32

The Islamic University of Gaza

Faculty of Science

Chemistry Department

Practical Inorganic Chemistry

2014

Prepared

by

Dr. Salman Saadeh Dr. Nabil El-Halabi(I). Synthesis of Cr(acac)3

This is accomplished by reacting Cr(H2O)63+ in aqueous solution with acetylacetone and urea. Urea hydrolyzes (reacts with water) at a convenient rate to ammonia and carbon dioxide:

(NH2)2CO + H2O 2NH3 + CO2

Ammonia reacts further with water to give NH4+ and OH-, making the solution slightly basic. This allows formation of the more reactive (H2O)5CrOH2+ species, which may react with acac to give (H2O)4Cr(acac)2+. Continued hydrolysis of urea, though, keeps the solution slightly basic and effects conversion of (H2O)4Cr(acac)2+ to give (H2O)2Cr(acac)2+ and finally to Cr(acac)3. The product is nonpolar and electrically neutral and has very limited solubility in water. It thus precipitates from the aqueous solution shortly after its formation.

Procedure:

1) Place 100 mL of distilled water in a 200 mL-flask and add to it 0.02 moles of

chromium(III) chloride hexahydrate.

2) Add 0.33 mole of urea and 0.12 mole of acetylacetone.

3) Cover with a watch glass and heat the mixture by placing the reaction flask in a large

beaker of boiling water for 90 minutes.

4) Cool the reaction flask and collect the product by suction filtration, then wash the

crystals while still on the filter with several small quantities of distilled water and air

dry.

5) Transfer the crystals to a screw-cap bottle for storage. Calculate percent yield. Keep

the product in a closed bottle for future use in other experiments.

Questions:

1) What is the percent yield? Is it possible to get better than 100 percent yield of a crude

product? Explain.

2) Why is NaOH not used ( instead of urea) to make the solution basic?

3) Why was a steam bath used to heat the reaction flask instead of a direct flame?

4) How does urea make the solution basic?

(II). Recrystallization and melting point determination

of Cr(acac)3

The Cr(acac)3 complex is soluble in many solvents less polar than water ("organic solvents") and may be purified using a technique called recrystallization. One method of recrystallization uses two solvents- one in which the material is quite soluble, and one in which it is not very soluble. The original batch of impure crystals is dissolved in the "good" solvent and then recrystallized by adding "poor" solvent. The melting point of the product is taken for comparison with that of Cr(acac)3 as reported in the literature.

Procedure:

Test the solubility of the product in several solvents as follows:

Place about 0.1 g of the crude crystals in a test tube and add 1 mL of solvent and shake briskly. Record your observations. If the product does not dissolve in cold solvent heat the tube gently in hot water bath, then shake it and notice if the complex dissolves. Try these solvents: toluene, hexane, ethanol, chloroform, ...etc.

Make a table of your results by writing + for soluble and - for insoluble then decide which solvent or solvent combination is suitable for recrystallization.

A good combination for recrystallization is toluene-hexane. Dissolve the product in the minimum amount of warm toluene. Add hexane (or petroleum ether) dropwise with swirling until crystals begin to appear. Stopper the flask and allow to stand. Cool the flask in an ice bath. Collect the crystals by filtration and air dry. Calculate the percent yield on the amount of your product. Determine the melting point.

Questions:

1) Define: a good solvent, supersaturation.

2) What conditions must two solvents meet to be a good two-solvent recrystallization

system?

3) Write formulas for all solvents used in your solubility study of Cr(acac)3, and arrange

them in order of increasing polarity.

4) Many inorganic salts are very soluble in water but insoluble in alcohol. Suggest

a solvent pair to recrystallize such salts.

(III). Volumetric Determination of the Percentage of Cr in Cr(acac)3

The quantitative determination of the chromium content in a compound may be accomplished by the volumetric method presented in this experiment. This titration approach is accomplished by converting Cr(acac)3 to Cr2O72- in acid solution, adding an excess of Fe(II) ion, which reduces the Cr2O72- to Cr(III), and then titrating the excess Fe(II) ions with a standard K2Cr2O7 solution (back titration).

Cr(acac)3 Cr2O72-

14H+ + 6Fe2+ + Cr2O72- 2Cr3+ + 6Fe3+ +7H2O

Note: Although the indicator is colorless before the end point but the solution is green due to the reasonably high concentration of Cr3+ (green) in solution.

Procedure:

1) Prepare 200 mL. of a solution that is 0.10 N in K2Cr2O7 (primary standard)

2) Prepare as accurately as possible, 200 mL. of approximately 0.10 N Fe2+ by dissolving

the required amount of FeSO4.7H2O in 100 mL. of water. Carefully add 50 mL. of conc.

H2SO4. After the solution has cooled, dilute to the mark with water. Calculate the

normality of the K2Cr2O7 solution based on the exact quantity used and the final

volume of the solution.

3) For titration, pipet 20.00 mL. of the 0.10 N Fe2+ solution into a 250-mL. Erlenmeyer

flask. Cautiously add 10 mL. of conc. H2SO4 and 10 mL. of 85% H3PO4 and dilute to

approximately 125 mL. with water. Add 6-8 drops of the indicator solution, sodium

diphenylamine sulfonate, Titrate this solution with the standard K2Cr2O7 solution until the color changes from green to violet.

The volume used is V1. The titration should be done twice and the volumes of titrant

required should agree within 0.05 mL.

4) Weigh accurately 0.11-g sample of dried recrystallized Cr(acac)3 and transfer it

quantitatively to a 250-mL. Erlenmeyer flask. Add 5 mL. of conc. H2SO4 and 5 mL of

conc. HNO3 and swirl until the sample dissolves. Place the sample, partially covered

with a watch glass, on a hot plate turned low. After 10 min., raise temperature-control

to medium for 45 min. Do not boil the sample to dryness. After the oxides of nitrogen

(red-brown fumes) have been driven off, the reaction mixture should gently be reflux

1/2 to 1 inch up the side of the Erlenmeyer. The sample must be a clear green to blue-

green at the end of the digestion. If a precipitate has appeared, you must start again

using less vigorous conditions.

CAUTION: Boiling H2SO4 causes severe burns on contact.

5) Cool the sample in air for a couple of minutes and then further cool in water. Carefully

add 5 mL. of distilled water to the flask and then cool under a tap.

6) Add in several small portions, 20 mL. of 30% H2O2. Cool under the tap to room temperature.

