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.
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(NH3)4]SO4
.H2O, 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.
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 partly-filled 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 hexa-halo complexes, e.g.[SnCl6]2- ,[TiCl6]2-. Vanadium(V) shows few resemblences to phosphorus(V) , although the tetrahedral oxychlorides VOCl3 and POCl3 are both liquids and are readly hydrolysed in water . Both chromium(VI) and sulphur(VI) form strongly acidic oxides CrO3 and SO3 , and the covalent oxyhalides CrO2Cl2 and SO2Cl2 are easily hydrolysed. Magness(VII) shows some formal resemblances to the halogens in the oxo-compounds Mn2O7, and MnO4
-. Both Cl2O7 and the corresponding
manganese compound are powerful oxidizing agents and they are of low stability. The tetrahedral oxyanions ClO4
- and MnO4- 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:
Cr2O72- + 14H+ + 6e- = 2Cr3+ + 7H2O Eo = 1.33V
The oxide, CrO3, 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,Cr2O3. It dissolves in mineral acids to give the hexa-aqua purple cation [Cr(H2O)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(H2O)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(H2O6]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 FeSO4(NH4)2SO4.6H2O (Mohr's salt) is used in volumetric analysis for
titrations with dichromate, permanganate and cerium(IV) solutions. In aqueous solution the octahedral ion [Fe(H2O)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+ + O2 + 2H+ = 2Fe3+ + H2O 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(H2O)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.[FeCl4]-. 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(C2O4)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(H2O)6]3+ + e- = [Co(H2O)6]2+ Eo = 1.85V [Co(NH3)6]3+ + e- = [Co(NH3)6]2+ 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 [CoCl4]2- shows a large absorption in the visible part of the spectrum, which accounts for its deep-bluecolour.
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.
le. Use the solution of 1 mol dm-3 cobalt(II) chroride or nitrate for the following tests.
olution (5 cm3) add slowly an aqueous solution of 60% sodium
2. Add 2 mol dm-3 sodium hydroxide (5cm3) to the cobalt(II) solution(5 cm3).
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 possib
1. To the cobalt(II) shydroxide (8 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(NO3)2
.6H2O, NiSO4.6H
NiCl2O and Ni(ClO4)2
.6H2O. note that
molar absorbances 1-10 and these are assigned to the three
octahedral complexes. Typical tetrahedral complexes are
ar complexes are [Ni(CN)4]2-, NiBr2(PEt3)2 and bis(dimethylglyoximato)nickel(II).
2.6H2O contains trans-NiCl2(H2O)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(NH3)6]2+ and [Ni(1,2-diaminoethanne)3]2+. The visible spectra of these complexese have three bands with lowspin-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[NiCl4]2- and NiCl2(PPh3)2. Square planar nickel(II) complexes are diamagnetic and often brown, red or yellow in colour. Typical nickel(II) square plan
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(H2O)6]2+. If ammonia added, up to four water molecules can be replaced stepwise to give the complex [Cu(NH3)4(H2O)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 [CuX4]2-(X = Cl,Br) which have distorted tetrahedral structures. For example, when lithium bromide or hydrobromic acid is added to cupric bromide the anion [CuBr4]2- is formed. Copper (III) occurs in a number of compounds, for example KCuO2 and K3CuF6. Diamagnetic complexes of the type K7Cu(IO6)27H2O 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 E0 = 0.52 V (9.1)
Cu2+ + e- = Cu+ 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(Ⅰ) 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(Ⅰ) chloride
CuCl + e- = Cu0 + Cl- E0 = 0.124 V (9.4)
Therefore by combination of eqns. (9.3) and (9.4) it follows that copper(Ⅰ) chloride is stable with respect to the disproportionation.
