Faculty of science / Chemistry department
Practical
ANALYTICAL
CHEMISTRY II
Chem. 390 manual
Prepared by:
Mrs. Danah Al Shaer
Mrs. Tasneem Iqbal
Supervised by:
Dr. Amaal Mansour Abdel-Satar
1
Experiment (1)
Redox Titration
Standardization of KMnO4 against Oxalic acid
1.0 OBJECTIVE:
To prepare and standardize a solution of an oxidizing agent, by titrating it with standard
solution of a reducing agent.
2.0 INTRODUCTION
A. Standardization process
A standard solution is one prepared by weighing a substance that can be obtained in very
pure form, and that does not change chemically when stored for a reasonable period of time.
In this experiment a solution of a primary standard, oxalic acid (written as; (COOH)2)
will be used as a reducing agent to determine the molarity of a solution of an oxidizing agent that
is KMnO4, the process of titration of the standard of known concentration against a solution of
unknown concentration is called standardization.
B. Oxidation-Reduction (redox) reactions
Reactions in which electrons are transferred from one species to another are called
oxidation-reduction (redox) reactions. When one species loses electrons by an oxidation
process another species simultaneously gains electrons by a reduction process in a chemical
reaction.
In this experiment we will consider the following redox equation:
2KMnO4 + 5(COOH)2 + 3H2SO4
10CO2 + 2MnSO4 + K2SO4 + 8H2O
We can separate this equation into two half cell reactions.
2MnO4- + 16H+ +10e+5(COOH)2
2Mn2+ + 8H2O
10 CO2 + 10 H+ + 10e-
(half cell reaction/reduction)
(half cell reaction/oxidation)
2
Note that; in the half cell equations, that the number of electrons gained (10) is equal to the
number of electrons lost (10). You will use the stiochiometry of the balanced equation to
calculate the molarity of the solution. That is;
2 KMnO4: 5 (COOH)2
CAUTION: YOU MUST WEAR DEPARTMENTALLY APPROVED EYE PROTECTION
AT ALL TIMES YOU ARE IN THE LABORATORY!! KEEP ALL REACTION VESSELS
WELL AWAY FROM YOUR FACE!!
3.0 PROCEDURE:
1. Prepare 50.0 ml of (0.02 M) KMnO4 solution (by dissolving 0.158g of KMnO4 in 50.0 ml
of distilled water).
2. Be sure that your burette is clean and the stopcock is closed. Pour a few milliliters of the
potassium permanganate solution into the burette. Slowly rotate the burette so that the
solution comes in contact with the walls of the burette. Open the stopcock and drain the
potassium permanganate solution into a waste beaker at your station.
3. Repeat step 3 two more times.
4. Close the stopcock and fill the burette with the potassium permanganate solution. Record
the initial volume of the potassium permanganate solution in the burette.
5. Take 10.0 ml of the previously prepared (0.05 M) Oxalic acid in titration flask.
6. Add 10.0 ml of (0.1 M) H2SO4 to Oxalic acid in titration flask.
7. Heat the mixture to 60-70 ◦C.
8. Titrate the solution in the flask against KMnO4 solution in burette.
9. Record the volume of KMnO4 consumed till purple-pink color appears.
3
Experiment (1) Report
Redox Titration
Standardization of KMnO4 against Oxalic acid
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet:
-Volume of Oxalic acid used =
-Concentration of Oxalic acid =
- Volumes of KMnO4:
Trial (1)
Trial (2)
Trial (3)
Initial reading
Final reading
Volume used
Average volume of KMnO4 consumed =
Calculations:
-
Calculate the molar concentration (molarity) of the standardized KMnO4 solution (Hint:
use the following equations)
2KMnO4 + 3H2SO4 + 5H2C2O4
K2SO4 +2MnSO4 + 8H2O + 10CO2
=
4
-
Suppose that the KMnO4 solution you’ve prepared is primary standard; what should be
its molar concentration? (use the following equations)
Molarity =
-
, number of moles=
Do both concentrations differ from each others? If so state why?
5
Experiment (II)
Redox Titration
Determination of Iron by Reaction with Permanganate
1.0 OBJECTIVE:
To determine the percentage of iron in iron(II) ammonium sulfate solid salt, Fe(NH4)2(SO4)2
• 6 H2O, as well as to determine Iron (II) concentration in an Unknown iron-containing solution
sample.
2.0 INTRODUCTION
Potassium permanganate, KMnO4, is widely used as an oxidizing agent in volumetric
analysis. In acid solution, MnO4- ion undergoes reduction to Mn2+ as shown in the following
equation:
8 H+(aq) + MnO4-(aq) + 5 e-  Mn+2(aq) + 4 H2O(l)
Since the MnO4- ion is violet and the Mn+2 ion is nearly colorless, the end point in
titrations, using KMnO4 as the titrant, can be taken as the first permanent pink color that appears
in the solution (that’s why; KMnO4 is considered as self-indicator).
KMnO4 will be employed in this experiment to determine iron(II) in given samples. The
oxidation reaction of the iron(II) ion is as follows:
Fe2+(aq)  Fe3+(aq) + eThe titration, which involves the oxidation of Fe2+ ion to Fe3+ by permanganate ion, is
carried out in sulfuric acid solution to prevent the air-oxidation of Fe2+. The end point of the
titration is sharpened markedly if phosphoric acid is present, because the Fe3+ ion produced in
the titration forms an essentially colorless complex with the acid.
The number of moles of potassium permanganate used in the titration is equal to the
product of the molarity of the KMnO4 and the volume used. The number of moles of iron present
in the sample is obtained from the balanced equation for the reaction and the amount of MnO4ion reacted. The percentage by weight of iron in the solid sample follows directly.