7) Using a medicine dropper, add fairly rapidly approximately 35 mL. of 6 M NaOH. Cool

the flask frequently during this addition; otherwise it will get quite hot. It is necessary

to add enough base to make the solution strongly basic at this point.

(CAUTION: When the solution becomes strongly basic , there is a vigorous exothermic reaction, which will cause excessive foaming. The flask should be cooled to prevent the solution from foaming over.)

The solution should be allowed to stand until any foaming has subsided. As NaOH is added the solution will go through several color changes and will become clear red-brown when it becomes basic.

8) Add 3 boiling chips to the sample. Heat gently at first then to boiling for 1/2 hour. This

should completely destroy the excess of H2O2. The sample should be a bright, clear

yellow.

9) Cool to room temperature and make the solution acidic by slow addition of 10 mL.

conc. H2SO4 followed by the addition of 10 mL. of 85% H3PO4. Cool the solution

again if it is warm.

10) Pipette 20 mL. of 0.10 N Fe2+ solution into the flask followed by 6-8 drops of the

above indicator. Titrate with K2Cr2O7 solution, the volume used is V2. Calculate the

percentage of Cr using this equation:

(V1 - V2)NCr2O72-(17.33)(100%)/weight of sample = % Cr

Questions:

1) Balance the equation: MnO4- + Fe2+ Mn2+ + Fe3+ ( in acidic solution)

2) Using your answer above, calculate the grams of Fe2+ in the sample if 0.123 g

KMnO4 was used.

3) What is a primary standard?

4) If H2O2 was not completely removed by boiling, what effect would this have on the

percentage of chromium?

5) Any undecomposed ligand in initial digestion will react with the oxidizing agent

Cr2O72- used in titrating the excess Fe2+. What effect will this leave in the percent of

chromium in the sample?

(IV). Synthesis of Cr(acac-NO2)3

The replacement of the 3-H in acetylacetonate ligand in Cr(acac)3 by the nitro group to yield tris(3-nitroacetylacetonato) chromium(III) is a good example of a metal influencing the reactivity of a ligand. The unbound compound, 3-nitroacetylacetone, cannot be made directly from acetylacetone.

Procedure:

1) Stir 8.0 mmoles of finely ground copper(II) nitrate trihydrate in 40 mL of acetic

anhydride at 0oC for 15 min. in a fume hood. Some of the copper compound will

dissolve to a deep blue solution.

2) Add 2.8 mmoles of Cr(acac)3 and stir the slurry at 0oC for 1 hr. and then at room

temperature for 45 min.

3) Pour the reaction mixture slowly and carefully into a mixture of 120 g of ice, and 15

g of sodium acetate while stirring the latter mixture.

4) Collect the resulting dark red precipitate by suction filtration, and wash it with small

amounts of water and ethanol.

5) Recrystallize the crude material from chloroform-ethanol. The expected amount of

pure product is 1.0 g. Calculate percentage yield.

Questions:

1) Why is copper nitrate used in excess?

2) Write the formula for each of these anhydrides: Ca(OH)2, H2SO4, HNO3,

CH3COOH.

3) Complete the reaction : CH3COONO2 + Cr(acac)3

4) What precautions are advised in handling large amounts of Cr(acac- NO2)3?

(V). Synthesis of Co(acac)3

To synthesize Co(acac)3, we make use of the fact that there are two stable oxidation states of cobalt, Co(II) and Co(III). The lower oxidation state, Co(II), undergoes substitution much more readily than does Co(III). Thus we can form Co(II)-acac complex, Co(acac)3-, by direct and rapid substitution and then oxidize Co(acac)3- to Co(acac)3.

Procedure:

1) Slurry 0.021 mole of powdered Cobalt (II) carbonate and 0.2 mole of acetylacetone in

a 125-mL. flask.

2) Heat the mixture to 100oC on a hot plate and transfer the flask immediately to

a magnetic stirrer and add 15 mL. of 10% H2O2 at a rate of 1 to 2 mL./min. while

stirring.

3) Reheat the reaction mixture to incipient boiling, then add 15 mL. more of H2O2 as

before, then heat to boiling .

4) Cool the reaction mixture in an ice-salt bath for 0.5 hr. and filter the green-black

sludge with suction . Wash with water then with small amounts of cold ethanol. The

yield is about 6.0 g.

5) Recrystallize from toluene- petroleum ether as done for Cr(acac)3.

6) Dry at 110oC and calculate percent yield.

Questions:

1) Suggest a mechanism of the reaction that proceeds through two steps, the first to

produce Co(acac)3- and the second to oxidize it. Designate the ligand as acacH.

2) What is the molar ratio acac: CoCO3 and H2O2:CoCO3 in the overall equation? What

is the actual ratio used?

Account for the large divergence between the stoichiometric and actual ones.

3) How might the compound Na[Co(II)(acac)3] be converted to Co(acac)3?

(VI). Synthesis of Copper(II) Acetate Monohydrate

Introduction

The first compound with metal-metal bonds was found to have the formula Ta6Cl14.7H2O. Its structure was shown to consist of an octahedron of mutually bonded tantalum atoms, with each edge of the octahedron being bridged by a chlorine. The formula would be represented by [Ta6Cl12]Cl2. Hundreds of compounds with metal-metal bonds are now known. Compounds are currently known containing not only M-M single bonds, but also double, triple, and even quadruple bonds.

Structure of rhodium(II) acetate alcoholate

Rh

Rh

O

O

C

L

L

O

O

C

C

H

3

O

O

C

C

H

3

C

H

3

O

C

C

H

3

O

Copper(II) acetate dihydrate exhibit a structure similar to that of rhodium(II) acetate alcoholate which has a short metal-metal bond that is quite short (2.386 Ang.) which indicates the presence of a rhodium-rhodium bond in this compound. Copper(II) acetate dihydrate may contain a copper-copper bond where there is only a weak coupling of the unpaired electrons on the Cu(II) ions (d9). Thus, while the ground state is diamagnetic, there is a low energy excited state that is paramagnetic. This excited state is appreciably populated at room temperatue and the complex therefore appears to be paramagnetic. As the temperature increases, the magnetic moment increases as well. An alternative explanation to this behavior is described as an antiferromagnetic coupling of the unpaired spins. Structural determination shows that the two copper atoms are separated by a distance of 2.64 . This compares with a 2.56 . interatomic distance in metallic copper, making the assumption of a Cu-Cu bond more controversial.

Copper(II) acetate dihydrate is prepared by a route involving the intermediate formation of a tetraamine complex, conversion of the tetraamine complex to a precipitated hydroxide, and subsequent reaction with acetic acid.

