×Ve mL
Ve = mL
3) Titration curve :
A. Initial Va = 0
NH3 + H2O = NH OH–
Kb = Kw / Ka = [NH4+][OH– ]/[NH3]
1.79×10–5 = [OH– ]2 / –[OH– ]
[OH– ]=  KbF =1.4×10–3
pH = 11.15
0.1000M HCl volume (mL) 10 20 30 40 50 pH
0.00
2.00
4.00
6.00
8.00
10.00
12.00
14.00 A
The calculated titration curve for
the titration of mL of M
ammonia with M HCl.

33. B Ve/2 B. Before the equivalence point 0<Va< Ve
Major constituents: NH3 + NH4+Cl–
[OH]= Kb [NH3]/ [NH4+]
initial NH3 amount = F×Ve
– added HCl amount
= used amount NH3 = F×Va
Remaining NH3 amount =F(Vi–Va)
[NH3] = {(F×Ve) – (F×Va)}/(Vi +Va)
[NH4+] = (F×Va) / (Vi +Va)
pH = pKb + log [NH4+] / [NH3]
at Va = Ve/ pOH= pKb = 4.74
pH = 9.26 B
Ve/2 Vi Va Ve

34. C. Equivalence point Va = Ve
NH3 + HCl  NH4+Cl–
initial
final
NH4+ = NH3 + H+
[H+] = KaF’
F’=( F×Vi) / (Vi +Va) C

35. D. After equivalence point Va> Ve
added HCl amount = F×Va
initial NH3 amount
= used HCl amount = F×Ve
Excess HCl amount = F×(Va –Ve)
[H+] = FHCl ×(Va– Ve) / (Vi +Va) D Vi Ve Va

36. The effect of base strength (dissociation constant) on titration curves. Each curve represents the titration of ml of M base with M HCl.

37. Comparison of Weak Acid/ Base with Strong Base/Acid
Titration of Weak Base with Strong Acid
Comparison of Weak Acid/ Base with Strong Base/Acid
[H+] = KaF’
F’=( F×Vi) / (Vi +Va)
[OH–] = KbF’ =  KwF’ /Ka
Equivalence point
After equivalence point
(Va>Ve)
pH = pKb + log[NH4+]/[NH3]
pH = pKa + log [A–] /[HA]
Before the equivalence point (0<V a<V e)
[OH-] =  KbF =1.4×10–3
[H+] = KaF
Initial
B + H2O → BH+ + OH-
HA + OH- → H2O + A-
Titration reaction
Weak Base with Strong Base
Weak Acid with Strong Base
(Vi + Va) }
(Va– Ve) {
FHCl =
[H+]
FNaOH
[OH-]

38. Determining the pK values for amino acids
Amino acids contain both an acidic and a basic group.
alanine
All naturally occurring amino acids are left-handed (L) form.
The amine group behaves as a base, while the carboxyl group acts as an acid.
Aspartic acid

39. Curves for the titration of 20. 00ml of 0. 1000M alanine with 0
Curves for the titration of 20.00ml of M alanine with M NaOH and M HCl. Note that the zwitterion is present before any acid or base has been added. Adding acid protonates the carboxylate group with a pKa of Adding base reacts with the protonated amine group with a pKa of 9.89.

40. Plots of relative amounts of acetic acid and acetate ion during a titration.

41. Acid base titration in non-aqueous media
HOOCCHNH2 + HClO  HOOCCHNH3+ClO4–
R CH3COOH R
(solvent)
HClO4 + CH3COOK  CH3COOH + K+ClO4–
CH3COOH(solvent)
Amino acid
in acetic
acid +
0.020M HClO4
in acetic acid
0.010M CH3COOK

42. Titrants used in non-aqueous titrimetry
Acidic titrants Perchloric acid
p- Toluenesulfonic acid
2,4-Dinitrobenzenesulfonic acid
Basic titrants Tetrabutylammonium hydroxide
Sodium acetate
Potassium methoxide
Sodium aminoethoxide

43. Selected solvents for non-aqueous titration
Solvent Autoprotolysis constant Dielectric
(pKHS) constant
Amphiprotic Glacial acetic acid
Ethylenediamine
Methanol
Aprotic or basic Dimethylformamide
Benzene
Methyl isobutylketone
Pyridine
Dioxane
n-Hexane

44. neither acid is completely dissociated.
Acid and base strengths that are not distinguished in aqueous solution may be distinguishable in non-aqueous solvents.
Ex. Perchloric acid is a stronger acid than hydrochloric acid in acetic acid solvent,
neither acid is completely dissociated.
HClO CH3COOH = ClO4– CH3COOH K = 1.3×10–5
strong acid strong base weak base weak acid
HCl CH3COOH = Cl– CH3COOH K = 5.8×10–8
Differentiate acidity or basicity of different acids or bases
differentiating solvent for acids …… acetic acid, isobutyl ketone
differentiating sovent for bases …… ammonia, pyridine

45. Hammett acidity function, Ho, for aqueous solutions of acids.
Acid Name Ho
H2SO4 (100%) Sulfuric acid –11.93
H2SO4 SO Fuming sulfuric acid –14.14
H SO3F Fluorosulfuric acid –15.07
H SO3F+10%SbF5 Super acid –18.94
H SO3F+7%SbF53SO –19.35
Hammett acidity function, Ho, for aqueous solutions of acids.

46. Titration of a mixture of acids with tetrabutylammonium hydroxide in methyl isobutyl ketone solvent shows that the order of acid strength is HClO4 > HCl > 2-hydroxybenzoic acid > acetic acid > hydroxybenzene. Measurements were made with a glass electrode and a platimum reference electrode. The ordinate is proportional to pH, with increasing pH as the potential becomes more positive.

47. Summary Acid-Base titration Titration curve Indicator
Detection of end-point
Potentiometry
Gran plot
Conductometry
Non-aqueous titration
Hammett acidity function, Ho