CAUSION:


The sulfuric acid, H2SO4, and the phosphoric acid, H3PO4, are both strong acids that are
corrosive. Laboratory coat and splash goggles are required. In the event that either of
these acids gets on your skin or clothing, rinse with copious amounts of water and notify
the teacher immediately.
If either sulfuric acid, H2SO4, or phosphoric acid, H3PO4, gets on your face or in your
eyes, immediately begin flushing the face/eyes with water and notify the teacher.
6


If either acid spills on the floor or lab counter, sprinkle acid-neutralizer on the spill.
When the acid has been absorbed, clean up the residue. Notify the teacher.
The potassium permanganate solution stains both clothing (permanently) skin (takes
several days to wear off). Use caution to avoid contact with the solution.
3.0 PROCEDURE:
Part A: Determination of Iron(II) percentage in Fe(NH4)2(SO4)2 • 6 H2O salt:
1. Mass on the analytical balance three samples of about 0.2 gram apiece of the standard,
iron(II) ammonium sulfate hexahydrate, Fe(NH4)2(SO4)2 • 6 H2O. Record the mass to
0.0001 grams on the data table.
2. Obtain about 100 mL of the previously standardized potassium permanganate solution
(about 0.02M), in a 150-mL beaker.
3. Close the stopcock and fill the burette with the potassium permanganate solution. Record
the initial volume of the potassium permanganate solution in the burette.
4. Add 50-mL of the 1 M sulfuric acid, H2SO4, to one of the samples of iron(II) ammonium
sulfate hexahydrate, Fe(NH4)2(SO4)2 • 6 H2O that is in a 250-mL Erlenmeyer flask. (The
sample should dissolve completely.)
5. Without delay, titrate this iron solution with the KMnO4 solution in the burette. When a
light yellow color develops in the iron solution, add 3 mL of the 85% phosphoric acid,
H3PO4.
6. Continue the titration with the KMnO4 until the first pink color shows up and lasts for 1530 seconds of mixing. Record the final volume of the KMnO4 solution in the burette.
7. Repeat the titration (steps 6-8) with the other two samples of the standard, iron(II)
ammonium sulfate hexahydrate, Fe(NH4)2(SO4)2 • 6 H2O.
Part B: Determination of Iron in the Unknown
1. Follow the same procedure as above (steps 5-8 in the Standardization of the Potassium
Permanganate solution), to determine the concentration of iron in the unknown sample
that has been provided to you. For the unknown solution, pipet 10.0 ml for each trial.
7
Experiment (1I)
Redox Titration
Determination of Iron by Reaction with Permanganate
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet:
Part A: Determination of Iron(II) percentage in Fe(NH4)2(SO4)2 • 6 H2O salt:
-Concentration of KMnO4 =
- Volumes of KMnO4 consumed:
Trial (1)
Trial (2)
Trial (3)
Trial (2)
Trial (3)
Mass of salt
Initial reading
Final reading
Volume used
-
Average sample mass=
-
Average KMnO4 volume consumed=
Part B: Determination of Iron in the Unknown
-
Volume of unknown sample taken=
-
Volume of KMnO4 consumed:
Trial (1)
Initial reading
Final reading
Volume used
-
Average volume of KMnO4=
8
Calculations:
-
Write down the balanced equation for the reaction between Fe2+ and MnO4- :
Part A: Determination of Iron(II) percentage in Fe(NH4)2(SO4)2 • 6 H2O salt:
-
Calculate mass percentage of Fe in salt sample.
-
Calculate the theoretical mass percent in Fe(NH4)2(SO4)2 • 6 H2O
Part B: Determination of Iron in the Unknown
-
Calculate the molar concentration of Fe2+ in the unknown sample
-
Calculate the concentration of Fe2+ in (g/liter)
9
Experiment (III)
Redox Titration
Estimation of Iodine using Thiosulphate ion
1.0 OBJECTIVE:
To determine the concentration of Iodine (I2) in an Unknown solution sample, through
titrating sample aliquot with a previously standardized sodium thiosulphate (Na2S2O3) solution.
2.0 INTRODUCTION
Iodometry is one of the most important redox titration methods. In which the appearance or
disappearance of elementary iodine indicates the end point.
This importance is based on one hand on the oxidation effect of the elementary iodine (I2),
and on the other hand on the reduction effect of the iodide ion (I-). This process is completely
reversible:
I2 + 2e-
2I- ……………….. (1)
Basically; there are two possibilities in iodine titrations:
1) Reducing agents could be titrated directly with iodine solution. They are oxidized,
whereas iodine is reduced to iodide. For example:
S2- + I2
S + 2I-
……………….. (2)
2) Oxidizing agents could react with an excess of an acidic iodine solution. They are
reduced, whereas iodide is oxidized to elemental iodine. For example:
2Fe3+ + 2I-
2Fe2+ + I2 ……………….. (3)
10
Then, the librated iodine is titrated with a standardized reducing agent. The mostly used
reducing agent is sodium thiosulphate. The equation for this reaction is:
2S2O32- + I2
S4O62- + 2I- …………….. (4)
This reaction goes to completion only in neutral or slightly alkaline solutions due to the
disproportionatin of I2 in highly alkaline solutions:
I2 + 2OH-
IO- + I- + H2O
……………….. (5)
In highly acidic solution, I- produced from the titration reaction can react with dissolved
O2 to reform I2, resulting in ambiguous endpoints.