CuSO4.5H2O(aq) + 4NH3(aq) [Cu(NH3)4]2+(aq) + SO42-(aq)

[Cu(NH3)4]2+(aq) + 2NaOH Cu(OH)2(s) + 4NH3(aq) + 2Na+(aq)

2Cu(OH)2(s) + 4CH3COOH [Cu(CH3COO)2.H2O]2(aq)

The solvent ligands can be easily removed by heating in vacuum to yield the non-adducted complexes.

Procedure:

1) In a 10-mL beaker equipped with a magnetic stirring bar, dissolve 160 mg (1.0 mmol) of copper(II) sulfate or 250 mg (1.0 mmol) of copper(II) sulfate pentahydrate in 5.0 mL of water. Stir the mixture, and warm it to 40-50o C on a sand bath (or a hot plate) to aid the dissolution.

2) Using a Pasteur pipet, add 50% NH3 (aq) to the warm, stirred, light blue solution, until the intense blue color of the copper ammonium complex is evident. During this addition, a precipitate of copper hydroxide may form initially, but will dissolve on further addition of the NH3 (aq) solution.

3) Add 80 mg (2.0 mmol) of sodium hydroxide flakes to the deep blue solution, and stir the mixture for 15-20 min at 55-65oC. A light blue solid of copper(II) hydroxide precipitates during this time.

4) Allow the mixture to cool to room temperature and collect the precipitate by vacuum filtration using a Hirsch funnel. Wash the blue solid with three 2-mL portions of warm water.

5) Transfer the amount of Cu(OH)2 to a 10-mL beaker, and dissolve it in the minimum amount of 10% acetic acid. Warming on the sand bath with stirring aids the dissolution process. Concentrate the solution nearly to dryness (HOOD) by warming it on a sand bath under a slow stream of nitrogen.

6) Collect the beautiful deep blue crystals that form by filtration using a Hirsch funnel. Dry the product on a clay plate or on filter paper.

Characterization of Product:

Acquire the IR spectrum of the product as a KBr pellet. Determine the magnetic moment of the product. Does it correspond to that of a diamagnetic or paramagnetic complex at room temperature? If paramagnetic, how many unpaired electrons seem to be present?

Questions:

1) What physical indications, other than magnetic moment, might lead one to conclude

that a metal-metal bond is present?

2) Even when the magnetic susceptibility seems to indicate that a metal-metal bond is

present, the low magnetic susceptibility might be due to other reasons. Discuss this

point.

3) One of the largest classes of metal-metal bonded compounds are the metal carbonyl

clusters. Discuss the bonding in two such members of this class.

4) Metal clusters are under active investigation as "mimics" to bulk metals in catalysis.

Perform a literature search and discuss several examples of metal clusters that were

studied in this manner.

References:

1) a. Cotton, F. A, . Multiple Bonds and Metal Clusters" in Reactivity of Metal-Metal

Bonds, M.H. Chisholm, Ed., ACS Symposium Series No. 155, American Chemical

Society, Washington, DC, 1981.

b. Cotton, F. A., Walton, R. A. Multiple Bonds Between Metal Atoms, Krieger:

Malabar, FL, 1988.

2) Catterick, J.; Thornton, P. Adv. Inorg. Chem. Radiochem 1977, 20, 291.

Kato, M.; Jonassen, H. B.; Fannin, J. C. Chem. Rev. 1969, 64, 99.

3) Cotton, F. A., Walton, R. Multiple Bonds Between Metal Atoms, second edition, Oxford University Press, 1993.

(VII). Preparation of Sodium Tetrathionate

Introduction:

Sulfur forms a large variety of oxo-anions such as the sulfate and the sulfite anions. The polythionates are a second class of sulfur-oxygen anions having a general formula SnO62- where n ranges to greater than 20. These anions contain more than one sulfur, are normally named according to the number of sulfur atoms present. Thus, the anion S4O62- is named the tetrathionate anion. Polythionates are stable only as salts-the free acids cannot be isolated as they decompose rapidly into S, SO2 and sometimes SO42-.

In general, polythionates are obtained by the reaction of thiosulfate, S2O33-, solution with sulfur dioxide in the presence of As2O3. Oxidizing agents like H2O2 and I2 also react with thiosulfate solutions to form polythionate salts.

In the thiosulfate anion, sulfur is in the II oxidation state. This anion can easily be oxidized to the tetrathionate anion where the sulfur atoms are in a mean oxidation state of 2.5. In this reaction, iodine is used as the oxidizing agent:

2Na2S2O3 + I2 Na2S4O6 + 2NaI

The reaction, which is generally carried out in aqueous medium, is a quantitative on. This is the basis of the quantitative application (iodometric method of titration) of this reaction in analytical chemistry.

Procedure:

The experiment works best when a slight excess of iodine is present.

In a 10-mL beaker, weigh out 240 mg of iodine and dissolve it in two mL of 95% ethanol. Using an agate mortar and pestle, grind a sample of sodium thiosulfate to a fine powder. Weigh out 250-300 mg of powdered sodium thiosulfate.

Add the powdered thiosulfate in several parts to the iodine solution, vigorously agitating the mixture using a glass rod or spatula before the addition of the next portion.

Since thiosulfate is insoluble in alcohol, thorough mixing of the reactants is necessary to ensure completion of the reaction. The mixture may be warmed (but don't boil, iodine may sublime off) to increase the rate of reaction.

At the end of the reaction, the slight excess of iodine should be left unreacted as indicated by a faint yellow color of the solution. If the mixture turns completely colorless, add one or two crystals of iodine to the solution to regenerate a persistent faint yellow color of iodine. This will ensure that no thiosulfate is left unreacted.

Isolation of product:

Collect the microcrystals of sodium tetrathionate by suction filtration using a Hirsch funnel. Wash the beaker and the product with several 500- L portions of ethanol, transferring the product as quantitatively as possible to the funnel. Wash the product with 500-L of ether and dry the product on the filter under suction. Determine the percentage yield of the product.

Characterization of Product:

Obtain the IR spectrum of the product as a KBr pellet and compare the spectrum to that of sodium thiosulfate. Obtain a TGA thermogram of the product and of sodium thiosulfate, and also determine the level of hydration of both materials.

Determination of the Presence of Sodium Iodide:

Sodium iodide is a byproduct of this reaction. Although it is fairly soluble is alcohol, it is important to make sure that the product is not contaminated with traces of NaI.

Take a small amount of the product in a test tube. Add 1 or 2 drops of concentrated H2SO4 directly to the product, followed by a small amount of solid MnO2. Warm the mixture on a flame. If no violet fumes of I2 are observed there is no iodide in the mixture.

