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UNIVERSITI MALAYSIA SABAH
SEKOLAH SAINS DAN TECHNOLOGI
PANDUAN AMALI
SK 2161
KIMIA AMALI IV
NAMA:
NO. MATRIK:
KUMPULAN:
CONTENTS
EXPERIMENT 1
PREPARATION OF POTASSIUM TRISOXALATOALUMINATE (III)
TRIHYDRATE {K3[Al(C2O4)3].3H2O}
EXPERIMENT-2
DETERMINATION OF THE OXALATE CONTENT OF THE COMPLEX
{K3[Al(C2O4)3].3H2O}
EXPERIMENT 3
PREPARATION OF POTASSIUM BISOXALATODIAQUACHROMATE (III)
DIHYDRATE {K[Cr(C2O4)2(H2O)2].2H2O}
EXPERIMENT 4
PREPARATION OF A COMPLEX COMPOUND [Cu(NH3)4]SO4.H2O
EXPERIMENT 5
DETERMINATION OF ZINC ION WITH ETHYLENEDIAMINETETRA
ACETIC ACID (EDTA): A COMPLEX-FORMING CHELATING AGENT
EXPERIMENT 1
PREPARATION OF POTASSIUM TRISOXALATOALUMINATE (III)
TRIHYDRATE {K3[Al(C2O4)3].3H2O}
Introduction
Aluminium occurs widely in nature as aluminosilicate minerals and as bauxite,
Al2O3.xH2O from which the metal can be produced by electrolysis after dissolving in
molten cryolite, Na3AlF6. The metal is mainly used in aluminium alloys. The
organoaluminium compounds (e.g., Et3Al) are used in the catalysts involved in the
polymerisation of ethene.
Group 13 (IIIb) of the Periodic Table includes the elements boron (B), aluminium (Al),
gallium (Ga), indium (In) and thallium (Tl). The electronic configuration is ns2np1 and
the group valency is three. Aluminium (At. No. = 13) is present in the third period and
contains empty 3d orbitals and favours the formation of complexes with six coordination
number. It forms stable octahedral complexes with such ligands as 8-hydroxyquinoline
also known as “oxine” (C9H7NO) and with the oxalate anion (C2O42-). In this
experiment, you are to prepare an oxalato- aluminium (III) complex and study its
reactions.
Preparation
(1)
Dissolve 7 g (7/666.42 = 0.0105 mole) Al2(SO4)3.18H2O in 100 mL distilled
water. Prepare a solution of 4.0 g (4.0/40.0 = 0.100 mole) of NaOH in 20 mL of
distilled water and add to the above solution dropwise with stirring.
(2)
Filter the freshly precipitated Al(OH)3 using a vacuum filtration technique (see the
demonstrator) and wash the precipitate 6 times using 20 mL distilled water each
time.
(3)
Weigh 4 g (4/126.04 = 0.0317 mole) oxalic acid dihydrate and 6 g (6/184.2 =
0.0326 mole) potassium oxalate monohydrate and dissolve them in
100
mL distilled water in a 250 mL beaker. Add Al(OH)3 to the solution and heat
the mixture on a steam bath with stirring using a glass rod to dissolve
Al(OH)3
(sometimes a little precipitate may remain).
(4)
Evaporate the solution to about 30 mL using a rotary evaporator (see the
demonstrator) or a hot plate. Gravity filter the concentrated solution into a
250 mL beaker.
(5)
Add 70 mL of ethanol to the solution dropwise over 10 minutes and cool in ice to
precipitate completely colourless crystals of the product.
(6)
Vacuum filter the product as in step (2) and wash four times with 10 mL ethanol,
dry completely by vacuum suction.
(7)
Keep the dry product for the next experiment.
(8)
Weigh the product on a top loading balance. Transfer about 0.5 g into a small
beaker to carry out the following tests for the presence of potassium, aluminium
and oxalate in the product and the rest into a properly labelled vial and store over
CaCl2 in a dry jar. Hand in the jar to the Demonstrator/Technician.
Qualitative Analysis
(1)
Test for potassium
Your Demonstrator will show you the flame-test for potassium to be performed in
a fume cupboard using a platinum wire. Note the colour of the flame.