4I- + O2 + 4H+
2I2 + 2H2O
……………….. (6)
Standardization of sodium thiosulphate:
Sodium thiosulphate is standardized against iodine librated from a primary
standard agent such as potassium iodate, according to the following equation:
KIO3 + 5KI + 6HCl
3H2O + 3I2 + 6KCl …….. (7)
Iodine librated is reacted with sodium thiosulphate according to equation 4.
End point detection:
Iodine may act as self indicator through the appearance or disappearance of the color of
iodine or triiodide ion (I3-). Mostly, starch is used as an indicator as it gives deep blue color in
the presence of iodine.
11
3.0 PROCEDURE:
Part A: Standardization of 0.1M sodium thiosulphate (S2O32-) solution:
1. Fill your burette with sodium thiosulphate solution.
2. Pipet 10.0 ml of KIO3 solution into 250 ml Erlenmeyer flask.
3. Add about 0.25 g potassium iodide (KI) and swirl to dissolve.
4. Add 5 ml of sulfuric acid (H2SO4).
5. Titrate immediately with sodium thiosulphate solution with continuous stirring until the
color becomes pale yellow.
6. Add 10 ml distilled water and 2 ml of starch solution and continue the titration carefully
until the color changes from blue to colorless.
7. Repeat the titration two more times and calculate the molarity of potassium thiosulphate.
Part B: Determination of Iodine (I2) in an unknown sample:
1. Pipet 10.0 ml of the iodine solution into an Er. Flask and dilute with 10 ml with distilled
water.
2. Titrate with standardized sodium thiosulphate till the color becomes pale yellow.
3. Add 2 ml starch and continue the titration till the blue color disappears.
4. Repeat the titration two more times and calculate the molarity and the number of grams
per liter of Iodine solution.
12
Experiment (III)
Redox Titration
Estimation of Iodine using Thiosulphate ion
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Data Sheet:
Section No.: ________________
Part A: Standardization of 0.1M sodium thiosulphate (S2O32-) solution:
- Concentration of KIO3 used=
- Volume of KIO3 used=
-volumes of S2O32- consumed:
Trial (1)
Trial (2)
Trial (3)
Initial reading
Final reading
Volume used
-
Average S2O32- volume consumed=
Part B: Determination of Iodine (I2) in an unknown sample:
- Volume of unknown I2 used=
-volumes of S2O32- consumed:
Trial (1)
Trial (2)
Trial (3)
Initial reading
Final reading
Volume used
13
-
Average S2O32- volume consumed=
-
Concentration of S2O32-used= (from calculations)
Calculations:
-
What is the molar ratio for the reaction between S2O32- and I2?
Part A: Standardization of 0.1M sodium thiosulphate (S2O32-) solution:
-
Calculate the molarity of sodium thiosulphate (S2O32-) solution.
Part B: Determination of Iodine (I2) in an unknown sample:
-
Calculate the concentration of Iodine solution in
- (mole / liter):
-
(g/liter):
-
Why we didn’t add the starch indicator before beginning the titration process?
14
Experiment (IV)
Redox Titration
Iodometric Determination of Ascorbic Acid (Vitamin C)
1.0 OBJECTIVE:
To determine the percentage of iron in iron(II) ammonium sulfate solid salt, Fe(NH4)2(SO4)2
• 6 H2O, as well as to determine Iron (II) concentration in an Unknown iron-containing solution
sample.
2.0 INTRODUCTION
Vitamin C (or ascorbic acid) is an important antioxidant. In the cells it is easily oxidized to
dehydroascorbic acid, removing oxidizing agents before they can do damage to other substances
present. This reaction is the basis of the iodometric titration of ascorbic acid - it is quantitatively
oxidized by iodine.
Beside; is the chemical structure
of ascorbic acid (Vitamin C).
A suitable method for the determination of vitamin C (C6H8O6) quantities is a titration with
potassium iodate (KIO3). Potassium iodate is used as a titrant and it is added to an ascorbic acid
solution that contains strong acid and potassium iodide (KI). Potassium iodate reacts with
potassium iodide, liberating molecular iodine (I2):
KIO3 + 5KI + 6H+ → 3I2 + 6K+ + 3H2O ……………………(1)
As long as the solution contains ascorbic acid, the iodine (I2) produced in (1) is used up in a
rapid reaction with ascorbic acid (C6H8O), during which dehydroascorbic acid (C6H6O6) and
iodide ion are formed:
C6H8O6 + I2 → C6H6O6 + 2I- + 2H+ …………………………(2)
Potassium iodide must be added in excess to keep iodine dissolved.
15
End point detection:
Once all the ascorbic acid has been consumed, any excess iodine will remain in solution.
Since aqueous iodine solutions are brown in colour, iodine can act as its own indicator.
However, it is quite difficult to detect endpoints using iodine colouration alone, and it is more
usual to add starch, which forms an intensely blue coloured complex with iodine but not with the
iodide ion.
3.0 PROCEDURE:
1. Fill your burettete with 0.05 M potassium iodate (KIO3) solution; record the intial
reading of the burette.
2. Use a mortar and pestle to crush a vitamin C tablet.
3. Weigh accurately about 0.5 g of ground vitamin C powder (two portions) , and record
their masses
4. Transfer each portion to a clean, 250 ml Er. Flask.
5. Add 75 ml of distilled water to flask 1 and swirl to dissolve most of the powder (do
not bother if all of the powder does not dissolve. Tablets contain various binders and
fillers, which may not dissolve in water).
6. Add about 0.25g potassium iodide and swirl to dissolve.
7. Add 5 ml of 2 M hydrochloric acid (HCl).
8. Add 2 ml of starch indicator.
9. Titrate with KIO3 solution till black-blue color appears; record the final reading of the
burette.