Questions:

1) What effect does the increase in oxidation state of sulfur from 2 (thiosulfate) to 2.5

(tetrathionate) have on the IR frequency of the S-O stretch?

2) The polythionate free acids are not stable. What do they decompose into?

3) The polythionates can be viewed as being derivatives of the sulfanes. What is a sulfane

and how is one prepared?

4) Dithionic acid, H2S2O6, appears to be a simple acid analog to the polythionates;

however, the acid and its salts do not show similar chemical behavior. From a search

of the literature, discuss the similarities and differences of dithionic acid and the

polythionates.

Reference:

1) Janickis, J. Acc. Chem. Res. 1969, 2, 316.

(VIII). Preparation of Copper-Glycine Complexes

Introduction:

Like acetylacetone, the amino acid glycine, NH2CH2COOH, dissociates to form an ion (gly) which can coordinate to a wide variety of metal complexes. A major difference is that the glycine anion is not symmetric and cis and trans structural isomers can arise depending on the relative orientation of the ligands.

The direct reaction of copper(II) acetate monohydrate and glycine results in an equilibrium mixture of the two isomers.

[(CH3CO2)2Cu.H2O]2 + H2NCH2CO2H

cis-Cu(gly)2.H2O + trans- Cu(gly)2.H2O

The isomer precipitates much more quickly than the trans, leading to a shift in the equilibrium away from the trans, producing only the cis product. Interestingly, even though the cis isomer is the kinetically favored product, the trans isomer is thermodynamically favored. The cis isomer may be converted to the trans simply by heating at 180oC for 15 min.

Procedure:

Preparation of cis-Bis(glycinato)copper(II) Monohydrate:

In a 10-mL flask containing a magnetic stirring bar, dissolve 100 mg (0.5 mmol) of copper(II) acetate monohydrate in 1.5 mL of hot deionized water. Add 1.0 mL (calibrated Pasteur pipet) of hot 95% ethanol to the solution. Maintain the temperature of the solution at 70oC.

In a separate 10-mL Erlenmeyer flask, dissolve 75 mg (1 mmol) of glycine in 1 mL of hot deionized water by swirling the flask. Transfer the glycine solution (Pasteur pipet) into the copper(II) acetate solution and stir briefly. Discontinue the stirring and allow the solution to cool to room temperature.

Isolation of Product:

Complete the precipitation of product by transferring the Erlenmeyer flask into a ice-water bath for 10 min. Collect the product by suction filtration using a Hirsch funnel. Wash the product with a 100 mL portion of ice-cold ethanol and dry the crystals on a clay tile or on filter paper.

Preparation of trans-Bis(glycinato)copper(II):

Place ~35 mg of the cis product produced in part A into a stoppered side arm test tube. Place the test tube on an aluminum block set on a magnetic stirring hot plate and heat the block to approximately 220oC for 15 min.

Note: If a muffle furnace is available, it may be used instead of the aluminum block.

Remove the test tube from the block and allow it to cool to room temperature.

Characterization of products:

Obtain an IR spectrum (KBr pellet) of the cis and trans products. If a far-IR spectrometer is available, prepare the material as a Nujol mull. The Cu-N stretches may be observed from 450-500 cm-1, and the Cu-O stretches from 250-350 cm-1. Differential scanning calorimetry is also informative, as the temperature for cis to trans conversion may be easily determined.

Questions

1) Assign the heavy atom framework for the cis and trans isomers to its proper point

group.

2) Why is the IR spectrum of the trans compound much simpler than that of the cis,

especially in the fingerprint region (800-1200 cm-1)?

3) When a carboxylic acid ligand is monodentate, the separation of the symmetric and

antisymmetric C=O stretching frequencies increases, compared to the free acid. Why?

4) Which of the other amino acids would you expect to give rise to similar complexes

having geometric isomers? Search the literature to determine if any were prepared.

References:

1) Delf, B. W.; Gillard, R.D.; O'Brien, P. J. Chem. Soc. Dalton Trans. 1979, 101.

2) O'Brien, P. J. Chem. Edu. 1982, 59, 1052.

(IX). Acylation of Ferrocene

Introduction:

Ferrocene has extensive chemistry. It very readily undergoes Friedel-Crafts acylation and alkylation reactions. In fact, ferrocene is 106 times more reactive than benzene towards acetylation. Thus, in the presence of a Friedel-Crafts catalyst, e.g. phosphoric acid, ferrocene can be conveniently acetylated. The catalyst generates the electrophile [CH3CO]+ from acetic anhydride, which then attacks the cyclopentadienyl ring of ferrocene forming monoacetyl and diacetyl products. At the end of the acetylaion reaction, one obtains a product mixture of unreacted ferrocene, and the mono- and diacetylated ferrocenes. Thin-layer chromatography may be used to quickly identify how many components there are in the product mixture, and which solvent is the best for efficient separation of these components. Column chromatography is used to separate and purify the products. High-performance liquid chromatography may be used to follow course of the reaction, as well as for the detection of reaction products.

Procedure:

Place 150 mg (0.81 mmol) of ferrocene in a 10-mL round-bottom flask equipped with a magnetic stirring bar. Using a Pasteur pipet, add 2-3 drops of 85% phosphoric acid. Add 500mL (0.525 g, 8.7 mmol) of acetic anhdydride, using an automatic delivery pipet. Attach an air condenser and a CaCl2 drying tube. Heat the mixture in a sand bath at 100oC for 10 min. Add ~2 g of crushed ice to the reaction mixture. When the ice has all melted, neutralize the mixture by adding solid NaHCO3 in small quantities (to avoid sudden flow of the solution because of rapid evolution of CO2 gas) the flask until CO2 evolution ceases.

Isolation of Product:

Cool the flask in an ice-water bath for 20-30 min. A brown solid consisting of a mixture of ferrocene and its mono- and diacetylated derivatives will precipitate out from the orange solution. Collect the solid by suction filtration using a Hirsch funnel and wash the brown crystals with a 100-mL portion of water. Dry the product on a clay tile.

Separation of Products by Thin-Layer Chromatography (TLC):

Prepare a concentrated solution of the product prepared above in 2-3 drops of toluene. Using a fine capillary spot a TLC plates (2.5x7.5 cm) with silica absorbent and elute five of them in small bottles as developing chambers 2-3 mL of different trial solvent [such as petroleum ether(bp 60-70 oC), diethyl ether, a mixture of petroleum ether, a mixture of petroleum ether and diethyl ether, ethyl acetate, and 10% ethyl acetate-90% petroleum ether by volume] in each bottle. After the solvent has moved about three quarters of the way up remove the plates, allow to dry in air and put the developed plates in a closed container containing few crystals of iodine. The spots on the plates indicate the components . By inspecting the slides and location of spots, determine which solvent gives the best separation of the products.

