THE FEDERAL POLYTECHNIC, EDE

SCHOOL OF SCIENCE AND TECHNOLOGY

DEPARTMENT OF CHEMICAL SCIENCES AND TECHNOLOGY

STC 124 [Analytical chemistry aspect]

TABLE OF CONTENT

1. EXPERIMENT 1: CALIBRATION OF A PIPETTE

2. EXPERIMENT 3: SEPARATION OF MIXTURE (P/Thin-layer Chromatography)

3. EXPERIMENT 4: Preparation of 500ml standard (0.1 M HCl and 0.1 M NaOH)

Solutions

4. EXPERIMENT 5: (STANDARDIZATION) ESTIMATION OF WORKING HCl

SOLUTION CONCENTRATION

5. EXPERIMENT 7: DETERMINATION OF PURITY OF ASPIRIN SAMPLES

6. EXPERIMENT 10 (REDOX TITRATION): DETERMINATION OF FeSO4 IN

SAMPLE

7. EXPERIMENT 12- COMPLEXOMETRIC TITRATIONS: DETERMINATION

OF HARDNESS OF WATER (Ca2+ EDTA TITRATION)

1
 EXPERIMENT 1

CALIBRATION OF A PIPETTE

1.1 Aim: To determine the accuracy and precision of pipette

1.2. Principle: An ordinary laboratory pipette may be expected to deliver its nominal volume with good
precision and good accuracy if it is used in the way recommended. In this experiment we investigate the
precision and accuracy of such a pipette by making accurate determinations of the mass of water it
delivers in repeated operations.
1.3 Materials: 25.0-cm3 pipette, 50-cm3 beaker, 250-cm3 conical flask, thermometer, pipette filler.

1.4 Procedure:
1. Clean the beaker and the pipette and dry them.
2. Obtained distilled water in the conical flask and let it stand for 15minute at room temperature before
determining its temperature.
3. Weigh the beaker on a balance to determine its mass to the nearest tenth.
4. Fill and discharge the sample pipette into the beaker, and determine the mass of water discharged by
taking the difference.
5. Determine also the temperature of the water pipetted into the beaker.
6. Repeat step 4 until you have the results of four (4) such trials.
Note: You do not need to empty and dry the beaker between trials.
7. Take the mean of the temperatures you have measured as the effective temperature of the water during the
calibration.
1.5 CALCULATIONS/ACTIVITIES
a. Correct the mean mass for air buoyancy by adding to the mean mass, 1.06 mg per gram of water
discharged by the pipette.
b. Determine the density of water, at the temperature measured in 5 above, from the tables given.

c. Find the mean volume of the four trials from the mean mass and the density of water at the temperature
determined in 5 above.
d. Comment on the accuracy of the pipette

2
 Fig. 1: density of water versus Temp
DETERMINATION OF PRECISION
We neto determine the precision of the pipette. You will be given four data points add this to your data
values and perform the statistical analysis to find the standard deviation of the eight data points.

Table 1: Mass and Volume of Water obtained in 8 Pipette Trials

Trial No. Mass of water Vol. of water, Vi (Vi – Vavg.) =di (Vi – Vaveg.)2 = di2
( ) ( ) ( ) ( )
1.
2.
3.
4.
5.
6.
7.
8.
Average
The standard deviation, σ is given by: σ = √ ¿ Ʃ ( di )2 /( N−1)}¿

g. Write your result for the volume you determined in 4 with the error you have now estimated from the
data; Volume of pipette = VO (± s)

7. Comment now, on the accuracy and precision of the pipette.

EXPERIMENT 2

3
 SEPARATION OF MIXTURE (P/Thin-layer Chromatography)
2.1 Aim: To separate mixture of compound into their components
2.2 Principles: Separation of mixtures involve movement of a solvent across a flat surface; components
migrate at different rates due to differences in solubility, adsorption, size or charge; elution is halted when
or before the solvent front reaches the opposite side of the surface and the components examined in situ
or removed for further analysis. Basically, the (R f) is the retention factor is the distribution ration =
Distance moved by solvent
Distance moved by component
Rf is dependent on relative solubility for partition systems or relative polarities for adsorption systems.
2.3 Apparatus and Instrumentation: Thin layers of powdered cellulose (paper), silica gel, alumina, ion-
exchange or gel,