10. Repeat with flask 2 steps 5-9.
16
Experiment (IV)
Redox Titration
Iodometric Determination of Ascorbic Acid (Vitamin C)
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet:
- Concentration of KIO3 used=
- Volumes of KIO3 consumed:
Trial (1)
Trial (2)
Trial (3)
Mass of powder
Initial reading
Final reading
Volume used
-
Average KIO3 volume consumed=
-
Average mass of tablet sample=
-
Reported mass percent of ascorbic acid in vitamin C tablet:
Calculations:
-
What is the molar ratio between the reacted KIO3 and Vitamin C in the sample?
-
Calculate mass of ascorbic acid per gram tablet.
-
What is the mass percent of ascorbic acid in vitamin C tablets?
17
GRAVIMETRIC ANALYSIS
INTRODUCTION:
Gravimetric analysis is one of the most accurate and precise methods of macro
quantitative analysis. In gravimetric analysis the measurement signal is mass. The analyte is
selectively converted to an insoluble form, which is separated, washed, dried or ignited then
weighed accurately. From the weight of the precipitate and knowledge of its chemical
composition, the concentration of the analyte in the sample can be calculated.
Gravimetric analysis normally involves processes of:
1. Preparation of the analyte solution.
2. Quantitative precipitation: it is necessary to ensure complete precipitation.
3. Digestion: the precipitate must be digested for purposes of coagulation or increase in the
particle size.
4. Filtration: quantitative transfer of the precipitate using proper techniques as will be
illustrated later.
5. Washing: this is as important step of removal of other co-precipitated species. Proper
choice of washing solvent is necessary for this purpose.
6. Drying or ignition: the precipitate must be stable at the drying temperature. Drying time
should ensure complete removal of solvent.
7. Weighing: the washed and dried precipitate is then weighed to “constant weigh” on an
analytical balance.
Equipments used in Gravimetric Analysis:
a. Desiccators:
Desiccator is a vessel, usually of glass which is used to equilibrate objects within a
controlled atmosphere.
b. funnels and Filter paper:
In gravimetric procedures the desired constituent is often separated in the form of
precipitate. This precipitate must be collected, washed free of undesirable contaminants from
the mother liquor, dried and weighed. Filtration is carried out with either funnels and filter
papers or filtering crucibles.
18
Experiment (V)
Gravimetric Analysis
Gravimetric determination of Sulphate ion
1.0 OBJECTIVE:
To determine the concentration of Sulphate ion in an unknown solution using gravimetric
techniques, through precipitating it as BaSO4.
2.0 INTRODUCTION
The analysis of a soluble sulphate is based upon precipitation with Barium ion:
Ba2+ (aq) + SO42-(aq)
BaSO4(s)
The crystalline barium sulphate precipitate is collected on an ashless filter paper, washed
with water and strongly ignited to constant weight.
The sulphate content is calculated from the weight of BaSO4 using the gravimetric factor:
Mass of SO42- = mass of BaSO4 X
PROCEDURE:
A. Precipitating Sulphate ion from unknown solution
1.
2.
3.
4.
Heat about 30 ml of barium chloride solution BaCl2.
Pipet 10.0 ml of the unknown sulphate solution into 400 ml beaker.
Add 4ml of 6M HCl and dilute to about 200 ml with distilled water.
Heat the solution to boiling, and quickly with vigorous stirring add the hot barium
chloride solution.
5. Digest the precipitated BaSO4 at just below the boiling point for half an hour on a small
flame.
19
B. Preparation of the crucible
6. Clean and mark a porcelain or silica crucible and cover.
7. Ignite them for 15 minutes in the muffle furnace at 900ᵒC.
8. Handle the crucible only with tongs.
9. Cool in a desiccator then weigh.
10. Repeat the ignition until a constant weight is achieved (±0.3 mg).
C. Filtration and Washing of the Precipitate
11. Decant the hot supernatant through a fine ashless filter paper.
12. Test the filtrate with few drops of barium chloride to ensure complete precipitation.
13. Quantitatively transfer the precipitate to the filter paper.
14. Wash the precipitate until it is chloride free (this can be tested by adding a drop of silver
nitrate solution to a test portion of the washing collected in a test tube, absence of
turbidity indicates chloride free precipitate).
D. Ignition
15. Fold the moist paper around the precipitate and place it in the previously prepared
crucible.
16. Dry the paper by placing the loosely covered crucible upon a triangle above a small
flame.
17. Gradually increase the heat until the paper chars and volatile matter is expelled (do not
allow the paper to burst into flame)
18. When charring is complete, raise the temperature of the crucible to dull redness and burn
off the carbon (crucible is slightly inclined with cover displaced)
19. When the precipitate is white, ignite the crucible for 10-15 minutes in muffle furnace at
900ᵒC.
20. Allow the crucible to cool somewhat in the air, transfer to a desiccator and when cold
weigh the contents.
20
Experiment (V) Report
Gravimetric Analysis
Gravimetric determination of Sulphate ion
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet and Calculations:
Volume of unknown solution
Mass of empty crucible
Mass of crucible + precipitate
Mass of precipitate (BaSO4)
Mass of Sulphate
Concentration of Sulphate in (mg/liter)
Questions:
-
Why should the sulphate precipitation be carried out in slightly acidic solution?
-
Why the BaCl2 solution is added quickly to the sulphate containing solution?
-
Why filtration is carried out while the solution is still hot?
-
Why the filter paper should be charred without inflaming?
21
Experiment (VI)
Gravimetric Analysis
Gravimetric determination of Nickel as Bis(dimethylgloximato) nickel(II)
Ni(DMG)2
1.0 OBJECTIVE:
To determine the concentration of Nickel (II) ion in an unknown solution using gravimetric
techniques, through precipitating it as Ni(DMG)2.
