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Inorganic Chemistry Exp

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INORGANIC CHEMISTRY EXPERIMENTExperiment 1 : PURIFICATION OF SODIUM CHLORIDE Experiment 2 : DOUBLE SALTS Experiment 3 : GROWING CRYSTAL FROM AQUEOUS SOLUTION Experiment 4 : REACTION OF Cr, Fe, Co, Ni, AND Cu ION Experiment 5 : COMPARISION WITH Cu(I) AND Cu(II) COMPOUND Experiment 6 : Preparation of [Co(NH 3 ) 6 ]Cl 3 and [Co(en) 3 ]Cl 3 Experiment 7 : PREPARATION OF [Co(NH 3 ) 4 CO 3 ]NO 3&

[Co(NH 3 ) 5 Cl]Cl

Experiment 8 : Geometrical isomerism ( cis and trans-[CoCl 2 (en) 2 ]Cl 2 ) Experiment 9 : LINKAGE ISOMERISIM ( Nitro and Nitrito Covalt(III) Complex) Experiment 10 : CYCLE OF COPPER REACTION Experiment 11: Bioinorganic Chemistry: Synthesis and Study of an OxygenCarrying Cobalt Complex Which Models Hemoglobin. Experiment 12 : The Synthesis and Characterization of YBa 2 Cu 3 O x - the "1-2-3" Superconductor.

IONIC SALTS Experiment 1. PURIFICATION OF SODIUM CHLORIDE Subject: Pure Sodium Chloride from Common Salt materials required dropping funnel, round buttom flask, Erlenmeyer flask, funnel, beaker, glass tube, rubber tube common salt(NaCl), BaCl 2 , (NH 4 ) 2 C 2 O 4 , dye (paranitrobenzen-azo-resorcinol), NaOH,c-H 2 SO 4 ,c-HCl procedure Dissolve 50 grams of common salt in 150 ml of hot water. Cool, filter, and test small portions of the solution qualitatively for sulfate, calcium, and magnesium. To test for sulfate, add barium chloride and dilute hydrochloric acid; for calcium,add ammonium oxalate; for magnesium, add a few drops of a dilute solution of the dye paranitrobenzene-azo-resorcinol followed by 2N sodium hydroxide.In the last test, magnesium gives magnesium hydroxide colored blue by the dye; the blue is distinct from the purple color the dye itself gives with sodium hydroxide.

Fig.1. Pure sodium chloride from common salt Place the remaining solution in an Erlenmeyer flask and pass over it a slow stream of hydrogen chloride gas, generated by dropping concentrated sulfuric acid into concentrated hydrochloric acid. To prevent the hydrogen chloride gas from passing out into the room, it is absorbed in sodium hydroxide solution by use of the funnel arrangement shown in Fig.1. Note: (a) The gas is passed over the salt solution and not

through it. (b) The funnel dips only a millimeter or so below the surface of the sodium hydroxide. The reasons for these arrangements should be clear. Pure sodium chloride precipitates from the solution, the impurities remaining in solution. Note the shape of the crystals as they form. When a substantial precipitate has formed, filter it on the Buchner funnel, rinse with a very little cold concentrated hydrochloric acid, then dry in an evaporating dish on the steam bath. Use a glass spatula, not steel, for the manipulations. Dissolve a gram or so of product in 15 ml of distilled water and test this solution as before for sulfate, calcium, and magnesium. It is not necessary to report the yield in this preparation, but turn in your product.

Questions 1. Why does sodium chloride precipitate, whereas the impurities do not? 2. Why not purify the sodium chloride by simply recrystallizing it from hot water? 3. Why does concentrated sulfuric acid displace hydrogen chloride from its solution in water? Account for the lack of any considerable evolution of heat when the sulfuric acid is added to the hydrochloric acidsolution.