2.4 Procedure:
1. Draw a thin line 2cm from the edge of chromatographic paper or thin-plate as the case may be;
2. Spot the stock solution provided by the technologist using the supplied capillary tube;
3. Place the spotted plate or paper in the mobile phase carefully without allowing the spot getting immersed
into the solvent;
4. Allow the mobile phase to travel along the plate or paper by capillary action carrying the components of
the sample across the plate’s surface at different rates into individual spots or bands. This procedure is
called development;
5. Observe the different colour bands on the developed plate if the sample is coloured;
6. Measure the Rf of the spots or bands on the plate; and
7. Compare the Rf with the references given and predict the components of the sample.

EXPERIMENT 3
Preparation of 500ml standard (0.1M HCl and 0.1M NaOH) Solutions
3.1 Aim: To prepare 500ml of standard solutions (0.1mol/dm3) of HCl and NaOH
3.2 Materials: Standard flask, stock acid, distilled water, measuring cylinder
3.3 Procedure: A
1. Measure the calculated (required) amount of stock HCl by graduated measuring cylinder and
quantitatively pour into the clean 500 ml capacity standard flask;
2. Dilute the acid until 500 ml and finally mixed.
3. Label the prepare solution indicating the date of preparation and concentration of the acid;
4
3.4 CALCULATONS FOR A

From the stock solution

HCl stock solution, 37% purity, d=1.19 g /cm3
m=d xV
The weight of concentrated HCl of 1000 mL is:
M =1.19 x 1000=1190 gr .
i .e ∈100 g=37 g HClis pure ,therefore ,
In 1190 g = x HCl is pure (value of x is obtained by cross multiplication)
x = 440.3 g pure HCl
1 M ∈1 L HCl 36.5 g HCl required
0.1 M ∈1 L will 3.65 g HCl requires . Similarly ,
1000 mL=440.3 g pure HCl
xml=3.65 g HCl (By cross multiplication the value of x is obtained ).
x=8.3 ml
Alternatively
To calculate the amount of conc. Acid needed to prepare 1M of 1000ml
molecular weight X 100
Amount (ml) =
%purity X density
For HCl
36.5 X 100
= = 82.9ml
37 X 1.19
Therefore to prepare 0.1M HCl
If 1M HCl requires 82.9ml
0.1M HCl requires x
82.9 X 0.1
= ~ 8.3ml
1
Therefore, if you take 8.3 ml acid and dilute it to 1000 ml it will be 1L, 0.1 M HCl.
Important note: never add acid to water*(put some water to volumetric flask first, and then add 8.3 mL of
HCl. And mix it well; finally complete it to 1000 mL till the line)
*If you add water onto acid, then high amount of heat is produced and it may explode!

Procedure: B:
1. Weigh a calculated quantity of NaOH pellets into the beaker using analytical balance.
2. Dissolve the pellets in 50-60ml distilled water
5
3. Quantitatively transfer the solution into 500 ml standard flask, mix with some distilled water and make to
mark.
4. Label the prepare solution indicating the date of preparation and concentration of the acid;

The required amount of NaOH pellet to be measure can be estimated form the molar mass of the compound
or use the relation
To calculate the amount of NaOH pellet needed
mass required X 1000
Molarity (M) =
molar mass X Amonut ∈volume
1.0M NaOH solution required 40g NaOH per 1L
1.0M NaOH solution will require 20g NaOH in 0.5 Liter.
0.1M NaOH will require 2g NaOH in 0.5Liter
Be careful! NaOH is a strong irritant. In case of contact, please wash the affected area with large amount of
water.