2.0 INTRODUCTION
In the gravimetric determination of a substance; the precipitating agent should be as
specific as possible for that substance. Dimethylglyoxime (DMG) has a specificity to
precipitate only palladium from acidic medium and only nickel from weekly basic solutions.
In ammonia solution, nickel is precipitated quantitatively by DMG as a brigh strawberry-color
complex, Ni(DMG)2.
The specificity of DMG arises from the unusual nature of the nickel chelate. Two DMG
ligands are coordinated into each nickel ion, with the nickel and all four coordinating nitrogen
atoms in the same plane:
Ni(DMG)2 is bulky when first precipitated; the precipitate also tends to creep up the walls
of the vessel containing it. For these reasons only small quantities of nickel can be handled.
This method is used for gravimetric determination of nickel in steel.
22
PROCEDURE:
1. Pipet 10.0 ml of the unknown nickel salt solution into a 400 ml beaker.
2. Add about 5 ml of dilute HCl then dilute to about 200 ml with distilled water.
3. Heat to 70-80 ᵒC (use a thermometer).
4. Reduce the flame and add 30 ml of DMG.
5. With good stirring add dilute ammonia dropwise, until precipitation takes place and then
add a slight excess.
6. Heat very gently on a small flame and test the solution for complete precipitation when
the red precipitate has settled out.
7. Allow the precipitate to stand for half an hour at room temperature.
8. Clean up a sintered glass crucible and heat it to constant weight at 110-120 ᵒC.
9. Filter the cold solution through the crucible, wash the precipitate with warm water and
then with 30% alcohol, to dissolve any excess DMG from the precipitate.
10. Finally, put the crucible in an oven at 110-120 ᵒC for 30 minutes.
11. Allow to cool in a desiccator and weigh. Repeat drying until constant weight is achieved.
Record your results.
23
Experiment (VI)
Gravimetric Analysis
Gravimetric determination of Nickel as Bis(dimethylgloximato) nickel(II)
Ni(DMG)2
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet and Calculations:
Volume of unknown solution
Mass of empty crucible
Mass of crucible + Ni(DMG)2
Mass of precipitate Ni(DMG)2
Mass of Ni
Concentration of Ni(II) in (mg/liter)
Hint: use the following equation in calculating mass of Ni(II):
Mass of Ni2+ = mass of Ni(DMG)2 X
Questions:
-
Why precipitation is carried out in ammoniacal solution?
-
Why a large excess of DMG should not be added?
-
Why 30% of ethanol is used for washing?
24
Experiment (VII)
Gravimetric Analysis
Gravimetric determination of water of hydration
3.0 OBJECTIVE:
The purpose of this experiment is to determine the empirical formula of a hydrate using
gravimetric techniques.
4.0 INTRODUCTION
Many solid ionic compounds contain weakly bound water molecules within their crystal
structures. Such solids are called hydrates and the bound water molecules are called water of
hydration.
Water of hydration is included in the name and the chemical formula for the hydrate,
connecting the water molecules to the rest of the formula with a raised dot. For example the
chemical formula for copper(II) sulphate pentahydrate is CuSO4. 5H2O.
Water of hydration can be removed from the hydrate using a process called dehydration,
leaving the anhydrous (“without water”) form of the compound. For example, when we heat
blue CuSO4. 5H2O, the water of hydraion is released as water vapor, and solid, white anhydrous
CuSO4 remains, as shown in Equation 1.
CuSO4. 5H2O (blue solid)
heat
CuSO4 (white solid) + 5 H2O (g)
Other examples of hydrates are; MgSO4.7H2O: magnesium sulphate heptahydrate and
Alum. Where alum is both a specific chemical compound and a class of chemical compounds.
The specific compound is the hydrated potassium aluminiumsulphate (Potassium alum) with
the formula KAl(SO4)2.12H2O. The wider class of compounds known as alums have the related
empirical formula, AB(SO4)2.12H2O.
25
PROCEDURE:
1. Clean two crucibles and lids using soap, water and test tube brush. Rinse with water and dry
with a paper towel.
2. Set up a ring stand with 2 rings and place a Bunsen burner under each ring. Adjust the height
of the rings to about 2 inches above the top of the Bunsen burners. Place a clay triangle on each
of the rings. Then carefully light both Bunsen burners.
3. Using crucible tongs, place the crucibles with their lids on the clay triangles and heat intensely
for 10 minutes to dry the crucibles.
4. Allow the crucibles to cool for one minute and, using crucible tongs, place them in a
desiccator for 10 minutes or until they reach room temperature.
5. Keep crucible and lid together as a matched pair throughout the entire experiment. Each
matched pair will have its own column of data on the data page.
6. Using crucible tongs, transfer a matched pair to the balance. It is important to use the same
balance throughout the experiment. Weigh and record the mass to 3 decimal places on the data
page. NOTE: Using crucible tongs prevents mass inaccuracies due to contamination of the
crucibles and lids with oil and moisture from the skin.
7. Place slightly more than one gram of the sample in each crucible. Weigh and record the mass
to 3 decimal places on the data page. Place the matched pair in the desiccator using crucible
tongs.
8. Repeat steps 5 and 6 for the other matched pair.
9. Place each matched pair on a clay triangle. Heat by a Bunsen burner using a low flame for five
minutes. NOTE: The low flame prevents spattering of the sample during dehydration. If
spattering occurs, mass may be lost when the crucible and lid are transferred from the ring to the
desiccator.