2CuCl = Cu0 + Cu2+ + 2Cl- E0 = -0.442 V
3 . 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(Ⅰ) compounds is not immediately apparent. Stability can be achieved by complex formation. If the complex is sufficiently stable for the equilibrium [MLx]+ → M+ + xL to lie well to left then the concentration of Cu+ will be reduced. Ligands which, in addition to forming σ bonds with the metal ion, also function as π-electron acceptors are more likely to form stable complexes. When the metal ion has a high electron density it will more readily donate π-electrons to the ligand. The strength of the π-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 π-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(Ⅰ) 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 E0 = 0.52 V
H+ + e- = 1/2 H2 E0 = 0 (1 mol dm-3 HCl)
For the reaction Cu0 + H+ = Cu+ + 1/2 H2 to proceed, the electrode potential for Cu0 + e- = Cu0 under the conditions of the reactions must be << 0. This can be achieved if the activity of Cu+ ions in solution is greatly reduced. From the equation
when T = 298.15 K, [Cu0] = 1 , E0 = 0.52 V
then
Therefore the maximum concentration of Cu+ ions in solution is < 10-8.8 mol dm-3. This means that in the presence of thiourea the concentration of Cu+ ions has been greatly reduced, due to complex formation. When stabilization of a higher oxidation state occurs, e.g. Ag(Ⅱ), then a ligand which can form π-bonds with the metal ion by donation of electrons to the metal ion may form a stable complex. When the metal ion has a high electronegativity it will more readily accept π-electrons from the ligand. The strength of the π-bonding and hence the stability of the complex will be increased. Electronegativity increases with increasing oxidation state of a given metal. Therefore metal ions of high oxidation state are more able to participate in this type of π-bonding. The increased stability of the complex again effectively reduces the concentration of metal ions in solution. The Ag(Ⅱ) oxidation state is unstable relative to Ag(Ⅰ), but coordination with pyridine to form the tetrapyridine silver(Ⅱ) ion results in the formation of a number of moderately stable compounds. See also p.88.
The examples of soluble complexes referred to above have used π-bonding ligands in order to obatain stable complexes. However, it is possible to stabilize a given oxidation state by complex formation in whiche no π-bonding occurs. Thus the formation of hexammine complexes of cobalt results in an increased stability of the Co3+ ion compared with the hexaquo complex
Co3+(aq) + e- = Co2+(aq) E0 = 1.842 V [Co(NH3)6]3+ + e- = [Co(NH3)6]2+ E0 = 0.1 V
The increased stability of the Co(Ⅲ) oxidation state in the cobalt hexammine complex compared with the cobalt hexaquo complex can be understood qualitatively, in terms of the crystal field stabilization energy. Awwuming that the electron-pairing energy and the 3rd ionization potential of the Co(Ⅱ)ion is approximately constant, then the gain in crystal field stabilization energy for Co(Ⅱ) in the hexammine compared with the hexaquo complex will be
-0.8(Δ"0 - Δ'
0) where Δ'
0 = crystal field splitting by H2O, Δ"0 = crystal field splitting by NH3
and for Co(Ⅲ) will be -2.4(Δ"0 - Δ'
0). Higher values of Δ0 will have a more marked effect on the stabilization of the Co(Ⅲ) oxidation state in a given complex and the position of the ligand in the spectrochemical seriex. For further information on the spectrochemical series see p.176. any effective change in the free energy for the reaction Co(Ⅲ) → Co(Ⅱ) will result in a change in the electrode potential, and the equilibrium constant as shown by the expression
ΔG0 = -nFE0 = -RTlnKp
An increase in crystal field stabilization energy will be a contributory gactor to a more positive ΔG0 and hence a less positive electrode potential. Note again (See p. 32), That a distinction must be made between thermodynamic and kinetic stability. The above discussion has refferd to the thermodynamic stability of a complex, that is its tendency to exist under equilibrium conditions. This is not neccessarily a measure of the kinetic stability, i.e. lability, of the complex with regard to ligand replacement reactions.
Procedure Copper(Ⅰ) chloride is prepared by reducing copper(Ⅱ) ions with sulphur dioxide or sulphite ions in the presence of chloride ions. The copper(Ⅰ) ions once formed, react with chloride ions to form the insoluble copper(Ⅰ) chloride. Prepare three solutions: (a) dissolve sodium sulphite (10g) in 50 cm3 of water, (b) dissolve copper(Ⅱ) chloride (13g) in 25cm3 of water, (c) prepare a wulphurous acid solution by dissolving sodium sulphite(1g) in 1 dm3 of water and add 12 cm3 of 2M hydrochloric acid.