(2)
Test for aluminium
Dissolve about 0.2 g of the product in 5 mL water and add 2 mL of dilute sodium
hydroxide solution. Note any precipitate in the solution. Now prepare a fresh
solution of 0.2 g of the product in 10 mL dilute HCl, warm the solution and add
NaOH solution dropwise until a precipitate is formed and redissolved in excess
NaOH. Shake the solution after each addition.
(3)
Test for oxalate
Dissolve 0.5 g of the product in 20 mL distilled water. Divide the solution into
two almost equal portions. Add 5 mL dilute H2SO4, in one of the solutions. Heat
both the solutions and add in each solution 3-4 drops of dilute solution of
KMnO4. Note the colour of the resulting solutions.
Report:
(Result sheet is provided. Anwer the questions in the result sheet according to the
information given below)
(1)
Write an ionic equation for the formation of Al(OH)3 using Al2(SO4)3.18H2O
and NaOH. Indicate which is the limiting reagent in the reaction.
(2)
The formation of the trisoxalatoaluminate (III) complex anion is carried out by
reacting Al(OH)3 with the exact number of moles of the oxalate ion and oxalic
acid. The number of moles of each reactant is given in step (3) of the preparation
section. Using this information, balance the equation for the reaction:
Al(OH)3 + (C2O42-) + H2C2O4
(3)
=> [Al(C2O4)3]3- + H2O
Aluminium is present as a free cation when Al2(SO4)3.18H2O is dissolved in
water. Which ions are present when K3[Al(C2O4)3].3H2O is dissolved in water?
The function of acid used in carrying out test (ii) is to release Al3+ from the
complex ion, [Al(C2O4)3]3-. The [Al3+(H2O)6] ion written in simple form as
Al3+ is then identified as described in test (ii). Write equations for the reaction of
the acid with the product and for the reaction which indicates the presence
of aluminium (III) in the product.
(4)
Aluminium (III) hydroxide is an amphoteric substance i.e., it reacts both with
acids and alkalis to form the aluminium salt of the acid and AlO2- anion
respectively. Write equations for the reactions of Al(OH)3 with HCl and NaOH
to indicate the amphoteric nature of Al(OH)3.
(5)
The oxalate ion is oxidised by the permanganate ion in the acid solution. Follow
the following steps to derive the two half ionic equations for this reaction:
Oxidation of oxalate ion to carbon dioxide
(i)
Write C2O42- on the left and CO2 on the right of =>.
(ii)
Balance the carbon and oxygen atoms.
(iii) Balance the charge by writing electrons (e-) on the right hand-side.
Why is it called an oxidation reaction?
Reduction of MnO4- to Mn2+ in acid solution
(i)
Write MnO4- on the left and Mn2+ on the right of =>.
(ii)
(iii)
(iv)
Balance oxygen by writing H2O on the right hand-side.
Balance hydrogen by writing H+ on the left hand-side.
Balance the charge by writing electrons (e-) on the left hand-side.
Multiply with the appropriate factor to obtain the same number of electrons in each halfequation. Add the two half-equations to obtain the overall equation for the oxidationreduction reaction. Why is it called reduction reaction?
What evidence did you see for this redox reaction in test (iii) to indicate the presence of
oxalate in your product?
(6) The presence of water of crystallation in the complex can be tested by heat a small
quantity of the complex in a test tube and holding a spetula containing anhydrous CuSO4
in the vapours. Note the colour of CuSO4. If it turns blue, water is confirmed. Write a
balanced equation for this test after finding the information from the book about copper
sulphate. Note the presence of water can also be tested from the I.R. spectra of
K3[Al(C2O4)3].3H2O.
(7)
Calculate the % yield of the product.
(8)
Draw the structure of the oxalatoaluminate anion (read the Introduction Section),
taking oxalate as a bidentate ligand.
EXPERIMENT-2
DETERMINATION OF THE OXALATE CONTENT OF THE COMPLEX
{K3[Al(C2O4)3].3H2O}
Introduction
The complex, potassium trisoxalatoaluminate (III) trihydrate, contains potassium as the
counter ion and the complex anion, [Al(C2O4)3]3- with three moles of water. The purity
of the complex can be determined by analysing the quantity of one or more of its
constituents (potassium, aluminium and oxalate) and comparing the experimental value
with the calculated value. The relative percentages of all the constituents can be
determined in order to confirm the above stoichiometry of the complex.