1. IONIC SALTS Experiment 2. DOUBLE SALTS Double salts are formed when two simple salts crystallize together in definite, simple molecular proportions. They have their own crystal form, which need not be the same as that of either of their component salts. They are a phenomenon of the solid state; in solution they are decomposed completely, or nearly so, into the ions of their component salts. In this respect double salts are distinguished from complex salts, which give complex ions of their own in solution. Double salts are extremely numerous. Two examples will be prepared. subject A: CuSO 4 (NH 4 ) 2 SO 4 .6H 2 O. materials requireds Erlenmeyer flask,beaker,Buchner funnel,filter paper, CuSO 4 . 5H 2 O,(NH 4 ) 2 SO 4 procedure Dissolve 0.1 mole of copper sulfate pentahydrate and 0.1 mole of ammonnium sulfate in 10 ml of hot water. Cool slowly. When the solution is cold, filter off the crystals on a Buchner funnel and dry then on filter paper in the air. Weigh, and record the yield. Examine them to see if they are homogeneous. Compare their appearance and behavior in solution with that of the complex salt cupric tetrammine sulfate,[Cu(NH 3 ) 4 ]SO 4 .H 2 O, which may be obtained from the side shelf or prepared as a separate exercise. This salt will separate in large well-formed crystals if a cold saturated solution is left to evaporate slowly in the air. The crystals belong to the monoclinic system.

Fig.2. Double Salts

Subject B: CuCl 2 .KCl.2H 2 O. materials required Erlenmeyer flask,beaker,Buchner funnel,filter paper, CuCl 2 , KCl procedure Dissolve 0.25 mole of copric chloride hydrate and 0.30 mole of potassium chloride in 40 ml of hot water, and cool slowly. When the solution is cold, filter it on a Buchner funnel and dry the crystals as well as possible with filter paper. The crystals are efflorescent and should not stand in the air long; otherwise they lose water and become brown. When dry but completely hydrated, they are green-blue in color. They are dried without losing their water of hydration by putting them in a desiccator over wet calcium nitrate, which gives a constant relative humidity of 51 % at 25oC. Dissolve about 5 grams of the salt in just enough warm water (about 50oC) to bring about solution; then cool until the first crystals appear. Filter these crystals under sucti.on, using a very small filter(a filter made by putting a small plug of glass wool in the neck of a glass filter funnel, and then covering the glass wool with a slight coating of asbestos fibers, is satisfactory). Alternatively, pour off the mother liquor, quickly transfer the mass of moist crystals to a small centrifuge tube packed tube for balancing. Note the appearance of the crystals. What are they? Can this double salt be purified by direct recrystallization? Prepare 10 to 15 ml of dilute solutions of each of these double salts and a similar solution of cupric sulfate. Compare their colors and explain. Questions 1. Make a list of half a dozen double salts that are important in industry or in laboratory practice. 2. How is potassium chloride extracted from carnallite? How would you make a sample of carnallite in the laboratory?

1. IONIC SALTS Experiment 3. GROWING CRYSTAL FROM AQUEOUS SOLUTION There are two general methods for growing large single crystals from aqueous solutions. In one method, crystal growth occurs ad a saturated solution is gradually cooled to a temperature at which the solution is appreciably supersaturated. In the other case, crystal growth occurs as a saturated solution is allowed to gradually evaporate at a constant temperature at which crystal growth is to occur (usually near room temperature) and to prepare some "seed crystal", one of which will be suspended by a thread in the saturated solution. A saturated solution may be conveniently prepared by the following procedure. In about 500mL of water at 50, dissolve somewhat more of the salt whose crystals are to be grown than will dissolve at the expected growing temperature. While stirring vigorously, cool the solution to the expected growing temperature (to not cool below this temperature). If crystallization does not take place, add a small crystal to induce it. Stir the suspension of crystals for about 15min and than let the solution stand in contact with the crystals in a covered beaker or stoppered flask in the crystal-growing room for a day or longer. Finally decent the solution from the precipitated crystals. These crystals may be spread out on a piece of filter paper to dry, and among them may be found a suitable seed crystal. A seed crystal must be al single crystal, and so that it may be easily suspended by a thread, it should be least 3mm long. Smaller crystals are not only difficult ot attach to the thread. Save all good seed crystals, for your initial attempt at crystal growing may not be successful.