In order to determine which spot is due to unreacted ferrocene, make two spots on the sixth TLC plate-one each of the solution of the product mixture and of pure ferrocene in toluene. Place the plate in the developing chamber containing the solvent that gave the best separation of the products. Develop the plate as described before.

Column Chromatographic Separation of the Products:

Prepare a microchromatographic column in a Pasteur pipet by pushing a small wad of cotton down to the tip of the pipet. Fill the pipet nearly to the top with silica gel, which serves as the solid phase. Begin dripping the solvent selected in the TLC analysis through the column, at the rate of one drop per second. Place 40-60 mg of product in the solvent selected and add this dropwise to the column. Continue eluting with the solvent. Do not allow the column to run dry.

The column may exhibit two colored bands because of the two principal components, which are ferrocene and monoacetylated ferrocene. Unreacted ferrocene travels at a faster rate than the other product(s). When the first band begins eluting, place a clean 10-mL Erlenmeyer flask under the column, and fraction in a separate 10-mL Erlenmeyer flask. When separation is complete, evaporate the solvent from each fraction collected (HOOD) with a gentle stream of N2 on a warm sand bath.

Characterization of Products(s):

Obtain IR spectra (KBr pellets) of both components. Compare the spectra with that of pure ferrocene and determine the identity of each fraction.

Questions:

1) Draw the reaction mechanism for the acetylation of ferrocene.

2) What is the geometric orientation of the Cp rings in ferrocene?

Propose an explanation.

3) The second acetylation of ferrocene invariably occurs on the second ring, never on the

same ring as the first acetylation. Explain.

4) Many more cyclopentadienly metal compounds are known than arene metal

compounds. Explain why.

5) Cyclopentadiene can bond in other ways than seen here with iron. From the literature,

describe the bonding in titanocene, a rather controversial subject. (Hint: More than one

form of titanocene is known.) Can cyclopenatdiene bond in other ways as well?

(X). Preparation of Tin Iodide

Sn + 2I2 SnI4 (reflux in carbon tetrachloride)

Procedure:

1) Place 1.0 g of tin metal and 3.2 g of solid iodine (weigh iodine on a watch glass) in a

250-mL. round bottomed flask.

2) Add 50 mL. of carbon tetrachloride.

3) Set up the flask for reflux on a heating mantle with an open water-cooled condenser fit

to the flask.

4) Continue reflux until the iodine completely reacts ( vapor of solvent becomes

colorless).

5) Stop heating and filter the contents of the flask while hot.

6) Leave the clear solution to cool to room temperature, and then cool it in an ice bath.

7) Filter the orange crystals formed.

8) Calculate percentage yield of tin iodide and determine its melting point.

Questions:

1) How can tin iodide be recrystallized?

2) Test the solubility of tin iodide in water. Is the aqueous solution acidic or basic?

3) To a potassium iodide solution in a test tube add a little amount of tin iodide. Shake

the tube and notice any change. Explain this result.

4) What is the structure of tin iodide?

(XI). Geometrical Isomerism

Preparation of potassium cis- and trans- dioxalatodiaquachromate(III)

Octahedral complexes exhibit two types of geometrical isomerism MA4B2 and MA3B3 types. In the cis- and trans- isomers, it is difficult to isolate one isomer from the mixture, since isomerization may take place. However, some complexes form stable isomers, for example, crystallization of chlorobis (triethylstibine) palladium (II) from benzene solution gives only the cis-isomer.

In case of chelates, geometrical isomers are expected to be more stable to isomerization, why ?

In this experiment you are going to prepare the cis-and trans- isomers of dioxalato diaquachromate (III).

A- The cis- isomer:

Precedure:

1) Thoroughly powder 12 g of oxalic acid and 4 g of potassium dichromate separately in

a dry mortar. Mix them as intimately as possible by grinding them together.

2) Moisten a 10 cm evaporating basin with water, and then allow nearly all the water to

drain away. Put the mixture in the center of this basin in a compact heap.

3) Cover with a watch glass, and warm gently. Soon a vigorous reaction sets in, with

evolution of steam and CO2.

4) Extinguish the flame. Without waiting for the mixture to cool, pour over it 20 mL

ethanol. Leave the dish on the gauze to keep it warm.

5) Stir the mixture vigorously with a nickel spatula until the product solidifies. Break up

the lumps with a pestle if necessary.

6) If the product solidifies completely pour off the ethanol then add 20 mL of fresh

ethanol and stir again until the product is entirely crystalline and granular.

7) Wash the purple black solid with ethanol and dry it in air. Weigh the product and

determine the yield.

K2Cr2O7 + 7H2C2O4.2H2O 6CO2 + 17H2O + 2KCr(C2O4)2.2H2O

B- The trans-isomer:

Procedure:

1) Dissolve 12 g of oxalic acid in the minimum amount of boiling water in a 400 mL

beaker.

2) Dissolve 4 g of potassium dichromate in the minimum amount of boiling water, and

add the solution gradually to the solution of oxalic acid. Cover the beaker while the

reaction proceeds.

3) Transfer the liquid to an evaporating basin and leave it to evaporate spontaneously at

room temperature. When the liquid is about one-third of its original bulk, rose-colored

crystals of the trans-isomer will have separated.

4) Filter these off by suction on a sintered glass crucible. Wash them with a little cold

water and then with methanol.

5) Dry in air, weigh the product and determine the yield.

Questions:

1) Which of the two isomers is more stable?

2) Suggest a mechanism for isomerization of one geometrical isomer to the other.

3) Mention two physical methods to differeniate between the two isomers.

4) Record the spectra of the two isomers and notice the differences between these spectra.

(XII). A. Determination of 10Dq for [Ti(H2O)6]+3

B. Construction of part of spectrochemical series

A closer look at the spectra, shows that from quantum mechanics, the d-d transition is forbidden. However, experimentally there is an absorption for [Ti(H2O)6]3+.

In this experiment, the absorption spectra of this complex ion and some other comp-lexes will be determined in order to measure 10Dq and the spectrochemical series of these ligands.