NEUTRALIZATION TITRATIONS
The reaction between an acid and a base is called as neutralization reaction. Titration is a laboratory
technique that measures the concentration of an analyte using reaction between analyte and standard
solution (solution of known concentration).
Acid-base titration is also called neutralization titrations. Acidimetry is the determination of concentration
of basic substances by titration with a standard acid solution, and alkalimetry is the measurement of
concentration of acid substances by titration with a standard base solution. The end-point (equivalence
point) of acid-base reactions are observed by using indicators (colour indicators) which are
substances that changes colours near their pKa. Therefore, a suitable indicator should be selected for acids
and bases that are reacted.
A titration curve is a plot of pH vs. the amount of titrant added. Shape of titration curves differ for weak and
strong acid-bases or for polyprotic acids and bases. (Post-Lab activity: Sketch the titration curves of:
(i) Weak acid vs strong base, (ii) Polyprotic acids and bases)

EXPERIMENT 4
6
 STANDARDIZATION: ESTIMATION OF WORKING HCl SOLUTION CONCENTRATION
4.1 Aim: To Standardize to working HCl solution using borax solution (primary standard)
4.2 Theory: First borax NaB4O7 salt of weak acid and strong alkali hydrolyses and then reacts with Acid.
4.3. Materials: beakers, burette, pipette, tripod stand, white tile, boric acid solution, working solution and
indicator solution.
4.4 Procedure:
1. Wash the burette with small volume of boric acid solution and fill it with boric acid solution of known
concentration.
2. Wash the pipette with l0 ml of investigative solution (in this case HCl).
3. Pipette 25ml of the working HCl solution in to three Erlenmeyer flasks.
4. Add 2-3 drops of methyl orange to each flask.
5. Note the initial burette reading before starting the titration.
6. Put Erlenmeyer flask with working solution on a white paper to observe the changes in color.
7. Titrate with borax solution until yellow colour appears.
8. Note the burette reading at endpoint and record.
9. Perform the titration in triplicate and record the average titre.
10. Calculate concentration of prepared working HCl solution according to mathematical expression of
equivalent point.

4.5 POST‐LAB QUESTIONS

1. How would it affect your results if you used a beaker with residual water in it to measure out your
standardized HCl solution?
2. How would it affect your results if you used a wet conical flask instead of a dry one when transferring
your acid solution from the volumetric pipette?
3. How do you tell if you have exceeded the equivalence point in your titration?
4. Write a balance chemical equation for the reaction
4. Vinegar is a solution of acetic acid (CH 3COOH) in water. For quality control purposes, it can be titrated
using sodium hydroxide to assure a specific % composition. If 25.00 mL of acetic acid is titrated with
9.08 mL of a standardized 2.293 M sodium hydroxide solution, what is the molarity of the vinegar?

EXPERIMENT 5
7
 DETERMINATION OF PURITY OF ASPIRIN SAMPLES
5.1. Aim: To determine the purity (assay) of Aspirin tablets
5.2 Materials: beakers, burette, pipette, tripod stand, white tile, boric acid solution, working solution and
indicator solution.
5.3 Experimental procedure:
1. Carefully weigh 0.2 gram of solid aspirin sample into a conical flask.
2. Dissolve it in 25 ml 60 % (v/v) ethanol. (Since the solubility of aspirin is very low in water, ethanol
solution is used for its preparation.)
3. Add 1-2 drops of phenolphthalein and titrate with standardized NaOH until the pink color is observed.

Result
1st Titration 2nd Titration 3rd Titration
Initial burette reading (ml)
Final burette reading (ml)
Titre value (ml)
Average titre value (ml)

Calculation:
First, the moles of NaOH consumed during titration can be calculated from the following equation
using the volume of NaOH that is consumed in the titration and the molarity of the NaOH.
According to reaction equation:
If 1 mol NaOH reacts with 1 mol Aspirin
nNaOH will react with x mol Aspirin
Percent amount of the pure aspirin will be calculated. (MW aspirin = 180 g/mol)
The amount of pure aspirin is calculated as the % of the sample weighed.
m aspirin
% purity = x 100
weighed aspirin sample