Crystal Growth by Cooling A solution, just saturated at the temperature of the crystal-growing room, is heated to about 15 above this temperature; a small additional amount of the salt is dissolve, and the solution is filtered. The solution (which would now be supersaturated at the growing temperature) is carefully poured into a clean 600mL beaker. When the temperature is about 3 above growing temperature, the seed crystal is suspended in the middle of the solution by a thin piece of sewing thread (a monofilament thread, such as nylon, is best) the upper end f which is attached to a piece of wood that completely covers the beaker (see fig. 3). The thread is most conveniently attached to the seed by means of a slip knot. The free end of the thread should be cut off as close as possible to the kont. The beaker should be allowed to stand in a room in which the temperature fluctuates less than 2 throughout the entire day. To avoid rapid temperature fluctuations, the beaker may be covered with several cardboard boxes or a alrge crock. After an hour or two, examine the beaker to see whether or not the seed crystal has dissolved. If it has dissolved, it will be necessary to begin again, using a solution containing a little more salt. A seed crystal generally has several other tiny crystals adhering to it. Thus, if the seed crystal is placed in an undersaturated solution, these tiny crystals will dissolve away, leaving only one crystal. Of course, one hopes that the solution cools rapidly enough so that the solution becomes supersaturated. The crystal should grow to a good size in about 3 to 6 days. Remove the fully grown crystal from the solution, and carefully dry it with filter paper or absorbent tissue. Crystal Growth by Evaporation A solution, saturated at the growing temperature, is heated to about 10 above this temperature and carefully filtered into a clean 600mL beaker. When the temperature is 1 to 2 above growing temperature, introduce the seed crystal suspended by a lass support. A suspension of this type is pictured in Fig 4. The seed crystal is hung by a monofilament thread from the gallows-shaped support, the top of which must always be below the surface of the solution. Cover the beaker with a cloth, and hold it in place with a string or rubber band. The beaker should stand in a room in which the temperature fluctuates less than 5 throughout the entire day. The rate of crystal growth depends on the rate at which water evaporates from the solution. When the crystal has reached a satisfactory size, or when the top of the suspension is about to protrude through the surface of the solution, remove the crystal and dry it with filter paper or absorbent tissue.

2. THE ELEMENTS OF THE TRANSITION SERIES Experiment 4. REACTION OF Cr, Fe, Co, Ni, AND Cu ION The transition elements have partly-filled d or f electronic shells. Elements at the end of the transition series may have filled d or f shells in the atomic state but the same shells may be partly filled in well characterized oxidation states. For example, Cu, [Ar]3d104s1 shows the transition element in the +2 state [Ar]3d9 and it is convinient to regard it as a transition element. In the same way, elements at the beginning of each transition series often have unfilled d or f shells in common oxidation states and partlyfilled shells in the atomic state. Scandium , [Ar]3d14s2 exhibits only the +3 oxidation state with the configuration [Ar]3d0. there are three main or "d"transition series corresponding to the filling of the 3d, 4d and 5d shells. The filling of the 4f and 5f shells gives the lanthanide and actinide series of elements. The first transition series compriese the elements scandium to copper. These are listed in Table 1 together with the eletronic configurations of the atoms and the common oxidation states. The transition elements show characteristic properties that, considered together, distinguish them from the main group elements. (a) They are hard, dense metals with high melting and boiling points. They are good conductors of heat and electricity and some are among the most familiar metals of commerce and everyday life. (b) Many alloys with commercially useful properties are formed between these elements. (c) Small, non metallic elements, such as carbon, boron and nitrogen, form interstitial compounds with the transition metals. These compounds are non- stoichiometric and show metallic properties in that they are hard and high melting. (d) The metals sre eletropositive and reactive although these propties decrease with increasing atomic number in each series. Many are sufficiently electropositive to dissolve in non-oxidizing mineral acids. The insolubility or "passivity" of some of these elements in acid is normally explained by the formation of a thin coherent oxide film on the surface of the metal. (e) The metals each exhibit several oxidation states (with a few exception), see Table 1. The compounds, complexes and ions are colored in one or more oxidation states and absorb in the ultra-violet or visible region of the spectrum. (f) Trantion metal cations like main group cations readily form complexes with a large number of neutral and anionic ligads.

(g) The ions of these metals in their different oxidation states often contain unpaired electrons and are paramagnetic. Table 1. Electronic configuration and oxidation states of the elements in the first transition series.

