Introduction to Volumetric Analysis for Basic Chemistry

1. VOLUMETRIC ANALYSIS

2. LESSON OUTCOMES
• Types of volumetric analysis
• Principles of titration: equivalence point and
end point
• Back-titration
• Acid base titration curve

3. TYPES OF VOLUMETRIC ANALYSIS

4. VOLUMETRIC ANALYSIS
• A method in quantitative chemical analysis in which the
amount of a substance is determined by the
measurement of the volume that the substance occupies
• It is commonly used to determine the unknown
concentration of a known reactant
• Volumetric analysis is often referred to as titration, a
laboratory technique in which one substance of known
concentration and volume is used to react with another
substance of unknown concentration.

5. TYPES OF TITRATION
Types Example
Acid-Base Determine % acetic acid in
vinegar
Complexometric Water hardness (Determine
Ca2+ in water)
Precipitation Determine of Cl- in water
Redox Quantify hydrogen peroxide
(H2O2)

6. Titration- Preparation
 Two solutions are used:
 The solution of unknown concentration.
 The solution of known concentration which is also known as the standard
solution.
 Write a balanced equation for the reaction between
your two chemicals.
 Clean all glassware to be used with distilled water. The
pipettes and burettes will be rinsed with the
solutions you are adding to them.

7. Titration Set up
 The burette is attached to a clamp
stand above a conical flask.
 The burette is filled with one of the
solutions (in this case a yellow
standard solution).
 A pipette is used to measure an
aliquot of the other solution (in this
case a purple solution of unknown
concentration) into the conical
flask.
 Prepare a number of flasks for
repeat tests.
 Last, an indicator is added to the
conical flask.

8. The Titration Process
 Read the initial level of liquid in the
burette
 Turn the tap to start pouring out liquid of
the burette into the flask. Swirl the flask
continuously. When the indicator begins to
change colour slow the flow.
 When the colour changes permanently,
stop the flow and read the final volume.
The volume change needs to be calculated
(and written down).This volume is called a
titre.
 Repeat the titration with a new flask now
that you know the ‘rough’ volume required.
Repeat until you get precise results.

9. Acid-Base Titration
• When the reaction involves an acid and a base, the
method is referred to as an acid-base titration.
• The aim of acid-base titration is to determine the
volume of two solutions that will exactly react with one
another
• A known volume of a solution is placed in a conical flask.
The other solution is added from a burette until there is
enough of it to completely react with the solution originally
in the flask
• However, most acids, bases and their salts are colourless.
Therefore an indicator is needed to determine the end-
point and equivalence point of the titration. Acid-base
indicator would change colour at the end-point of a
titration

10. Procedure of Acid- Base Titration

11. Definition of terms
Volumetric Analysis
 Measurement of volume of a solution of known
concentration, which is used to determine the concentration
of the analyte
VolumetricTitrimetry
 Quantitative chemical analysis which determines volume of a
solution of accurately known concentration required to react
quantitatively with the analyte (whose concentration to be
determined).
 The volume of titrant required to just completely react with
the analyte is theTITRE

12. What is a titration
A procedure of carefully controlled
addition of reagent (titrant) to an analyte
until the reaction between the two is
judged complete
Known concentration
which is called
standard
Usually in
burette
Usually in a
conical flask

13. Successful Volumetric Titration
must employed…..
 Reaction must be stoichiometric, well defined
reaction between titrant and analyte.
 Reaction should be rapid.
 Reaction should have no side reaction, no
interference from other foreign substances.
 Must have some indication of end of reaction, such
as color change, sudden increase in pH, zero
conductivity, etc.
 Known relationship between endpoint and
equivalence point.

14. Direct and Back Titration
1. DirectTitration
A titration process where a titrant (standard
solution) is added to the analyte until the
reaction is between analyte and titrant is judged
to be complete.
2. BackTitration
A titration process in which the excess standard
solution used to react with analyte is determined
by titration with a second standard solution.

15. End Point VS Equivalence Point
END POINT EQUIVALENCE POINT
The point at which the reaction is
observed to be completed
The point at which an equivalent or
stoichiometric amount of titrant is
added to the analyte
The end point signal frequently occurs at
some point other than the equivalence
point which tells the analyst to stop
addingTITRANT and record the volume.
The point at which the reaction is
complete
The selected indicator should change
color very near to the equivalent point.
Theoretically at the equivalence
point we can calculate the amount
of titrant that is required to react
EXACTLY with the amount of
analyte present.