Procedure for Part A:

TiCl3 may be purchased as an aqueous solution. Measure the spectrum of the aqueous soution using a 1.0 cm cell. A 3.0% solution gives a band of suitable intensity. Record the absorbance of this solution over the range 300-700 nm.

Complementary work:

1) Determine the position of maximum absorption and hence calculate the crystal

environment of water ligands.

2) Calculate the C.F.S.E. value for [Ti(H2O)6]+3 in its excited state?

Procedure for Part B:

Carry out preparations on a selection of two complexes besides the ones provided. Obtain the spectra of each complex as you have done in part (A) before. The recommended concentrations are:

CrCl3.6H2O 0.04 M

[Cr(H2O)6](NO3)3.3H2O 0.04 M

Cr(acac)3 0.01 M ( preferably in EtOH or MeOH)

K3[Cr(C2O4)3].3H2O 0.005 M

From the spectra determine the positions of the absorption maxima. Obtain the approximate values of 10 Dq for each species and thus derive a "spectrochemical series" of the ligands involved.

Questions:

1) Why does Cr(acac)3 appear to have only one absorption band in the range studied?

2) What is the structure of CrCl3.6H2O?

3) Why does the spectrum of an aqueous solution of this species vary with time?

Preparation of K3[Cr(C2O4)3].3H2O

Dissolve 6 g oxalic acid dihydrate in 15 mL of water and slowly add 2 g of potassium dichromate. After the reaction subsides, heat the solution to boiling and dissolve in it 2.4 g of potassium oxalate monohydrate. Cool to room temperature and add 2-3 mL of methanol. Cool and filter (sinter). Wash with a mixture of equal parts water and methanol, and finally with methanol alone.. Suck dry. Yield: 4 g.

In your account, include a balanced equation for the preparation and draw the two isomeric forms of the anion.

(XIII). Preparation and conductivity of

[Co(NH3)5Cl]Cl2 and [Co(NH3)4CO3]NO3

The synthesis of [Co(NH3)4CO3]NO3 will be carried out according to the balanced equation:

Co(NH3)2 + NH3(aq) + (NH4)2CO3 + H2O2 [Co(NH3)4CO3]NO3 + NH4NO3 + H2O

The preparation of [Co(NH3)5Cl]2+ is accomplished from the carbonato complex according to the following series of equations:

[Co(NH3)4CO3]+ + 2HCl [Co(NH3)4(OH2)Cl]2+ + CO2 + Cl-

[Co(NH3)4(OH2)Cl]2+ + NH3(aq) [Co(NH3)4(OH2)]3+ + Cl-

[Co(NH3)4(OH2)]3+ + 3HCl [Co(NH3)5Cl]Cl2(s) + H2O + 3H+

The molar conductance, LM , is defined as the conductance of 1 cm3 of solution that contains one mole of solute. The specific conductance, L, is the conductance of 1cm3 of solution. The conductance per mole of solute may be calculated by dividing L by the number of moles present in 1 cm3 of solution:

LM = 1000L/M

where M is the molarity of the solution. Comparison of molar conductances with those of known ionic substances allow one to determine the number of ions present in a given salt. General ranges of LM for 2, 3, 4 , and 5 ion conductaors at 25oC in water solvent are as follows:

number of ionsLM

2118-131

3235-273

4408-435

5~560

Carbonatotetramminecobalt(III) nitrate, [Co(NH3)4CO3]NO3

Procedure:

1) Dissolve 20g (0.21 mole) of (NH4)2CO3 in 60 mL of water and add 60 mL of

concentrated aqueous NH3.

2) While stirring, pour this solution into a solution containing 15g (0.052 mole) of

[Co(OH2)6](NO3)2 in 30 mL of water.

3) Slowly add 8 mL of a 30 percent H2O2 solution.

4) Pour the solution into an evaporating dish and concentrate over a gas burner in a fume

hood to 100 mL. Do not allow the solution to boil. During the evaporation time add, in

small portions, 5g (0.05 mole) of (NH4)2CO3.

5) Filter the hot solution by suction using a glass frit filter and cool the filtrate in an ice

water bath.

6) Under suction, filter off the red product crystals then wash it in the filtration apparatus

first with a few milliliters of water and then with a similar amount of ethanol. Calculate

the yield. (Save a portion of your product for the conductance derermination).

Chloropentamminecobalt(III) chloride, [Co(NH3)5Cl]Cl2

Procedure:

1) Dissolve 5g of [Co(NH3)4CO3]NO3 in 50 mL of water and add concentrated HCl (5 to

10 mL) until all of the CO2 is expelled.

2) Neutralize with concentrated aqueous NH3 and add about 5mL excess.

3) Heat for 20 minutes, again avoiding boiling; [Co(NH3)5(OH2)]3+ is formed. Cool the

solution slightly and add 75 mL of concentrated HCl.

4) Reheat for 20 to 30 min and observe the change in color. Purple-red crystals of the

product separate on cooling to room temperature.

5) Wash the compound several times, by decantation with amounts of ice-cold distilled

water, then filter under vacuum with a glass sintered funnel (medium porosity). Wash

with several milliliters of etnanol. Dry in an oven at 120oC to remove solvent yields

[Co(NH3)5Cl]Cl2. Calculate the yield. Obtain the infrared spectra of the products.

(Save a portion of your product for the conductance derermination).

Electrical conductance of solutions of ionic compounds:

1) Prepare a 0.02 M KCl solution and obtain the cell constant, k, from the resistance

miasured for the solution.

2) Prepare 500 mL of 0.001M aqueous solutions of [Co(NH3)4CO3]NO3 and complexes,

and measure their resistances while fresh. Calculate the molar conductances of the two

Co(III) complexes.

Questions:

1) The conductance of an aqueous solution of [Co(NH3)5Cl]Cl2 changes on standing

overnight. Would you expect it to increase or decrease? Why?

2) In the final isolation of [Co(NH3)4CO3]NO3 and [Co(NH3)5Cl]Cl2, why are the

solids washed with ethanol after having first been washed with water?

3) Why are 500 mL of the 0.001M [Co(NH3)4CO3]NO3 and [Co(NH3)5Cl]Cl2

solutions?

4) How do [Co(NH3)4CO3]NO3 and [Co(NH3)5CO3]NO3 differ structurally? How

would you experimentally distinguish between these two compounds?

(XIV). Kinetics of Aquation of

Chloropentammine Cobalt(III) Chloride

Substitution reactions of octahedral Co(III) complexes have been greatly studied, due to the stability of these complexes. the two extreme mechanisms for such reactions are:

i) SN1 (dissociative)

[N5CoL]n+ [N5Co](n+1)+ [N5CoE]n+

ii) SN2 (associative)

[N5CoL]n+ + E- [N5CoLE](n-1)+ [N5CoE]n+

The possibility of a concerted interchange mechanism tending toward the dissociative side is also possible.