Comparisions with main-group elements Since the formal maximum oxidation states of the corresponding main group and transitional elements are the same, we may expect some resemblances between them interms of chemical and are readly hydrolysed in water. The elements also form hexahalo complexes, e.g.[SnCl 6 ]2- ,[TiCl 6 ]2-. Vanadium(V) shows few resemblences to phosphorus(V) , although the tetrahedral oxychlorides VOCl 3 and POCl 3 are both liquids and are readly hydrolysed in water . Both chromium(VI) and sulphur(VI) form strongly acidic oxides CrO 3 and SO 3 , and the covalent oxyhalides CrO 2 Cl 2 and SO 2 Cl 2 are easily hydrolysed. Magness(VII) shows some formal resemblances to the halogens in the oxo-compounds Mn 2 O 7 , and MnO 4 - . Both Cl 2 O 7 and the corresponding manganese compound are powerful oxidizing agents and they are of low stability. The tetrahedral oxyanions ClO 4 - and MnO 4 - are strong oxidizing agents. It must be emphasized, however, that the diffrences between the transitional and main group elements greatly outweigh the similarities mentioned. In particular, the oxidation states of the transition metals that have d eletrons have nocounterpart in the main-group elements. The reactions of the elements The reactions described in this section have been selected to illustrate important features of the chemistry of these elements. Carry out the reactions on the small scale indicated. The important oxidation states of each metal are introduced, and the ease of oxidation

and reduction of these states is damonstrated. The reactions of the cations, with, for example, acids and bases may be compared across the transition series. It is useful to remember such comparisions so that the chemistry of each element if seen in the context of the series. When a metal ion gives a characteristic reaction then this is included. The student may wish to carry out other reactions of the cations. discuss these changes with the demonstrator before beginning the experimental work. A. Chromium(Cr)

Chromium, with the electronic configuration [Ar]d54s1, has a highest oxition state of +6. This corresponds to the total number of 3d and 4s electrons, as for titanium and vanadium. In the +6 state chromium is a powerful oxidizing agent and is reduced easily to chromium +3, the most stable state. The intermediate states +5 and +4 are poorly characterrized and have no solution chemistry as they disproportionate readily to chromium(VI) and chromium(III). divalent chromium is strongly reducing and is oxidized to chromium(III). The electrode potential of the metal indicates that it is quite active and should dissolve in dilute mineral acids. Cr3+ + 3e- = Cr Eo = -0.74V

While it is soluble in non-oxidizing acids it resists attack by nitric acid and other oxidizing acids. This passive character means that chromium is widely used in corrosion-resistant alloys and plated films. Chromium(VI) is stabilized by highly electronegative ligands and exists only as oxy compounds. All of its compounds are powerful oxidizing agents. dichromate finds application in analysis, especially in acid solution,as indicated by the potential: Cr 2 O 7 2- + 14H+ + 6e- = 2Cr3+ + 7H 2 O Eo = 1.33V

The oxide, CrO 3 , dissolves in water to give strongly acid solutions. In basic solution,chromium(VI) exists as the yellow, tetrahedral chromate anion. When acid is added to chromate solutions the colour deepens to orange as the binuclear dichromate ion is formed. In the stable trivalent state, chromium forms an amphoteric oxide,Cr 2 O 3. It dissolves in mineral acids to give the hexa-aqua purple cation [Cr(H 2 O) 6 ]3+ and in concentrated alkalis as chromites ,possibly [Cr(OH) 6 ]3-.

chromium(III) forms a very large number of octahedral 6 coordinated complexes which are characterized by their kinetic inertness. Even the co-ordinated water molecules in the ion [Cr(H 2 O) 6 ]3+ exchange quite slowly with the solvent. cationic complexes often contain ammonia, water, or amine ligands, for example the pale green pentaaquachloro ion. Chromium(II) is a basic and strongly reducing state as indicated by the eletrode potential: Cr3+ + e- = Cr2+ Eo = -0.41V

The bright-blue aqueous solutions contain the hydrated octahedral cation [Cr(H 2 O 6 ]2+. The solutions are rapidly oxidized by oxygen, and in the absence of air they slowly attack water and liberate hydrogen. Reactions of chromium These experiments show some of the reactions of chromium in the more important oxidation states. The oxidative stability of these states is indicated. For reactions 1 and 2 use a 5% solution of chromium potassium sulphate. 1. Slowly add 4 mol dm-3 sodium hydroxide to the solution of chrome alum (2 cm3) until no further reaction takes place. Explain the observed reactions. 2. Slowly add 4 mol dm-3 aqueous ammonia to the chrome alum solution (2cm3) until it is in large excess. Compare the results with those obtained in the previous experiment. For the reactions 3 and 4 use a 5% solution of potassium dichromate. 3. Stir the potassium dichromate solution (2cm3) with 4 mol dm-3 sodium hydroxide (2 drops). Add 4 mol dm-3 sulphuric acid (5 drops). Explain the colour changes. 4. Add the potassium dichromate solution (1 cm3) to a 5% solution of ferrous ammonium sulphate in 4 mol dm-3 sulphuric acid (2cm3). Outline the analytical implications of this reaction. B. Iron(Fe)