16. Good End Point
Bad End Point
(Overshot)

17. Sometime direct titration are not feasible due
to…..
1.Reaction kinetic or the reaction rate is slow.
2.No suitable indicator in the direct titration.
3.The color change is slow or delay.
4.The end point is far from the equivalent point.

18.  In a simple acid-base titration, a base (reagent) is
added in a known quantity – greater than the amount
required for acid neutralization.
 Acid and base reacts completely.
 The remaining base is titrated with a standard acid.
 The system has gone from beingACID, past the
equivalent point to the BASIC (excess base), and back
to the equivalence point again.
 The final titration to the equivalence point is called
BACKTITRATION.

19.  Example, the titration of insoluble organic acid with
NaOH is not practical because the reaction is slow.
 To overcome it add NaOH in excess and allow the
reaction to reach completion and then titrate the
excess NaOH with a standard solution of HCl.
 The system has gone from being ACID , past the
equivalence point to the BASIC side (excess base),
and then back to the equivalence point.
 The final titration to the equivalence point is called a
BACKTITRATION.

20.  The analyte may be in solid form
 The analyte reacts slowly with the titrant in
direct/forward titration
 The analyte may contain impurities which
may interfere with direct titration

21. 150.0 mL of 0.2105 M nitric acid was added in excess
to 1.3415 g calcium carbonate.The excess acid was
back titrated with 0.1055 M sodium hydroxide. It
required 75.5 mL of the base to reach the end point.
Calculate the percentage (w/w) of calcium carbonate
in the sample.
Examples On
Back Titration Method

22. First write a balance equation for the above
reactions.
2HNO3 + CaCO3  Ca(NO3)2 + CO2 + H2O ------ 1
HNO3 + NaOH  NaNO3 + H2O ------- 2
From Equations above:
2 mole HNO3 required 1 mole CaCO3
1 mole HNO3 required 1 mole NaOH
Initial mole of acid
= Molarity x volume (L)
= 0.2105 mol/L x 0.150 L
= 0.031575 mol

23. Remaining/excess acid during back titration.
Mole of excess acid
= molarity x volume (L)
= 0.1055 mol/L x 0.0755 L
= 0.007965 mol acid.
Mole of CaCO3 in sample= ½ x mmole acid
= ½ x 0.02361
= 0.011805 mol
Mole of acid reacted with CaCO3
= Initial – remaining/excess
= ( 0.031575 – 0.007.965 ) mol
= 0.02361 mol

24. Mass of CaCO3 in sample = mole x molar mass
= 0.011805 mol x 100 g/mol
= 1.1805 g.
% (w/w) of CaCO3 in sample
= 87.99 % (w/w)
%
100
sample
of
weight
CaCO
weight 3
x

%
100
1.3415
1.1805
x


25. Indicator and Choice of Indicator
Acid –Base Indicators:
• The acid-base indicator function by changing colour just
after the equivalence point of a titration; this colour
change is called the end point.
• The end point is most often detected visually.
• Acid-base indicator are usually weak organic acids or
bases that dissociate partially in water with the
undissociated molecules have different color from their
ions. (eg: Undissociated molecule of phenolphthalein is
colourless & it anion is pink)
• Indicators can be monoprotic (HIn) or diprotic (H2In) acids.

26. • The acid form of an indicator is usually coloured; when it
loses a proton resulting in anion (In-) , or base form of the
indicator, exhibiting different colour.

27. Acid Base Indicators
Common Name Transition range Colour Change
ACID BASE
Crystal violet 0.1 – 1.5 Yellow Blue
Thymol blue 1.2 – 2.8 Red Yellow
Methyl yellow 2.9 – 4.0 Red Yellow
Methyl orange 3.1 – 4.4 Red Orange
Bromocresol green 3.8 – 5.4 Yellow Blue
Methyl red 4.2 – 6.3 Red Yellow
Chlorophenyl red 4.5 – 6.4 Yellow Red
Bromothymol blue 6.2 – 7.6 Yellow Blue
Phenol red 6.8 – 8.4 Yellow Red
Thymol blue 8.0 – 9.6 Yellow Blue
Phenolpthalein 8.3 – 10.0 Colourless Red
Alizarin yellow 10.0 – 12.0 Colourless Yellow

28. Choosing a Titrant
• In theory, any strong acid or strong base can be used as titrant.
• The reason for this is that most reaction involving a strong acid or
a strong base is quantitative.
Strong Acid Titrant Weak Acid Titrant
• Hydrochloric acid (HCl)
• Nitric acid (HNO3)
• Perchloric acid (HClO4)
• Phosphoric acid (H3PO4)
• Acetic acid (CH3COOH)
• Ammonium ion (NH4-)
• Hydrogen fluoride (HF)
• Carbonic acid (H2CO3)
• Nitrous acid (HNO2)
• Hydrogen sulphide (H2S)
• Hydrogen cyanide (HCN
Acid Titrant

