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 => [Al(C2O4)3]3- + H2O

(3) 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) Balance oxygen by writing H2O on the right hand-side.

(iii) Balance hydrogen by writing H+ on the left hand-side.(iv) 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 half-equation. Add the two half-equations to obtain the overall equation for the oxidation-reduction 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 => Mn2+ + 4H2O (2)

C2O42- => 2CO2 + 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 powder is obtained.

(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) in ~ 2 mL warm NaOH solution (b) 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+ + OH-

is 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 labelon 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 = --------------------------- X 100

theoretical mass

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 4- charges 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+ + E4- => ZnE2- (1)(colourless) (colourless) (colourless)

ZnT+ + E4- => ZnE2- + T- (2)(pink) (colourless) (colourless) (blue)

Only a small amount of the indicator is used in the reaction. Thus, during the titration, the E4- reacts with the free Zn2+ to give colourless ZnE2-. At the end of the titration, the E4-

finally 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.