This element has the outer electronic configuration 3d64s2. The highest oxidation state is +6 but this is extremely rare. The most important oxidation states are iron(II) and iron(III). Iron(II) is a good reducing agent as shown by the oxidation potential:

Fe3+ + e- = Fe2+

Eo = +0.77V

THe double salt FeSO 4 (NH 4 ) 2 SO 4 .6H 2 O (Mohr's salt) is used in volumetric analysis for titrations with dichromate, permanganate and cerium(IV) solutions. In aqueous solution the octahedral ion [Fe(H 2 O) 6 ]2+ is present. It is blue-green. Molecular oxygen will oxidize iron(II) to iron(III) in both acidic and basic solutioin,as shown by the electroed potentials for these systems: Fe2+ + O 2 + 2H+ = 2Fe3+ + H 2 O Eo = +0.46V

Iron(II) forms complexes of witch the most important is the porphyrin complex haem, which exists associated with globular protein in haemoglobin. In iron(II) complexes quite strong ligand fields are requied to cause the electrons to pair. Nearly all iron(II) complexes are high spin but [Fe(CN) 6 ]4- and [Fe(CNPh) 6 ]2+ are low spin. Ferric salts are less ionized than the corresponding iron(II) salts, and ferric oxide is amphoteric. In aqueous solution the hydrated ferric ion [Fe(H 2 O) 6 ]3+ undergoes hydrolysis, and the purple hexaqua ion exist only when the pH is near zero. Iron (III) forms many complexes most of which are octahedral, and it also forms a few tetrahedral complexes,e.g.[FeCl 4 ]-. Iron(III) has little affinity for amine ligands but it readly forms complexes with ligands that coordinate via oxygen. For example, the oxalate ligand forms the octahedral anion [Fe(C 2 O 4 ) 3 ]3-. The reactions of iron These test illustrate the reactions and relative stabilities of Fe(II) and Fe(III). Record your observations, and explain the reactions that occur together with the equations where possible. Ferrous, Fe(II) ion Use a freshly-prepared solution of 2 mol dm-3 ferrous ammonium sulphate for the tests.(Prepare 50 cm3 of your own). 1. To the iron(II) solution (5 cm3) add an excess of 2mol dm-3 sodium hydroxide and allow the solution to stand for 10 min. To another portion (5 cm3) of the test solution add 2 mol dm-3 sodium hydroxide (5 cm3) followed by excess hydrochloric acid. 2. Add excess aqueous 2 mol dm-3 ammonia to the ferrous solution. 3. To the iron(II) solution (5 cm3) add an excess of sodium carbonate solution and

compare this with the action of the aqueous sodium hydroxide. 4. To the ferrous solution (5 cm3) add aqueous 2 mol dm-3 potassium thiocyanate. Ferric, Fe(III) ion Use a freshly-prepared solution of 2 mol dm-3 ferric sulphate for the tests. (Prepare 50 cm3 of your own). 1. Repeat the test 1,2 and 3 above,using the solution of ferric sulphate instead of the ferrous ammonium sulphate solution. 2. To the ferric solution (5 cm3) add 2 mol dm-3 potassium thiocyanate(5 cm3) followed by ammonium fluoride (0.5g). C. Cobalt (Co)

This element has the outer electronic configuration 3d74s2 and its highest significant oxidation states is +4. This reflects the trend towards decreased stability of the very high oxidation states on moving across the transition metal series. Cobalt(I) forms some complexes, most with p-bonding ligands. The chemistry of cobalt(I) is better characterized than any other unipositive oxidation state of the first transition series except copper. The two most important oxidation states are cobalt(II) [Ar]3d7 and cobalt(III) [Ar]3d6,and in an aqueous solution containing no complexing agent cobalt(III) is easily reduced to cobalt(II). however, cobalt(III) is more stable in the presence of a complexing agent such as ammonia as shown by the electrode potentials. [Co(H 2 O) 6 ]3+ + e- = [Co(H 2 O) 6 ]2+ [Co(NH 3 ) 6 ]3+ + e- = [Co(NH 3 ) 6 ]2+ Eo = 1.85V Eo = 0.1V