In this experiment, you are to determine the oxalate content of the complex and therefore
find the purity of the complex which is decomposed by addition of excess sulfuric acid.
K3[Al(C2O4)3] + 3H2SO4
=> 3H2C2O4 + 3K+ + Al3+ + 3SO42-
(1)
The oxalic acid thus released into the solution can be determined by titration with a
standard solution of KMnO4 which acts as an oxidant. The oxidation state of manganese
in KMnO4 is (VII) i.e., Mn+7 which is reduced direct to manganese (II) i.e., Mn+2 in
hot acid solution (or else it is converted to MnO2) by accepting five electrons from the
oxalate ion which in turn acts as a reductant and is oxidised to CO2 as shown in the
reaction equations:
MnO4- + 8H+ + 5e
C2O42-
=>
Mn2+
=> 2CO2
+
4H2O
(2)
+
2e
(3)
Exprimental
Carry out the following steps for the determination of oxalate.
(1)
Weigh out accurately using an analytical balance about 0.2 g of the complex in a
weighing bottle. Note the mass of the complex up to the fourth decimal place. Carefully
transfer the solid complex into a 250 mL conical flask. Add about 5 mL of 2M H2SO4 to
the weighing bottle using a small measuring cylinder. Use a glass rod to dissolve any
remaining complex in the weighing bottle and add the solution into the flask. Use a
further two times about 5 mL of 2M H2SO4 for washing the weighing bottle and add the
solutions to the flask. Now add an additional 10 mL of 2M H2SO4 to the flask.
(2)
Heat the solution on a hotplate, at a low dial setting, so as not to let the solution
boil but the flask be just “hot” enough to touch. While the solution is being heated
on the hotplate, wash a 50 mL burette with distilled water and then twice with 5 mL of
the standard (0.02M) KMnO4 solution provided and finally fill the burette to the mark.
Titrate the hot oxalate solution adding slowly KMnO4 solution from the burette to a
permanent pink colour. Note the volume of the KMnO4 solution. Repeat the above steps
(1) and (2) to obtain two concordant results i.e., two values should be within 0.1 mL for
the 50 mL burette.
(3)
Calculate the volume of the KMnO4 solution required in each titration for 1 g of
the complex by dividing each volume with the mass of the complex used for each
titration. Calculate the average volume and find the number of moles of KMnO4 in that
volume for 1 g of the complex.
(4)
Write the balanced overall reaction for the oxidation of oxalate by permanganate
from the above half equations (2) and (3) and calculate the equivalent number of
moles of oxalate ion reacted with the above number of moles of KMnO4. Now
multiply the number of moles of oxalate with its molar mass [mol. mass = 88] to
obtain the mass of oxalate in 1 g of the complex and then convert it to a percentage value.
(5)
Calculate the theoretical value of the percentage of oxalate in the complex,
K3[Al(C2O4)3].3H2O [mol. mass = 462.39]. Find the percentage purity of the
complex from the relationship:
% Purity = experimental value/ theoretical value x 100
Report
(Result sheet is provided. Anwer the questions in the result sheet according to the
information given below)
(1)
Show all the weighings and calculations carried out to find the purity of the
complex.
(2)
Comment on your result regarding the purity of the complex.
(3)
Comment as to why no indicator is used in this titration and state the difference in
reading the burette when using a coloured solution such as KMnO4 and a
colourless solution such as NaOH.
(4)
Read the Introduction Section carefully and state the reason as to why is it
necessary to carry out the KMnO4 titration in hot acid conditions?
EXPERIMENT 3
PREPARATION OF POTASSIUM BISOXALATODIAQUACHROMATE (III)
DIHYDRATE {K[Cr(C2O4)2(H2O)2].2H2O}
Introduction
Chromium (Cr) is found in group 6 (VIA) of the Periodic Table along with molybdenum
(Mo) and tungsten (W). These metals, therefore, have 6 valence electrons. In the
chromium atom, the 3d-orbital is half-filled with 5 electrons and 6th electron is present in
3s orbital. The most common ore of chromium is chromite [chrome ironstone/iron (II)
chromite FeCr2O4 which can be used directly for the manufacture of steel and other
chromium compounds. An alloy of chromium and nickel, nichrome is used for electrical
heating elements.