In this experiment, you will first prepare the complex [(NH3)5CoCl]Cl2 and then carry out the displacement reaction of Cl- by water, in an acidic medium. From the kinetics of the reaction, you will see if the reaction mechanism is SN1, SN2 or something else, by establishing the kinetics of the reaction..

Procedure:

Carbonatotetrammine cobalt(III) nitrate, [Co(NH3)4CO3]NO3

(see the previous experiment)

Chloropentammine cobalt(III) chloride, [Co(NH3)4Cl]Cl2

(see the previous experiment)

The Kinetics of Aquation of [Co(NH3)5Cl]2+

In a bath set at 60oC, thermostat for at least 15 min two volumetric flasks, one containing 0.10 M HNO3 and the other 0.30 M HNO3 solution. (conc.HNO3 is about 16 M.). To each solution add sufficient [Co(NH3)5Cl]Cl2 to give a 1.2x10-2 M complex concentration. Shake the flasks until all of the complex has dissolved, and then return the flasks to the thermostat bath. When the solution has again reached 60oC (about 15 min) begin withdrawing samples with a pipette ( use a pipette bulb). At approximately 15-minute intervals withdraw 10 mL samples from each reaction mixture. Measure the visible spectrum of each sample immediately after removal, and recored the time of withdrawal. Take 8 aliquots from each reaction solution.

Either measurthe absorbance of each sample at the set wavelength of 550 nm, or run the spectrum over the same piece of chart paper for the samples of each solution (350-650 nm).

1) Record the absorbance measurements taken during the runs.

2) Calculate the absorbance at infinite time, A, given that the extinction coefficient of

[Co(NH3)5(OH2)]3+ is 21.0 cm-1 M-1 at 550 nm, then calculate the rate constant

from your measurements. The first order rate constant is equal to the slope of the plot

of ln(A - A) versus time where A is the absorbance of solution at any time, t. Justify

this plot as a method of obtaining the first order rate constant.

3) Plot the data to determine the first-order rate constants from absorbance and time data.

Rate = kobs[complex]

4) Determine the rate constant dependence on hydrogen ion concentration.

kobs = k[H+]n

5) Propose a mechanism for the kinetics of aquation.

(XV). Stability Constants of Ni(glycinate)n(2-n)+

The acidic character of glycine +NH3CH2COO- (HA) allows stability constants for Ni+2 complexation of the glycinate ion, NH2CH2COO- (A-), to be evaluated. The acid dissociation constant of glycine must first be determined.

H

+

+ A

-

HA

Ka = [H+][A-]/[HA](1)

This is done by measuring [H+] in a solution of known [HA].

By definitionaH+ = g

[H+],

where aH+ and g are the activity and the mean aH+ activity coefficient of the hydrogen ion respectively. Rearranging and substituting pH = - log aH+ gives

log[H+] = - pH - log g(2)

and log[OH-] = pH - pKw + log g(3)

Hence it is possible to evaluate [H+] and [OH-] for use in calculating the concentration equilibrium constants from pH measurements.

Since [H+] + [Na+] = [OH-] + [A-]

and Atot = [HA] + [A+]

where Atot is the total glycine conentration, then for glycine:

pKa = -log[H+] + log{( Atot - ([Na+] + [H+] - [OH-])/([Na+] + [H+] - [OH-])} (4)

All the quantities on the right-hand side of this expression are either known or may be calculated from equations (2) and (3), and the pKa of glycine may be evaluated.

The method of determining the stepwise stability constants K1, K2, and K3 for the reaction of Ni+2 with A- is as follows:

Experimentally, a solution containing 1 mmole of Ni+2 (as NiCl2.6H2O) and 1 mmole of H+ (as HNO3) will be titrated with a solution of sodium glycinate prepared by neutralizing glycine with NaOH. This will give a solution contianing an equilibrium mixture of H+, OH-, Na+, HA, A-, Ni+2, NiA+, NiA2 and NaA3-. From pH measurements and a knowledge of the quantities Ni+2, H+, HA and NaOH originally added it is possible to calculate the stability constants.

To facilitate the determination of the K's, a function,

n

, is defined as the average number of lignad molecules bound per metal ion.

Here,

n

= moles of bound A- / total moles of Ni+2

= ([NiA+] + 2[NiA2] + [NiA3-])/([Ni+2] + [NiA+] + [NiA2] + [NiA3-]) (5)

Experimentally,

n

may be expressed in terms of the total glycine concentration (Atot), the conentration of HA and A- , and the total Ni+2 concentration (Mtot)

n

= (Atot - [HA] - [A-])/Mtot(6)

since bound H+ = (added H+ from HNO3) - (H+ from dissociation of water) - free H+

so [HA] = CH - [OH-] - [H+](7)

substitution in (1) gives

[A-] = {Ka (CH - [OH-] - [H+])}/[H+](8)

substitution of (7) and (8) into (6) gives

n

= {Atot - (1 + Ka/[H+])(CH - [OH-] - [H+])}/Mtot(9)

A typical plot of

n

versus [A-] is linear. From such plots it is possible to estimate the values of K1, K2, and K3.

Since

Ni

+2

+ A

-

NiA

+

K1 = [NiA+]/[Ni+2][A-]

when[NiA] = [Ni+2] thenK1 =1/[A-]

i.e., when

n

= 1/2, the value of K1 can be estimated from the plot.

Likewise, at

n

= 3/2 and

n

= 5/2 the values of [A-] will allow the estimation of K2 and K3. In general, Kn can be estimated from the expression

Kn = [1/[A-]]

n

= n-0.5

A more precise graphical method for evaluating K1, K2, and K3 from

n

and [A-] was developed by F.J.C.Rossotti and H.S.Rosotti:

Using the overall stability constants, b, equation (5) is converted to equation (10)

n

= (b1[A-] + 2b2[A-]2 + 3b3[A-]3) / (1 + b1[A-] + b2[A-]2 + b3[A-]3) (10)

where b1 = K1, b2 = K1K2 =, and b3 = K1K2K3

Rearrangement gives:

n

/(1-

n

)[A-] = b1 + (2-

n

)[A-]b2/(1-

n

) + (3-

n

)[A-]2b3/(1-

n

)(11)

A plot of

n

/(1-

n

)[A-] vs. (2-

n

)[A-]/(1-

n

) tends to a straight line at low [A-] of intercept b1 and slope b2.