Add slowly, with constant stirring, the sodium sulphite solution to the copper(Ⅱ) chloride solution. Dilute the suspension of copper(Ⅰ) chloride so formed with about half the sulphurous acid solution, allow the precipitate to settle, and decant most of the supernatant solution. Filter the solid by suction on a sintered glass disc, wash the precipitate on to the sinter by means of sulphurous acid solution. Take care that the copper(Ⅰ) chloride is always covered by a layer of solution. Finally wash the product with portions of glacial acetic acid, alcohol, and ether. Dry the product in a warm oven. Copper(Ⅰ)chloride is slowly oxidized by moist air to given the basic copper(Ⅱ) chloride, CuCl2.3Cu(OH)2, so it must be stored in stoppered containers.
Complementary work (1) Prepare 10-20 cm3 of an aqueous of potassium chloride and to this add -1 g of copper(Ⅰ) chloride. Note what happens and interpret your observations. Now add a few drops of ethylendiamine, identify the precipitate so formed, and comment. (2) In what other ways may the copper(Ⅰ) oxidation state be stabilized?
4. Properties of Cobalt(III) Complexes
Experiment 6. Preparation of [Co(NH3)6]Cl3 and [Co(en)3]Cl3
Suject 1. Preparation of [Co(NH3)6]Cl3
Hexaamminecobalt(III) salts may be prepared by any of three methods that depend on oxidation of cobalt(II) ion in ammoniacal solution:(1) air oxidation, with formation of the pentamine ion, which is converted to the hexamine by heating with aqueous ammonia under pressure; (2) oxidation with an agent such as hydrogen peroxide, iodine, potasium permanganate, lead oxide, or hypochlorite solution; (3) oxidation in the presence of a catalyst that allows equilibrium between the pentamine and hexamine ions to be established at room temperature and atmospheric pressure. These compounds may also be prepared indrectly from other hexaminecobalt(III) salts.
In the best of the catalytic methods diamminesilver ion or decolorizing charcoal is used as a fcatalyst. The method devised by J.Bjerrum in which decolorizing charcoal is catalyst is simple, give high yields of pure product, and is not time- consuming. A high concentration of ammonium salt is sufficient to stabilize the hexaminecobalt(III) ion, and the carbon serves only to establish the equilibrium. Air is used as oxidant except when the cobalt(II) compound is slightly soluble in the ammoniacal solution, as in the preparation of [Co(NH3)6]Br3, for which hydrogen peroxide is preferable.
procedure 240g (1mol) of cobalt(II) chloride 6-hydrate and 160g (3mol) of ammonium chloride are added to 200ml of water. The mixture is shaken until almost of the salts are dissolved. Then 4g of activated decolorizing charcoal and 500ml of concentrated ammonia are added. Air is bubbled vigorously through the mixture until red solution becomes yellowish brown (usally about 4hour). The air inlet tube is of fairly large bore(10mm) to prevent clogging with the precipitated hexamminecobalt(III) salt.
The crystals and carbon are filtered on a Buchner funnel and then added to a solution of 15 to 30ml. of concentrated hydrochloric acid in 1500ml of water; sufficient acid
reaction. The mixture is heated on a hot plate to effect complete solution and is filtered hot. The hexamminecobalt(III) chloride is precipitated by adding 400ml of concentarated hydrochloric acid and slowly cooling to 0oC . The precipitate is filtered, washed first with 60% and then with 95% alcohol, and dried at 80 to 100oC.
Ethylenediamine coordinates with metallic ions through both nitrogen atomes. The five-membered chelate rings that are thus formed are very stable. Many cobalt(III) ammines are converted by aqueous ethylenediamine to ethylenediamine)cobalt(III) chloride. Thus,Jorgensen prepared the salt by heating [Co(NH3)5Cl]Cl2 with aqueous ethylenediamine. Grossman and Schuck obtained the salt by oxidizing a mixture of cobalt(II) chloride, ethylenediamine, and water. The method described below has been developed from the latter suggestion.
material required en, HCl, CoCl2.6H2O
Procedure 61g of 30% ethylenediamine is partly neutralized with 17 ml of 6N hydrochloric acid and the resulting mixture poured into a solution of 24 g of CoCl2
.6H2O in 75 ml. of water. The cobalt is oxidized by bubbling a vigorous stream of air through the solution for three hours. The solution allowed to evaporate on a steam bath until a crust begins to form over the surface (the volume will be about 15 to 20 ml.); then 15 ml. of concentrated hydrochloric acid and 30 ml.of ethyl alcohol are added. After cooling, the crystals of [Co(en)3]Cl3 are filtered and washed with alcohol until the washings are colorless. They are then washed with ether or dried in an oven.