Chromium (III), the most stable oxidation state of chromium, forms a number of
complexes, in which six monodentate ligands surround the cation octahedrally. In the
above compound, two oxalate anions (bidentate ligands) and two water molecules
coordinate the chromium cation to provide oxygen atoms at the corners of the
octahedron. Two geometrical isomers exist for the complex ion, [Cr(C2O4)2(H2O)2]-,
two (C2O4)2- ions [or two H2O molecules] can be on the same side of octahedron
structure of the complex ion giving the cis isomer or on the opposite sides giving trans
isomer.
Preparation
(1)
Using a pestle and mortar, powder 7 g (7/126.04 = 0.0555 mole) oxalic acid
dihydrate and separately 2 g (2/294.2 = 0.0068 mole) K2Cr2O7. Mix the
powders thoroughly, regrinding them gently.
(2)
Transfer the mixture into a 100 mL beaker. Cover with a small watch glass and
gently heat the beaker on the hot-plate. A vigorous reaction commences
with the
evolution of steam and carbon dioxide and with the formation of a
dark semi-solid compound.
(3)
Remove the beaker from the hot-plate and immediately add 15 mL absolute
ethanol. Stir with a spatula and transfer the mixture into a small mortar (or stir
with a beaker and glass rod). Continue stirring for further 5 minutes breaking up
the product and decant the ethanol. Add a fresh 15 mL absolute
ethanol. Grind
the product using a beaker and glass rod until a crystalline violet
obtained.
powder
is
(4)
Filter the crystals by vacuum suction using a sintered glass crucible. Wash the
crystals two times with 5 mL ethanol. Dry in the air by vacuum suction. Weigh
the product on a top loading balance, transfer into a properly labelled vial.
(5)
Carry out the following tests on the product and hand in the remaining compound
to the Demonstrator/Technician .
Qualitative Analysis
(i)
Test for potassium
You will be shown the flame-test for potassium to be performed in a fume
cupboard using a platinum wire. Note the colour of the flame.
(ii)
Test for chromium (III)
Dissolve about 0.5 g of the complex in 20 mL water in a 50 mL beaker. Take 10
mL of the solution in a small beaker and add 5 mL dil. NaOH. Note the colour if a
precipitate is formed. Take 10 mL of the fresh solution add 10 mL dil. H2SO4 and heat
the solution. Then add dropwise 6M NaOH until a precipitate is formed. Note the colour
of the precipitate. Filter to remove the precipitate. Keep the filtrate for the next test.
Redissolve a small quantity of the precipitate:
(a)
(b)
in ~ 2 mL warm NaOH solution
in ~ 2 mL conc HCl.
Note the colour of the solutions.
(iii)
Test for oxalate
Heat the filtrate from the above test with 10 mL dil. H2SO4 and using a litmus
paper, make sure that the solution is acidic. Add dropwise dilute KMnO4
solution shaking the solution after each addition. Note the colour of the resulting
solution.
(iv)
Test for cis and trans isomers.
The two isomers are in equilibrium when the solid complex is dissolved in water.
The cis-form is more soluble and exists in high proportion in solution but it is
slowly converted into the trans-form when the solution is evaporated slowly. In
the solid state only one isomer can exist depending on the method of
preparation of the complex.
Carry out the following test to determine the nature of the isomer prepared in this
experiment.
(a)
Moisten a few crystals of the product placed on a small filter paper resting on a
watch glass, first with a few drops of dilute ammonia solution and then with dilute
hydrochloric acid. Record your observations. With the cis -form, a deep-green
solution rapidly forms and spreads over the filter paper, no solid remains. With
the trans-form there is no green colour, but a light-brown solid forms and
remains undissolved. In each case, reaction with dilute HCl restores the
original
colour of the isomeric form.
Report
(Result sheet is provided. Anwer the questions in the result sheet according to the
information given below)
(1)
What is the oxidation state of chromium in K2Cr2O7 and in the product? Write
two half-equations for each reaction involved in the preparation of the complex
using oxalic acid dihydrate (H2C2O4.2H2O) and potassium dichromate
(K2Cr2O7) and a balanced equation for the overall reaction.
(2)
Which is the limiting reagent in the above preparation?
(3)
Calculate the % yield of the product.
(4)
Write the chemical equations for the reactions involved in the tests for the
detection of chromium (III) and oxalate ions. Indicate in each case whether it is an
acid-base or oxidation-reduction reaction.