10. Gradually increase the temperature of the flame until you are using a medium flame and
continue heating for an additional ten minutes. CAUTION: DO NOT ALLOW CRUCIBLES OR
SAMPLES TO BECOME RED HOT SINCE UNDESIRED SAMPLE DECOMPOSITION
MAY OCCUR.
11. Allow the crucibles to cool 1 minute and then place them in a desiccator for 10 minutes or
until they reach room temperature.
12. Weigh each matched pair and record the mass to 3 decimal places on the data page.
13. Heat again for 10 minutes on medium heat. Cool in desiccator, weigh to 3 decimal places and
record on the data page. If the mass is not within 0.005 g of the previous mass, repeat this step
until the mass is within 0.005 g.
14. Record your results on the "master data sheet" provided by your instructor.
26
15. Residue in the crucibles should be placed in the labeled waste container.
NOTE: Always perform weighing on the same balance so that constant errors will cancel out.
Always use crucible tongs to handle your crucibles.
Figure 1: Heating the crucible
27
Experiment (VII) Report
Gravimetric Analysis
Gravimetric determination of water of hydration
Name:____________________
ID No.: ___________________
Instructor Name: ______________
Section No.: ________________
Data Sheet and Calculations:
Hydrate: ____________.XH2O
Mass of crucible
Mass of crucible + hydrate
Mass of crucible + anhydrous
Observation:
Calculation
1- Mass of hydrate:
2- Mass of anhydrous:
3- Mass of water lost:
4- Number of moles of anhydrous salt:
5- Number of moles of water:
6- The value of “X” in the formula:
(X = number of moles of water /number of moles of anhydrous salt)
7- Percent mass of water of hydration:
28
Questions:
1- Calculate the percent mass of water of hydration in BaCl2.2H2O.
2- What is the effect of incomplete dehydration on the calculated value of “X”?
29
Experiment (VIII)
Potentiometric Titration
Acid – Base reaction
1.0 OBJECTIVE:
In this experiment, you will use a pH meter to follow the course of acid-base titrations.
From the resulting titration curves, you will determine the concentrations of the acidic solutions
as well as the acid-ionization constant of a weak acid.
2.0 INTRODUCTION
You have performed acid-base titrations in the past to determine the concentration of an
acidic or basic solution using a colored indicator. However, there are times when an appropriate
indicator does not exist, or where the color of the solution would obscure any color change
associated with the endpoint. In such cases, a pH meter can be used to monitor the acidity of the
solution throughout the titration. Recall the definition of pH:
pH = –log[H3O+]
The pH meter:
A pH meter consists of two electrodes: a glass electrode, which is sensitive to the
concentration of hydronium ions in solution, and a reference electrode. The reference electrode is
often a calomel electrode, which supplies a constant potential (E° = +0.24 V versus the standard
hydrogen electrode) as determined by the half-reaction
Hg2Cl2 + 2 e–
2 Hg + 2 Cl– ………….. eq(1)
Calomel is the trivial name for the compound Hg2Cl2. When both the reference and glass
electrodes are contained in a single unit, it is referred to as a combination electrode.
30
The potential of the glass electrode is proportional to the logarithm of the ratio of [H3O+] inside
and outside the electrode. The pH meter measures the total potential across the two electrodes
and displays this measurement on a scale calibrated in pH units. The pH meter is an accurate and
easy-to-use device for determining the pH of a solution.
Potentiometric titrations:
Figure 1 shows a plot of pH versus volume of base added for the titration of a strong
acid with a strong base. There is very little change in pH when the base is initially added. Below
the equivalence point, the pH is a function of the amount of excess acid present. Above the
equivalence point, the pH is a function of the amount of excess base present. The equivalence
point for the titration of a strong acid with a strong base occurs when [OH–] exactly equals
[H3O+] in the solution; pH = 7.0.
Figure 1. Titration curve for the titration of a
strong acid with a strong base.
31
The situation in the case of the titration of a weak acid with a strong base is somewhat
different due to the fact that a weak acid is only partially ionized in aqueous solution. A dynamic
equilibrium exists, which is represented by the following equation:
H3O+ + A– ……………………eq(2)
HA + H2O
The equilibrium expression for this reaction is:
Ka =
………………………Eq. 1
Or
[H3O+] = Ka (
)
Where, Ka is the acid-ionization constant for the weak acid. Let us assume that the initial
dissociation of the weak acid is negligible. The progressive addition of NaOH during the titration
decreases the concentration of HA and increases the concentration of its salt,
NaA:
HA (aq) + NaOH (aq)
H2O(l) + NaA (aq)
The presence of both HA and its salt, NaA, creates a buffer system, which resists a large
change in pH. The ratio of [HA]/[A–] changes only slightly; therefore, according to Eq. 1, the
change in [H3O+] (or pH) must also be small. The pH increases slowly until the equivalence
point is approached (see Figure 2).
32
At the halfway point in the titration, exactly half of the HA originally present will have
been neutralized, and therefore the concentrations of HA and A– will be equal, so [HA]/[A-]=1.
Substituting this information into Eq. 1, we obtain:
Ka = [H3O+]1/2…………………………..Eq. 2
Thus, the ionization constant of a weak acid is equal to the hydronium ion concentration
at the halfway point in the titration; pKa = pH1/2. This relationship is valid only if the initial
dissociation of the acid is negligible. When the degree of dissociation is appreciable as in the
case of a very dilute solution, the pH at the midpoint of the titration bears no relation to the value
of Ka.
33
The subsequent rapid increase in pH and the inflection in the titration curve at the
equivalence point can be accounted for. As the equivalence point is approached, the
concentration of unreacted HA becomes progressively smaller so that successive increments of
NaOH neutralize a greater fraction of the HA remaining. This produces a large change in the
[HA]/[A–] ratio and, therefore, in the pH of the solution.