Cobalt(II) forms both octahedral and tetrahedral complexes but they are labile and they have a strong tendency to be oxidized by molecular oxygen. The complexes are usually prepared in an inert atmosphere. In aqueous solution the cobalt(II) ion is pale pink because the absorption is weak and occurs in the blue region of the visible. Tetrahedral cobalt(II) complexes are often highly coloured owing to their lower order of symmetry relative to the octahedral complexes. The spectrum of [CoCl 4 ]2- shows a large absorption in the visible part of the spectrum, which accounts for its deep-blue colour.

Cobalt(III) salts are difficult to prepare because the ion is a strong oxidizing agent and the chemistry of this oxidation state is largely that of coordination compounds. Cobalt(III) usually forms octahedral complexes and it has a strong affinity for nitrogen donors such as ammonia, amines(e.g. ethylenediamine), nitro groups and nitrogen bonded -SCN groups as well as water molecules and halide ions. Reaction of cobalt these tests illustrate the reactions of cobalt(II). Record your observation and give the reactions that occur together with the equations where possible. Use the solution of 1 mol dm-3 cobalt(II) chroride or nitrate for the following tests. 1. To the cobalt(II) solution (5 cm3) add slowly an aqueous solution of 60% sodium hydroxide (8 cm3). 2. Add 2 mol dm-3 sodium hydroxide (5cm3) to the cobalt(II) solution(5 cm3). D. Nickel ( Ni )

This element has the outer electronic configuration 3d84s2 and its highest important oxidation states is +4. With nickel the trend towards the decreased stability of the higher oxidation states continues and the most common oxidation dtate in aqueous solution, where it forms the green hexaaqua nickel ion. This ion occurs in some hydrated nikel(II) salts, for example Ni(NO 3 ) 2 .6H 2 O, NiSO 4 .6H 2 O and Ni(ClO 4 ) 2 .6H 2 O. note that NiCl 2 .6H 2 O contains trans-NiCl 2 (H 2 O) 4 units. Nickel(II) readly forms complexes, the main structral types being octahedral, tetra hedral and squire planar. Amines form blue octahedral complexes, for example [Ni(NH 3 ) 6 ]2+ and [Ni(1,2-diaminoethanne) 3 ]2+. The visible spectra of these complexese have three bands with low molar absorbances 1-10 and these are assigned to the three spin-allowed transitions. The tetrahedral complexes of nickel(II) are usually intensely blue, and in there visible spectra they have molar absorbance of approximately 200. These values are in contrast to the low values found for the octahedral complexes. Typical tetrahedral complexes are [NiCl 4 ]2- and NiCl 2 (PPh 3 ) 2 . Square planar nickel(II) complexes are diamagnetic and often brown, red or yellow in colour. Typical nickel(II) square planar complexes are [Ni(CN) 4 ]2-, NiBr 2 (PEt 3 ) 2 and bis(dimethylglyoximato)nickel(II).

The reactions of nickel These tests illustrate the reactions of nickel(II) and should be carried out on a solution of 0.2 mol dm-3 nickel(II) nitrate or nickel(II) chloride. For these tests 1-3 record all observations, interpret the results and give equations where possible. 1. To the nickel(II) solution (3 cm3) add excess 2 mol dm-3 sodium hydroxide (4cm3) and then concentrated aqueous ammonia(5 cm3). 2. Add 2 mol dm-3 potassium thiocyanate solution(3 cm3) to the nickel(II) solution (3 cm3) followed by a few drops of pyridine. 3. Add am alcoholic solution of dimethylglyoxime to the nickel(II) solution, which has been acidified with 2 mol dm-3 hydrochroric acid (2 drops). Systemetically change the pH of the solution and note what happens. E. Copper ( Cu )