EXPERIMENT 6
REDOX TITRATION: DETERMINATION OF FeSO4 IN SAMPLE
8
6.2 Theory: Redox titration is a titration based on the oxidation-reduction reaction between
analyte and titrant. Fe2+ is oxidized to Fe3+ in acidic medium by MnO4-
6.3 Materials: beakers, burette, pipette, tripod stand, white tile, boric acid solution, working solution and
indicator solution.
6.4 Experimental procedure:
1. Dilute the 20 mL FeSO4 sample in the volumetric flask to 100 mL with distilled water and mix it well.
2. Transfer a portion of 25 mL diluted sample to an conical flask and add 10 mL ½ diluted H2SO4.
3. Add 50-100 mL distilled water and titrate with standardized KMnO4 solution until pink color.
Reaction equation: MnO4-+ 5Fe2+ + 8H+ Mn2+ + 5Fe3+ + 4H2O
1st Titration 2nd Titration 3rd Titration
Initial burette reading (ml)
Final burette reading (ml)
Titre value (ml)
Average titre value (ml)

Calculations:
Calculate the concentration of the iron (II) sulfate of the sample in g/L (MW FeSO4 = 152 g/mol)
Firstly, calculate the moles of reacted KMnO4 using the molarity of KMnO4 and the volume of
KMnO4 used in the titration:
n KMnO 4=M KMnO 4 ×V KMnO 4
According to the reaction:
If 1 mol KMnO4 reacts with 5 mol FeSO4
nKMnO4 mol KMnO4 will react with x mol FeSO4
Calculate the moles of diluted FeSO4 (𝑦 = nFeSO4) is using above proportion. Using x, calculate the
x
molarity of diluted FeSO4: MFeSO4 =
V FeSO 4
Calculate the molarity of sample by multiplying the molarity of diluted sample with dilution
factor:
Msample = MFeSO4 × DF
Finally, convert the molarity of sample to concentration in g/L:
C(g L ⁄ ) = Msample × 152

EXPERIMENT 7

9
 COMPLEXOMETRIC TITRATIONS: DETERMINATION OF HARDNESS OF WATER
(Ca2+EDTA TITRATION)
7.1 THEORY:
Complex is an ensemble of one or more central atom (usually metals) and ligands that surround
the central atom. They are called as mononuclear complex (one central atom) or polynuclear
complex (more than one central atom) depending on the number of central atoms. If a ligand contains
more than one electron donating group, those groups can coordinate with same central atom and
this type of complexes are called chelate complexes.
For example, protoporphyrin Fe+2 is a chelate complex. This complex is also called «heme» and it
forms hemoglobin with globin.

EDTA coordinate with +2 and +3 charged cations with a 1:1 mol ratio. EDTA forms
hexadentate complexes using its six unpaired electrons (two on nitrogens and four on
carboxyls). EDTA is commercially sold as water-soluble sodium salt (Na 2H2Y.2H2O). Regardless of
the charge of central atoms, 2H+ is produced in EDTA coordination:
7.2 Materials: beakers, burette, pipette, tripod stand, white tile, boric acid solution, working
solution and indicator solution.
7.3 Experimental procedure:
1. Dilute the sample to 100 mL with distilled water and transfer 20 mL of diluted sample to
conical flask.
2. Add 10 mL of pH 10 buffer in conical flask followed by 50 mL distilled water.
3. Add 2 drops of Eriochrome black T indicator. Eriochrome black T forms a red colored complex
with free calcium in the conical flask.
4. Titrate with EDTA until blue color occurs. Blue color formation shows that all calcium forms
complex with EDTA and free eriochrome black T gives its blue color to the solution.
5.
1st Titration 2nd Titration 3rd Titration
Initial burette reading (ml)
Final burette reading (ml)
Titre value (ml)
Average titre value (ml)

10
Calculations:
Calculate the concentration of the Ca2+of the sample in g/L. (MWCa2+ = 40 g/mol)
Firstly, calculate the moles of reacted EDTA using the molarity of EDTA and the volume of
EDTA used in the titration:
nEDTA = MEDTA× VEDTA
According to the reaction:
If 1 mol EDTA reacts with 1 mol Ca2+
nEDTAmol EDTA will react with x mol Ca2+
Calculate the moles of diluted Ca2+ (x = nCa2+) is using above proportion. Using x, calculate
the molarity of diluted Ca2+:
x
MCa2+ =
V Ca 2+¿ ¿
Calculate the molarity of sample by multiplying the molarity of diluted sample with dilution
factor: M Sample=MCa+2 × DF
Finally, convert the molarity of sample to concentration in g/L:
C (g L ⁄ )=M Sample × 40

11