At the equivalence point, the acid and base have reacted completely to yield the salt,
NaA. The pH at the equivalence point is determined by the strength of the base, A–. The
conjugate base of a weak acid is a strong base. It will react with water to produce hydroxide ions
(hydrolysis):
A– (aq) + H2O (l)
HA (aq) + OH– (aq)
For this reason, it is not surprising to see a pH which is greater than 7 at the equivalence
point.
Beyond the equivalence point, the pH is determined by the ion product for water:
Kw = [H3O+][OH–]
The first small excess of NaOH greatly increases the concentration of OH–,
concomitantly decreasing the H3O+ concentration, and causing the pH to continue to increase.
Well past the equivalence point, the concentration of OH– becomes so large that only slight
changes in pH are produced.
You will titrate a solution of HCl with a standardized solution of NaOH while measuring
the pH throughout the course of the titration. From your titration curve, you will determine the
concentration of the HCl solution.
34
You will also titrate a sample of commercial vinegar using a standard solution of NaOH.
The active ingredient in vinegar is acetic acid, which is a weak acid.
CH3COO– (aq) + H3O+(aq)
CH3COOH (aq) + H2O (l)
The acid-ionization constant of acetic acid is:
Ka =
= 1.74 x 10-5
From your titration curve, you will be able to determine the concentration of acetic acid in
commercial vinegar.
35
PROCEDURE:
Part I. Titration of a Strong Acid
1. Into a clean and dry 150 mL beaker, and with a carefully rinsed volumetric pipet,
dispense a 10.00 mL aliquot of the 0.5 M HCl solution (do not pipet directly from the
bottle).
2. Add exactly 75.0 mL of distilled water and 3 drops of phenolphthalein solution.
3. Fill a clean and carefully rinsed burette with the standardized 0.5 M NaOH solution
(record the exact molarity from the label). Record the initial burette reading in your
notebook.
4. Remove the pH electrode from the buffer solution. Thoroughly rinse the electrode with
distilled water, shake off the drops of water and place it in the acid solution such that the
tip is immersed. Stir the acid solution with the pH electrode (be careful not to break the
glass tip!).
5. Begin the titration by adding, with stirring, 1 to 2 mL of NaOH. Be careful not to splash
any liquid out of the beaker. When the pH reading is stable, stop stirring, and then record
the pH of the solution.
6. Continue to add base, record the pH. The pH values should increase in approximately 0.2
pH unit increments. Be certain the solution is stirred after each addition of titrant and that
the pH is stable before recording it. Slow down as you approach the equivalence point
(as indicated by the appearance of the pink color)!
7. As the pH readings approach 2, reduce the amount of base added to 0.1 mL increments.
Record in your notebook the pH reading when the pink phenolphthalein endpoint color
persists for 30 seconds. Add increments consisting of several drops of NaOH beyond this
endpoint; then increase the increments of base to 1-2 mL until 10 mL more of NaOH has
been added. Record your data!
Part II. Titration of vinegar
1. Refill the burette with the standardized 0.5 M NaOH solution provided.
2. Record the initial volume of NaOH in your notebook.
3. Using a carefully cleaned and rinsed transfer pipet, dispense a 10.00 mL aliquot of
vinegar into a clean, dry 150 mL beaker (do not pipet directly from the bottle).
36
4. Dilute this aliquot with exactly 75.0 mL of distilled water.
5. Add 3 drops of phenolphthalein solution.
6. Thoroughly rinse the electrode with distilled water, shake off the drops of water and
place it in the vinegar solution.
7. Stir the vinegar with the pH electrode (be careful not to break the glass tip!), and arrange
the burette above the beaker for a titration.
8. Titrate the vinegar as you did the HCl solution above. As you approach the equivalence
point, decrease the size of the base increments until single drops are being added.
9. Record in your notebook the pH reading when the pink phenolphthalein endpoint color
persists for 30 seconds.
10. Continue the titration until a pH of approximately 12 has been reached. Record your data.
11. All of the solutions in this experiment may be poured down the sink. Clean off your lab
bench before you leave and return the rinsed pH electrode to the plastic bottle.
12. Record the percent acidity of the vinegar from the label on the bottle.
37
Experiment (VIII)
Potentiometric Titration
Acid – Base reaction
Name:____________________
ID No.: ___________________
Instructor Name: ______________
Section No.: ________________
Data Sheet and Calculations:
Part I. Titration of a Strong Acid
1. Prepare a graph of pH versus milliliters of NaOH and locate the equivalence point.
2. Find out the concentration of HCl used.
Part II. Titration of a Weak Acid
1. Prepare a graph of pH versus milliliters of NaOH and locate the equivalence and halfway
points.
2. Find the Ka for acetic acid.
3. Calculate the concentration of Acetic acid in vinegar used
38
Experiment (IX)
Potentiometric Titration
Potentiometric halide titration with Ag+
1.0 OBJECTIVE:
The purpose of this experiment is to determine chloride and iodide simultaneously in the
unknown samples.
2.0 INTRODUCTION
If a metal is placed in a solution containing its own ions, a potential difference E is
established between it and the solution, the value of which is given by the Nernst equation.
For the electrode reaction:
Mn+ + ne
M,
E = E° + (0.0591/n) log [Mn+] at 25°.
In precipitation titrations there is a sudden change in the ion concentration (Ag+ in this
case) at each equivalence point because the concentration of X- determines (through the
solubility-product equations) the concentration of Ag+. To detect the end point we can use an
electrode which responds to a change in the Ag+ ion concentration during the titration and
measure its potential after successive additions of the titrating agent.