Properties [Co(en)3]Cl3 crystallizes in orange-yellow needles, which are readily soluble in water but insoluble in the usual organic sovents. Its solubility in 6 N hydrochloric acid is about 3 %. It is stable at temperatures as high as 200o and is decomposed only slowly by hydrogen sulfide and sodium hydroxide.
4. Properties of Cobalt(III) Complexes
Of particular importance to the development of coordination chemistry are metal complexes of the type to be synthesized and characterized in this experiment. Prior to 1950 reserch in this area was almost exclusively concerned with the investigation of complexes of transition metal ions with such monodentate ligands as Cl-, Br-, I-, NH3, pyridine, CN-, and NO2
- and bidentate ligands as ethylenediamine(H2NCH2CH2NH2), oxalate (-O2CCO2
-), glycinate (H2NCH2CO2-), and CO3
-. These complexes still form
the basis of a vast amount of research today, despite the more recent discoveries of the ligand properties of -H, -CH3, CO, H2C=CH2, and benzene, to mention a few.
Coordination compounds of Co(III) and Cr(III) have been of particular interest because their complexes undergo ligand exchange very slowly compared with complexes of many other transition metal ions. For example, Ni(NH3)6
2+ reacts virtually instantaneously with H2O to form Ni(OH2)6
2+. Under the same conditions, the analogous reactions of Co(NH3)6
3+ and Cr(NH3)63+ occur very slowly. This difference
in behavior of complexes of different metal ions has been qualitatively accounted for by ligand field theory and molecular orbital theory.
The slow reactivity of Co(III) complexes has made them suitable for extensive investigations. The structures of the octahedral Co(III) complexes which you will prepare are given below.
One important method of characterizing ionic substances is the determination of the ability of their solutions to conduct an electric current. Those substances whose solutions have the highest conductivity consist of the greatest number of ions. Thus, a one molar solution of [Co(NH3)4CO3]NO3 will have a lower conductance than a solution of [Co(NH3)5Cl]Cl2 of the same concentration. By measuring the conductivity of a solution of a compound, it is possible to determine whethera formula unit of that compound consists of 2, 3, 4, or more ions. Although measurements will be done on
water solutions of the complexes, the same information can frequently be obtained using organic solvents such as nitrobenzene or acetonitrile for ionic compounds that either are not very soluble in water or react with water.
Another very powerful method for establishing the identity of a complex is infrared spectroscopy. This technique examines the frequencies of the vibrational modes of a molecule. Thus, the infrared spectra of both of the preceding complexes exhibit absorptions at frequencies(commonly expressed in wave numbers, cm-1, i.e., reciprocal wavelength 1/ㄌ) characteristic of stretching and bending modes of the NH3 group. While the Co-N stretching modes are in principle also measurable, they sometimes occur at frequencies too low (below 650 cm-1) to be observed in most infrared spectrophotometers. More specialized and expensive infrared instruments are, however, available for studying these vibrations. The complex [Co(NH3)4CO3]NO3 also exhibits absorptions characteristic of the carbonate group. Because of its coordination to the metal ion, the CO3
2- group has a somewhat different spectrum in this complex than it has as the ion, as in NaNO3. The spectrum should also contain absorption bands resulting from vibrational modes of the NO3
- ion, very similar to those observed in NaNO3. In contrast, the spectrum of [Co(NH3)5Cl]Cl2 would be largely dominated by absorptions attributable to the NH3 groups.
In general, metal-Cl stretching frequencies are lower than can be observed on usual infrared instruments, and the ionic Cl- groups, of course, are not, in the solid state, strongly bonded to any other single atoms; thus, no absorption are expected in the infrared spectrum to indicate their presence in the compound. While infrared spectroscopy and other instrumental methods of compound characterization are emphasized in this and in other experiments in this book. it should be stressed that a quantitative elemental analysis is an absolutely essential step in determining the composition and structure of a new compound.