(5)
Draw the structures of the two possible geometrical isomers of the complex. What
is your conclusion regarding the nature of the isomer in the test (iv)? When dilute
ammonia is added to the solid complex, one molecule of the ligand H2O replaced
by OH- ion. Write equations for both the isomers to represent this reaction as well
as for the reaction of the hydroxo product and HCl to restore the original
complex.
Note :
By comparing the I.R. spectrum of your product with the I.R. spectra of calcium
oxalate monohydrate, anhydrous calcium oxalate, you can Identify oxalate and
water bands in the I.R. spectrum of your product.
EXPERIMENT 4
PREPARATION OF A COMPLEX COMPOUND [Cu(NH3)4]SO4.H2O
INTRODUCTION
Ammonia (NH3) is a colourless, pungent smelling gas. The molecule has a pyramidal
structure. It is more soluble in water than any other gas. The equilibrium constant for the
reaction
NH3(aq) + H2O <=> NH4+ + OHis only 1.8 X 10-5. This means that it mostly exists as NH3(aq) in water. Ammonia is one
of the most used chemical in industry. It is mostly used in the production of fertilizers
(ammonium sulphate and ammonium nitrate); manufacture of explosives, plastics, pulp &
paper, textiles, rubber, refrigerants, detergents, insecticides and food additive. It is also
used in window cleaning detergents.
Because of the presence of a lone pair of electron on nitrogen atom in ammonia, it forms
a number of complexes with cations involving dative or coordinate bond. It acts as a
monodentate ligand.
When ammonia is added to an aqueous solution containing copper(II) ions, a deep blue
complex cation [Cu(NH3)4]2+ is formed.
In this experiment the compound
[Cu(NH3)4]SO4.H2O is isolated by adding alcohol to the aqueous solution of the complex.
The alcohol is less dense than water and although it is soluble in water, it can float on the
top of water if the two liquids are not mixed. If alcohol is poured carefully onto the
surface of the liquid, the slow diffusion of the alcohol down into the water results in the
slow growth of the crystals. Slowly grown crystals of [Cu(NH3)4]SO4.H2O usually are
long needles and deep blue in colour.
PROCEDURE
1. Take a clean 100 mL beaker. Write your name on a label and put it on the beaker.
2. Weigh about 6.3 g of finely powdered CuSO4.5H2O in the beaker. Add 6 mL of
distilled water and stir with glass rod to dissolve the solid. Note the colour of the solution.
3. Carry out further procedure inside the fumehood, add 10 mL of concentrated aqueous
ammonia and stir thoroughly. Note the colour of the solution.
4. From the nozzle of a wash (plastic) bottle, carefully pour 10 mL of ethanol down the
side of the beaker so that the alcohol runs onto the top of the solution. Leave it overnight
in the fumehood undisturbed. Cover it with an inverted 400 mL beaker. The compound
will crystallize as long needles when alcohol slowly diffuses into the aqueous solution.
Leave the beaker overnight in the fumehood.
5. At the next session, get ready the following wash solutions.
a. A mixture containing 5 mL of ethanol and 5 mL of concentrated aqueous ammonia.
b. 10 mL ethanol.
c. 10 mL acetone.
6. First note point (7). Set up a suction filtration apparatus using a sintered glass crucible.
Use the solution in the beaker and a rubber policeman to transfer the bulk of the crystals
into the crucible. Use a rubber policeman to remove the crystals from beaker to the
crucible.
7. Add each wash solution quickly so that the crystals are not exposed to air until after
the acetone has gone through the crucible. Apply slow vacuum and quickly wash the
crystals first with ethanol ammonia mixture then with pure ethanol and finally with the
acetone.
8. Dry the crystals by pressing them between two pieces of filter paper. Then transfer the
dry crystals quickly to a clean, dry and previous weighed vial or plastic bag. Put a label
on the vial or bag.
9. Weigh the crystals and calculate the theoretical yield of [Cu(NH3)4]SO4.H2O from the
reaction equation.
CuSO4.5H2O + 4 NH3
=> [Cu(NH3)4]SO4.H2O + 4 H2O
10. Calculate the % yield of the product.
actual mass obtained
% yield = --------------------------theoretical mass
X 100
11. Save the product for the next experiment. Quickly remove a few crystals from the
container. Carry out the following tests.