The potential of the reference electrode is constant (Why?) and the change in emf is due
entirely to the indicator electrode. In the precipitation of Cl- and I- by Ag+, with the
concentrations used in the experiment, the AgI begins to precipitate first, and the E changes very
slowly until nearly all the I- is precipitated. At the first equivalence point [Ag+] will increase
more rapidly until the Ksp of AgCl is reached; here again the [Ag+] and E will remain fairly
constant until all the Cl- is precipitated, then another sharp increase in E occurs, indicating the
second equivalence point.
Figure 1: shows how a silver electrode can be used in conjunction with a calomel
reference electrode to measure Ag+ concentration. The reaction at the silver indicator electrode is
Ag+ + e-
Ag(s)
V
39
The reaction at the reference electrode is
Hg2Cl2(s) + 2e-
2Hg(l) + 2Cl-
V
The reference half-cell potential (E-, not E-°) is fixed at 0.241 V because the reference
electrode is saturated with KCl. The Nernst equation for the entire cell is therefore
E = 0.058 + 0.5916 log [Ag+]
As shown in Figure 1. The saturated calomel electrode has a double junction, like that in Figure
2. The outer compartment of the electrode is filled with KNO3, so there is no direct contact
between the KCl solution in the inner compartment and Ag+ in the beaker.
Figure 1. Use of silver and saturated calomel electrodes
to measure Ag+.
40
Figure 2. Double-junction silver-silver chloride reference electrode. The inner electrode is the
same as SCE. The outer compartment can contain KCl solution or any o ther solution that is
more compatible with the analyte solution. If you do not want Cl- to contact the analyte, the outer
electrode could be filled with KNO3 solution. Because the inner and outer electrolyte solutions
mix slowly, the outer electrode should be refilled occasionally with fresh KNO3 solution.
That is, the voltage of the cell in Figure 1 provides a direct measure of [Ag+]. Ideally, the
voltage changes by 59.16 mV (at 25¡É) for each factor of 10 changes in [Ag+].
In Figure 3 we used a silver indicator electrode and a glass reference electrode. The glass
electrode responds to pH. The cell in Figure 3 contains a buffer to maintain a constant pH.
Therefore, the glass electrode remains at a constant potential; it is being used in an
unconventional way as a reference electrode.
Figure 3. Apparatus for
measuring the titration curves in this
experiment. The silver electrode
responds to changes in the
Ag+ concentration, and the glass
electrode provides a constant
reference
potential
in
this
experiment. The measured voltage
changes by approximately 59 mV
for each factor of 10 changes in
[Ag+]. All solutions, including
AgNO3, were maintained at pH 2.0
by using 0.010 M sulfate buffer
prepared from H2SO4 and KOH.
41
PROCEDURE:
1.. Pour your unknown carefully into a 50- or 100-mL beaker. Dissolve the solid in about 20
mL of water and pour it into a 100-mL volumetric flask.
2. Rinse the sample vial and beaker many times with small portions of H2O and transfer the
washings to the flask. Dilute to the mark and mix well.
.
3.2 Dry 1.2 g of AgNO3 at 105 ᵒC for 1 h and cool in a dessicator for 30 min with minimal
exposure to light. Some discoloration, is normal (and tolerable in this experiment) but should
be minimized.
4. Accurately weigh 1.2 g and dissolve it in a 100-mL volumetric flask.
3
.
5. Set up the apparatus as shown in Figure 3. The silver electrode is simply a 3-cm length of
silver wire connected to copper wire. The copper wire is fitted with a jack that goes to the
reference socket of a pH meter. The reference electrode for this titration is a glass pH electrode
connected to its usual socket on the meter. If a combination pH electrode is employed, the
reference jack of the combination electrode is not used. The silver electrode should be taped to
the inside of the 100-mL beaker so that the Ag/Cu junction remains dry for the entire titration.
The stirring bar should not hit either electrode.
4
.
6. Pipet 25.00 mL of unknown into the beaker, add 3 mL of bisulfate buffer, and begin
magnetic stirring.
7. Record the initial level of AgNO3 in a 50-mL burette and add 1 mL of titrant to the beaker.
Turn the pH meter to the milivolt scale and record the volume and voltage. It is convenient (but
is not essential) to set the initial reading to + 800 mV by adjusting the meter.
5
.
8. Titrate the solution with 1-mL aliquots until 50 mL of titrant have been added or until you
can see two clear potentiometric end points. You need not allow more than 15-30 s for each
point.
42
10. Record the volume and voltage at each point. Make a graph of millivolts versus milliliters
to find the approximate positions (±1 mL) of the two end points.
6
.
11. Turn the pH meter to standby, remove the beaker, rinse the electrodes well with water, and
blot them dry with a tissue. Clean the beaker and set up the titration apparatus again. (The
beaker need not be dry.)
7
.
12. Now perform an accurate titration, using 1 drop aliquots near the end points (and 1-mL
aliquots elsewhere). You need not allow more than 30 s per point for equilibration.
8
.
NOTE: Silver halide adhering to the glass electrode can be removed by soaking in
concentrated sodium thiosulfate solution. This thorough cleaning is not necessary between Step 6
and 7 in this experiment.
43
Experiment (IX) Report
Potentiometric Titration
Potentiometric halide titration with Ag+
Name:____________________
Instructor Name: ______________
ID No.: ___________________
Section No.: ________________
Data Sheet and Calculations:
1. Prepare a graph of millivolts versus milliliters and locate the end points.
2. Calculate milligrams of KI and milligrams of KCl in your solid unknown.
44