Experiment 7 . PREPARATION OF [Co(NH3)4CO3]NO3 & [Co(NH3)5Cl]Cl
The synthesis of [Co(NH3)4CO3]NO3 will be carried cot according to the unbalanced equation.
e first reaction in the preceding sequence probably involves the following mechanism.
The Co(NO3)2 that is available commercially has the formula Co(NO3)26H2O and very probably is a coordination [Co(OH2)6](NO3)2. Since Co(II) complexes, like those of Ni(II), react very rapidly by ligand exchange, the first step in the reaction might be expected to be Co(OH2)6
On the basis of mechanistic studies of reactions of [Co(NH3)4CO3]+ with acids, th
That O-C bond fission occurs in the intermediate has been established from 18O isotopic exchange studies in several similar reactions of carbonato complexes. The subsequent steps in this preparation involve the substitution of one ligand in the coordination sphere by another. At first glace, one might expect these reactions to proceed according to SN1 or SN2 mechanisms, but even now there is considerable debate as to how these substitutions actually proceed. In Experiment 2, you will have an opportunity to
postulate a mechanism, based on your rate data, for the reverse of the last reaction in the preceding series, the conversion of [Co(NH3)5Cl]2+ to [Co(NH3)5(OH2)]3+.
he reaction mixtures from the atmosphere. (This is required for preparations tha involve reactants or products that react with
t necessitate heating of the solutions should be carried out in an efficient hood.
Precedure
The precautions are necessary to protect t
moisture or oxygen in the air.) Operations tha
Subject 1. Preparation of [Co(NH3)4CO3]NO3
Dissolve 20g(0.21 mole) of (NH4)2CO3 in 60ml of H2O and add 60ml of concentrated aqueous NH3. While stirring, pour this solution into a solution containing 15g(0.0532 mole) of [Co(OH2)6](NO3)2 in 30ml of H2O. Then slowly add 8ml of a 30 per cent H2O2 solution.(Handle H2O2 with rubber gloves. If the affected area is not washed immediately with water, hydrogen eroxide can cause severe skin burns.) Pour the solution into an evaporating dish and concentrate over a gas burner in a hood to 90 to 100ml. Do not allow the solution to boil. During the evaporation time add, in small portions, 5g(0.05 mole) of (NH4)2CO3. Suction filter (with water aspirator: for better control of the vacuum, use a pinch clamp on the rubber tubing between the trap and filtration flask: see Figur 1-1) the hot solution and cool the filtrate in an ice water bath. Under suction, filter off the red product crystals. Wash the [Co(NH3)4CO3]NO3 in the filtration apparatus first with a few milliliters of water (the compound is somewhat
anol. Calculate the yield. (Save a portion of your product for the conductance determination.) soluble) and then with a similar amount of eth
Subject 2. Preparation of [Co(NH3)5Cl]Cl2
Dissolve 5.0 g of [Co(NH3)4CO3]NO3 in 50 ml of H2O and add concentrated HCl (5 to 10 ml) until all of the CO2 is expelled. Neutralize with concentrated aqueous NH3 and then add about 5 ml excess. Heat for 20 minutes, again avoiding boiling; [Co(NH3)5(OH2)] is formed. Cool the solution slightly and add 75 ml of concentrated HCl. Reheat for 20 to 30 minutes and observe the change in color. Purple-red crystals of the product separate on cooling to room temperature. Wash the compound several times,
3+
by decantation, with small amounts of ice-cold distilled water: then filter under a water aspirator vacuum with a glass fritted funnel (medium porosity). Wash with several milliliters of ethanol. Drying in an oven at 120℃ to remove solvent yields
]Cl2. Calculate the yield. (Do not discard the compound because part of it etics experiment. Experiment 2.)
o(NH3)4CO3]NO3 and [Co(NH3)5Cl]Cl2.
3. IR spectra of [Co(NH3)4CO3]NO3 [Co(NH3)5Cl]Cl2, indicating the absorptions H3, CO3, and NO3. What are the spectral similarities and differences
QUESTIONS . (Note
-
aration of [Co(NH3)4CO3]NO3.