These element has the outer electronic configuration 3d104s1 and the single s electron can be removed to give copper(I). This oxidation state is somtimes compared with the alkali metal ions but there is little similarity and copper(I) is best condered as a typical transition metal ion. Cuprous compounds sre diamagnetic and colourless except when colour results from the anion or from charge transfer bands. The stability of copper(I) in aqueous solution is low, as shown by its oxidation potential, and it disproportionates to give copper(II) and copper metal: however many cuprous salts are insoluble in water. Cu+ + e- = Cu Eo = 0.52V Cu2+ + e- = Cu+ Eo = 0.15V 2Cu+ = Cu + Cu2+ Eo = 0.37V If potassium iodide is added to a copper(II) solution, cupric iodide is formed but this rapidly decomposese to give a precipitate of cuprous iodide and iodine. Most cupric salts dissolve in water to give the blue hexaaqua ion [Cu(H 2 O) 6 ]2+. If ammonia added, up to four water molecules can be replaced stepwise to give the complex [Cu(NH 3 ) 4 (H 2 O) 2 ]2+. However it is difficult to replace the last two water molecules because of the Jahn-Teller effect. There are many amine complexes of copper(II) and they are all more intensely coloured than the hexa-aqua ion. The stronger ligand field of the amine cases the single absorption band in the visible spectrum of the

hexa-aqua to move to shorter wavelengths as the water molecules are replaced by the ligand. Copper(II) forms halide ion complexes of the type [CuX 4 ]2-(X = Cl,Br) which have distorted tetrahedral structures. For example, when lithium bromide or hydrobromic acid is added to cupric bromide the anion [CuBr 4 ]2- is formed. Copper (III) occurs in a number of compounds, for example KCuO 2 and K 3 CuF 6 . Diamagnetic complexes of the type K 7 Cu(IO 6 ) 2 7H 2 O are formed when an oxidizing agent is added to an alkaline copper(II) solution contaning iodate or tellurate. The reactions of copper These tests illustrate the reactions of copper(II). They should be carried out on a 0.5M solution of copper(II) sulphate. For the tests record all observations are interpret the results and give the equations where possible. 1. To the cupric sulphate solution (4 cm3) add 2 mol dm-3 sodium hydroxide (8cm3), whereupon a precipitate is formed. Divide this suspension in to three portions. a. Boil the solution b. Add concentrated hydrochloric acid( 4cm3) carefully while shaking the solution to produce mixing. c. Add 60 % sodium hydroxide( 8 cm3 ) to the solution and warm the mixture gently. 2. To the copper solution add a small piece of zinc.

6. THE STABILIZATION OF OXIDATION STATES Experiment 5. COMPARISION WITH Cu(I) AND Cu(II) COMPOUND 1. Introduction When an element can exist in more than one oxidation state in aqueous solution each oxidation state will have a different thermodynamic stability. The relative of two oxidation states in aqueous solution is most conveniently expressed in terms of the electrode potential for the reaction Ma+ + (a-b)e- Mb+ where b < a

The electrode potential for a solution containing the ions Mb+ and Ma+ is given by the equation,

where

z = the number of electrons per ion transferred at the electrode

F = the Faraday = 96,480 C mol-1 E = the electrode potential of the solution E0 = the standard electrode potential [Ma+] = the activity of Ma+ ions in the solution [Mb+] = the activity of Mb+ ions in the solution Therefore any species added to the solution which reduces the concentration of either Ma+ of Mb+ and so alters the ratio [Ma+]/[Mb+] will cause an observable change in the electrode potential. If [Ma+] is reduced then the observable potential will become less positive, that is the higher oxidation state will become more stable. Alternatively, if [Mb+] is reduced the observed potential will become more positive and the lower oxidation state will become more stable. The electrode potential for a given reactoin indicates whether the stability of the higher oxidation state has been increased of decreased by the formation of a given compound. It does not provide any information as to how that increase or decrease in stability has been will now be considered.