29. Strong Base Titrant Weak Base Titrant
Sodium hydroxide, NaOH
Potassium hydroxide,
KOH
Magnesium hydroxide,
Mg(OH)2
Barium hydroxide ,
Ba(OH)2
Ammonium hydroxide,
NH4OH
Amine acetate.
Carbonate, CO3
2-
Fluoride ion,F-
Sodium
carbonate,NaCO3
Base Titrant

30. Acid- Base Titration Curve

33. Acid-Base
Titration
Strong Acid-
Strong Base
Weak Acid –
Weak Base
Strong Acid -
Weak Base
Strong Base –
Weak Acid

34. The Titration (or pH) Curve
• The change in pH during an acid-base titration
can be followed by measuring the pH of the
mixture using a pH meter
• The change is then plotted against the volume of
base (or acid) added from the burette
• These titration curves allow us to choose the
most suitable indicator for the particular titration

35. The Relationship Between pH & pOH
pH = -log [H+]
[H+] = [OH-]
[H+] > [OH-]
[H+] < [OH-]
Solution Is
neutral
acidic
basic
[H+] = 1 x 10-7
[H+] > 1 x 10-7
[H+] < 1 x 10-7
pH = 7
pH < 7
pH > 7
At 250C
pH [H+]
[H+][OH-]= Kw=1.0 x 10-14
-log [H+] – log [OH-] = 14.00
pH + pOH= 14.00

36. Titration curve for the titration of 25.00 mL of 0.20 M HCl with the 0.20 M NaOH
• Sudden sharp changes of pH
• Salt formed: NaCl
• Neutral salt: So pH is 7
Slow change of pH
Strong acid with strong base
Buffer Region
NaOH (aq) + HCl (aq) H2O (l) + NaCl (aq)

37. • Before addition of NaOH, the pH=1. As NaOH is added, the pH
increases gradually
• There is a sharp increase in pH (from 4 to 11) slightly before and after
equivalence point (pH=7)
• Beyond equivalence point, pH increases gradually as more NaOH
added
• The suitable indicators are:
Indicator pH Range Colour Change
Methyl red 4.2 – 6.3 Red- Yellow
Chlorophenyl red 4.8 – 6.4 Yellow- Red
Bromothymol blue 6.0 – 7.6 Yellow- Blue
Phenol red 6.4 – 8.4 Yellow -Red
Cresol red 7.2-8.8 Yellow -Red
Phenolpthalein 8.3- 10.00 Colourless- pink

38. HCl (aq) + NH3 (aq) NH4Cl (aq)
NH4
+ (aq) + H2O (l) NH3 (aq) + H+ (aq)
At equivalence point (pH < 7):
Weak base with Strong acid
• Sharp decrease in pH
• pH is less than 7
(usually around 5-6)
• Salt formed is acidic

39. • Before addition of HCl, the pH=11. When HCl is added,
the pH increases gradually as more NH3 neutralised and
the solution becomes more acidic.
• There is sharp decrease in pH (from 7 to 3) slightly
before and after equivalence point.
• Beyond equivalence point, pH decreases gradually as
more HCl is added
• The suitable indicators are:
Indicator pH Range Colour Change
Methyl Orange 3.1-4.4 Red-Yellow
Methyl Red 4.2-6.3 Red-Yellow
Chlorophenol blue 4.8-6.4 Yellow-Red

40. 41
CH3COOH (aq) + NaOH (aq) CH3COONa (aq) + H2O (l)
CH3COO- (aq) + H2O (l) OH- (aq) + CH3COOH (aq)
At equivalence point (pH > 7):
Weak acid with Strong base
• pH changes rapidly
• pH is higher than 7
(usually around 8-9)
• Salt formed is alkaline

41. • Before addition of NaOH, pH≈ 2.9 (ethanoic acid is a
weak acid). When NaOH is added, the pH increases
gradually as more CH3COOH is neutralised and the
solution becomes less acidic
• There is a sharp increase in pH (from 7 to 11) slightly
before and after equivalence point
• Beyond the equivalence point, the pH increases
gradually as more NaOH is added
• The suitable indicators are:
Indicator pH Range Colour Change
Cresol Red 7.2-8.8 Yellow-Red
Phenolpthalein 8.3- 10.00 Colourless- pink

42. Weak Acid- Weak Base
• Titrations involving a weak acid with a weak
base are not normally done
• This is because the equivalence point cannot
be accurately observed