Using a value of b1 obtained as the intercept, eqn.(11) is divided by (2-

n

)[A-]/(1-

n

) and rearranged to give

{

n

-(1-

n

)b1[A-]}/(2-

n

)[A-]2 = b2 + (3-

n

)[A-]b3/(2-

n

)(12)

The intercept of a plot of {

n

-(1-

n

)b1[A-]}/(2-

n

)[A-]2 vs.(3-

n

)[A-]/(2-

n

) gives b2 and the slope is b3.

From the relatively precise values of b1, b2 and b3 , the stability constants K1, K2, and K3 may be calculated. Because of the method of data treatment, the precision of the constants decrease in the order of K1 > K2 > K3.

Procedure:

The pH meter must first be standardized. Insert the electrodes into a pH 7 buffer and adjust the calibration knob until the meter reads the same as the pH of the buffer. Then, using buffers of pH 4 and 10, check the meter reading. If the meter reading differs by more than 0.1 pH unit from the buffer, recheck the reading of the buffer of pH 7. If it is correct but incorrect at pH 4 and pH 10, the meter reading must be corrected by the slope adjustment or the temperature compensator. Otherwise, it should be accomplished by determining a meter correction constant,Cmeter , from the linear expression:

pH= 7.00 + (pHmeter -7.00)Cmeter (13)

1) Standardize the pH meter in a pH 7 buffer using calibration knob. Using buffers of

pH 4 and 10 check the meter reading, which should agree within 0.1 pH unit. If no

2) Prepare the following solutions:

20 mL 0.4M glycine,HA

500 mL 0.2M KNO3

100mL 0.10 M HNO3

100mL 0.5M NaOH

3) Standardize the NaOH solution by titrating a solution of the primary standard,

potassium acid phthalate, using phenolphthalein indicator. Use the standardized NaOH

solution to standardize the HNO3 solution. Store all solutions in tightly stoppered

bottles.

4) To determine pKa of glycine:

Add a magnetic stirring bar and introduce the electrodes of the standardized pH meter

into a 400 mL beaker containing 90 mL of H2O, 100 mL of 0.2 M KNO3 and 10 mL of

0.4M glycine. If possible thermostat the beaker at a constant temperature. Rinse and fill

a 10 mL burette with the standardized 0.5M NaOH. With stirring add about 15mL

aliquots of NaOH in titrating from 0 to 100 percent of the glycine. After each addition

record the amount of NaOH added and the pH. Using points in the 20 to 80 percent

titration range, calculate the pKa of glycine from equations 20,9,10,.11. Obtain an

average Ka value.

5) Collecting data for calculating K1,K2, and K3

Prepare 20 mL of 0.4M sodium glycinte solution by neutralizing weighed solid glycine

with a calculated amount of standardized 0.5M NaOH and diluting with water to a total

volume of 20 mL. Rinse and fill the burette with this solution and titrate with stirring

a solution of 100 mL 0.2 M KNO3, 1 mmole NiCl2.6H2O (weigh this amount ), 10 mL

0.1M HNO3 and 90 mL of H2O in a 400 mL beaker. After each 0.2 mL aliquot addition,

record the amount of glycinate solution added and the pH of the solution. Add the full

10 mL. From the titration data calculate [A-] from eqn.(15) and

n

from eqn.(16) at

each of the 50 points. Note that [H+] and [OH-] may be negligible in some cases. In

these calculations, be sure to correct meter reading (eqn.20) if necessary and to convert

pH to [H+] using eqns.(9), (10). Make a plot of

n

versus log[A-]. From the values of [A-] at

n

= 1/2, 3/2, 5/2 calculate the approximate values of K1, K2, and K3..

Now calculate

n

/(1- n ){A} and (2-

n

)[A-]/(1-

n

) for those points with values of

n

between 0.8 and 0.2, then plot these quantities to obtain an accurate value of b1 from the intercept. Then evaluate.

{

n

-(1-

n

)b1[A-]}/(2-

n

)[A-]2 and (3-

n

)[A-]/(2-

n

)

For the points whose

n

values range from 1.1 to 1.7, then from the plot of these quantities determine b2 (intercept) and b3 (slope). Finally K1, K2, and K3 fromb1, b2, and b3 and compare these values with the approximate ones that you evaluated earlier by using the half-n method

Questions:

1 ) You wish to determine the formation constant K1, K2, and K3 for the coordination of ethylenediamine, NH2CH2CH2NH2 (en), to Ni+2 to form Ni(en)+2, Ni(en)22+, and Ni(en)32+.

a) How would you determine the first and second pKa values of en?

b) Define these two pKa values. Which of the two will be larger, and why?

c) Briefly explain what titrations you would carry out in order to determine the stability constants K1, K2, and K3. would the procedure differ from that used in this experiment?

d) Which value would you expect to be larger, K1 or K2?

2) Consider the coordination of Cu2+ by the tridentate iminodiacetate ligand,

HN(CH2CO2-)2, labeled HA-, is added to an aqueous solution of Cu2+, the following

equilibrium is rapidly established:

Cu2+ + HA CuA + H+

If a solution is prepared by combining 100 mL of 0.02 M Cu2+ and 100 mL of 0.02 M HA- is adjusted it pH 1.60 with NaOH, what are the concentrations of Cu2+, HA-, and CuA in the solution? ( The pKa for HA- is 4.70 x 10-10, and K1 for Cu2+ + A2- = CuA is 4.26 x1010 ). Draw the probable structure of the CuA complex.

Table of Contents

Experiment Page

I.Synthesis of Cr(acac)31

II.Recrystallization and melting point determination of Cr(acac)32

III.Volumetric determination of %Cr in Cr(acac)33

IV.Synthesis of Cr(acacNO2)35

V.Synthesis of Co(acac)36

VI.Synthesis of Copper(II) Acetate Monohydrate7

VII.Preparation of Sodium Tetrathionate10

VIII.Preparation of Copper-Glycine Complexes12

IX.Acylation of Ferrocene14

X.Preparation of Tin Iodide16

XI.Geometrical Isomerism17

XII.A) Determination of 10 Dq for [Ti(H2O6]3+

B) Construction of part of the spectrochemical series19

XIIIPreparation and conductance of [Co(NH3)5Cl]Cl2 and

[Co(NH3)5CO3]NO321

XIV.Kinetics of aquation of [Co(NH3)5Cl]Cl224

XV.Stability constants of {[Ni(glycinate)]2-n}26