11.1 Dissolve crystals in 3-4 drops of concentrated HCl in micro test tube and carry out
the flame test. Note the colour of the flame. [test for copper].
11.2 Dilute the solution in the micro test-tube and add a few drops of 0.1M BaCl2
solution. Note the colour of the precipitate. [test for sulphate].
11.3 Heat the crystals in a test tube and bring a glass rod dipped in concentrated HCl near
the mouth of the test-tube and note any dense fumes of NH4Cl. [test for NH3].
11.4 Suggest a test to confirm the presence water in the product.
EXPERIMENT 5
DETERMINATION OF ZINC ION WITH ETHYLENEDIAMINETETRA
ACETIC ACID (EDTA): A COMPLEX-FORMING CHELATING AGENT
INTRODUCTION: In the Coordination Chemistry, you learnt that a ligand is an ion or a
neutral molecule that forms a special covalent bond (dative or coordinate) with a metal
cation or a neutral metal atom by donating a pair of electrons which are shared by the
ligand and the metal. A chelate is a cyclic or ring compound formed when a cation is
bonded by two or more donor groups present in a single ligands. EDTA is a good
example of the ligand which form chelates with a number of metal cations.
An application of coordination chemistry is the determination of the concentration of
metal cations using EDTA as a chelating agent. The determination of the concentration of
a metal ion in a solution is an important analytical technique. In chemical industry it is
important to determine the concentration of metal cations discharged from the
manufacturing plants or calcium and magnesium in hard water.
In this experiment, the presence of Zn2+, a metal cation that might be discharged from an
industrial plant, will be determined with EDTA which can exist as an anion with 4charges by loss of its 4 protons. The anion reacts with many metal cations in 1:1 ratio
irrespective of the charge on the metal cation. For zinc the complex is represented as
ZnE2-.
[Zn2+ + E4- <=> ZnE2-]
It is necessary to determine when exactly one E4- per one Zn2+ is present in a titration.
This determination is done by titrating (gradually adding) E4- to a solution of Zn2+ in the
presence of an indicator. The indicator used in this experiment is a big organic molecule
called Eriochrome Black T. In the solution the indicator exists as an anion represented by
T-. In the presence of zinc ion, the indicator forms a 1:1 complex ion which is pink in
colour. [Zn2+ + T- + <=> ZnT+].Only a small amount of T- is used to give enough ZnT+ to
see the colour. All T- is complexed to the zinc ion. Since the Zn2+ is present in excess, the
majority of the zinc ion is not complexed. As EDTA is added to a solution during the
titration, the EDTA can react with both the zinc ion and the coloured complex ion, ZnT+,
as follows.
Zn2+
(colourless)
ZnT+
(pink)
+
+
E4(colourless)
E4(colourless)
=>
=>
ZnE2(colourless)
ZnE2(colourless)
+ T(blue)
(1)
(2)
Only a small amount of the indicator is used in the reaction. Thus, during the titration, the E4reacts with the free Zn2+ to give colourless ZnE2-. At the end of the titration, the E4finally reacts with the ZnT+ to replace T- with E4-. The T- released is blue and a colour of
the solution changes from pink to blue.
METHOD OF CALCULATION: According to the balanced equation (1). the number of
moles of E4- used equals the number of moles of Zn2+ present in the sample. The number
of moles of E4- used is calculated from its volume and its known concentration. The same
number of moles of Zn2+ are present in its volume. The molarity of the zinc in the
solution can be calculated.
moles of Zn2+
Molarity of zinc = -------------------------------------volume of zinc solution used
PROCEDURE
1. Transfer 10 mL of zinc solution from the automatic dispenser in a 250 mL titration
flask. Add about 10 mL distilled water using a graduated cylinder. Add 5 mL of the
buffer solution using a graduated cylinder. Add a few grains of the solid indicator. A light
pink colour should appears as the indicator dissolves.
2. Rinse a 50 mL buret with about 5 mL EDTA standard solution (0.01 M). Fill the buret
with the solution close to the mark.
3. Titrate the zinc solution with the EDTA solution shaking the flask carefully until the
pink colour changes to blue. Record the volume of the EDTA solution.
4. Repeat the titration two more times so that you obtain two concordant readings which
agree within 0.1 mL.
5. Calculate the molarity of the zinc solution as suggested above. Calculate the average
value.