The elements copper (3d10 4s1) and silver (4d10 5s1) have similar outer elecrton configurations but differ in the stability of the Cu1+ and Cu2+ oxidation states. For copper the 'normal' oxidation state in aqueous solution is Cu2+ shile for silver it is Ag1+. However, compounds in which the Cu1+ oxidation state has been stabilized can be prepared in aquesous solution and compounds containing Ag2+ can also be prepared. 2 . Insoluble compounds The standard electrode potentials using the IUPAC convention for the reactions Cu+ + e- = Cu0 Cu2+ + e- = Cu+ E0 = 0.52 V (9.1) E0 = 0.153 V (9.2)

show that Cu+ ions in aqueous solution with respext to disproportion into Cu0 and Cu2+ 2Cu+ = Cu0 + Cu2+ E0 = 0.367 V

Note that the IUPAC convention for electrode potentials differs from the convention used in a number of textbooks, where electrode potentials are written as oxidation potentials, e.g. Cu0 = Cu++e-, E0 = -0.52V. The numerical value of the electrode potential is the same but the sign is reversed. Consequently it is important to check which convention has been used for quoted values. From the equation

any species added to the solution which reduces the concentration of Cu+ but affects the concentration of Cu2+ to a lesser extent will cause an increase in the measured elecrtode potential. Such a reduction in concentration of Cu+ can be brought about by addition of an ion which forms an insoluble salt with Cu+ but not with Cu2+. Thus copper() chloride is insoluble in aqueous solution and the increase in stability of Cu+ is given by the electrode potential. Cu2+ + Cl- + e- = CuCl E0 = 0.566 V (9.3)

There will be a similar effect on the electrode potential Cu+ + e- = Cu0 E0 = 0.521 V

in the presence of Cl- due to the precipitation of copper() chloride CuCl + e- = Cu0 + ClE0 = 0.124 V (9.4)

Therefore by combination of eqns. (9.3) and (9.4) it follows that copper() chloride is stable with respect to the disproportionation. 2CuCl = Cu0 + Cu2+ + 2Cl3 . Soluble compounds The increased stability of the lower oxidation state of copper by precipitation as the chloride can be readily understood in terms of the reduced concentration of the Cu+ ion in solution. The reason for the stability of soluble copper() compounds is not immediately apparent. Stability can be achieved by complex formation. If the complex is sufficiently stable for the equilibrium [ML x ]+ M+ + xL to lie well to left then the concentration of Cu+ will be reduced. Ligands which, in addition to forming bonds with the metal ion, also function as -electron acceptors are more likely to form stable complexes. When the metal ion has a high electron density it will more readily donate -electrons to the ligand. The strength of the -bonding and hence the stability of the complex will be increased. For a given metal the lower oxidation state has the higher electron density and participates more readily in -bonding. Thiourea is a suitable ligand for the stabilization of Cu+ ions in solution, with coordination occuring through the sulphur atom. An indication of the stability of the tris(thiourea)copper() ion is given by the reaction between copper and hydrochloric acid. There is no reaction between copper and hydrochloric acid (see electrode potentials below) but on addition of thiourea hydrogen is liberated. The electrode potentials of the half reactions are, Cu+ + e- = Cu0 H+ + e- = 1/2 H 2 E0 = 0.52 V E0 = 0 (1 mol dm-3 HCl) E0 = -0.442 V

For the reaction Cu0 + H+ = Cu+ + 1/ 2 H 2 to proceed, the electrode potential for Cu0 + e= Cu0 under the conditions of the reactions must be

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INORGANIC CHEMISTRY EXPERIMENT Experiment 1 : PURIFICATION OF SODIUM CHLORIDE Experiment 2 : DOUBLE SALTS Experiment 3 : GROWING CRYSTAL FROM AQUEOUS SOLUTION Experiment 4 : REACTION OF Cr, Fe, Co, Ni, AND Cu ION Experiment 5 : COMPARISION WITH Cu(I) AND Cu(II) COMPOUND Experiment 6 : Preparation of [Co(NH 3 ) 6 ]Cl 3 and [Co(en) 3 ]Cl 3 Experiment 7 : PREPARATION OF [Co(NH 3 ) 4 CO 3 ]NO 3 & [Co(NH 3 ) 5 Cl]Cl Experiment 8 : Geometrical isomerism ( cis and trans-[CoCl 2 (en) 2 ]Cl 2 ) Experiment 9 : LINKAGE ISOMERISIM ( Nitro and Nitrito Covalt(III) Complex) Experiment 10 : CYCLE OF COPPER REACTION Experiment 11: Bioinorganic Chemistry: Synthesis and Study of an Oxygen- Carrying Cobalt Complex Which Models Hemoglobin. Experiment 12 : The Synthesis and Characterization of YBa 2 Cu 3 O x - the "1-2-3" Superconductor.
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