4. How would you experimentally establish the fission of the O-C bond, rather than the Co-O bond, in the conversion of
[Co(NH3)5Clwill be needed in the kin
REPORT Include the following:
1. Percentage yields of [C
2. Values of AM for the preceding complexes and your conclusions as to the number of ions in each compound.
characteristic of Nbetween these two compounds ?
1. Outline a method of analyzing [Co(NH3)5Cl]Cl2 for its percentage Cl contentthat the ionic Cl is much more reactive than that in the coordination sphere.)
2. Outline a method of analyzing [Co(NH3)4CO3]NO3 for NH3 and Co content.
3. Balance equation (1) for the prep
5. The conductance of an aqueous solution of [Co(NH3)5Cl]Cl2 changes on standing overnight. Would you expect it to increase or decrease ?
6. 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 ?
7. Why are 500ml of the 0.001M [Co(NH3)4CO3]NO3 and [Co(NH3)5Cl]Cl2 solutions prepared when only 50 to 100ml are required for the conductance measurements ?
lly distinguish between these two compounds ?
. Will the cell constant, k, change if the electrodes in a conductivity cell are bent or oved? Why?
8. How do [Co(NH3)4CO3]NO3 and [Co(NH3)5Cl]Cl2 differ structurally ? How would you experimenta
9m
5. Isomerism of Octahedral Complexes.
Experiment 8. Geometrical isomerism ( cis and trans-[CoCl2(en)2]Cl2 )
Among the inorganic complex compounds that can be resolved into optically active forms, cis-[CoCl2(en)2]Cl2 is one of the best known and easiest to prepare. It is often used in stereochemical studies and as an intermediate in the preparation of other cobalt complex salts. The best direction for the preparation of the trans salt are those given in the thesis of Vera Tupizina. Jorgensen has described the conversion of the trans
to the cis: the latter may easily be resolved in the manner described by Bailar and Auten.
rnight before the bright-green square plates of the hydrochloride of the trans form are filterd. These are washed with alcohol and ether and dried at 110
drogen chloride is lost, and the crystals crumble to a dull-green powder.
600g of a 10 % solution of ethylenediamine is added, with stirring,to a solution of 160 g. of cobalt chloride 6-hydrate in 500 ml. of water in a 2-1. beaker or bottle. A vigorous stream of air is drawn or passed through the solution for 10 or 12 hours. (Longer aeration causes undesirable secondary reaction to take place.) Three hundred and fifty millimeters of concentrated hydrochloric acid is added, and the solution is evaporated on the steam bath until a crust forms over the surface (750 ml.). The solution is allowed to cool and stand ove
degree. At this temperature, the hy
B. Conversion To The cis Form
trans-[Co(en)2Cl2]Cl → cis-[Co(en)2Cl2]Cl
Conversion to the cis form is brought about by evaporating a neutral solution of the trans to dryness on the steam bath. The unchanged trans form may be washed out with a little cold, or the transformation may be completed by repeating the evaporation. It
ould be repeated not more than two or three times, however, as some decomposition kes place.
shta
hedral Complexes. 5. Isomerism of Octa
Experiment 9. LINKAGE ISOMERISIM ( Nitro and Nitrito Covalt(III) Complex)
Material required Chloropentamminecobalt(III) chloride, Sodium nitrite, aqueous ammonia(S.G. 0.88)
Precedure
1. Nitropentamminecobalt(III) chloride
Dissolve chloropentamminecobalt(III) chloride(1.5 g) in 20 cm3 of 2M aqueous ammonia. Warm on a water bath until the salt dissolves. Cool, and acidify with 4M hydrochloric acid to pH 4. Add sodium nitrite (2 g) and heat heat gently until the red precipitate first formed has dissolved. In some cases the red precipitate only appears tran- siently. Cool the solution, and add carefully, concentrated hydrochloric acid(20
crystals,and wash with alcohol. Record the yield. cm3). Cool in ice, filter off the yellow brown
2. Nitritopentamminecobalt(III) chloride
Dissolve chloropentaaminecobalt(III) chloride(1.5 g) in aqueous ammonia (30 cm3 of water to 10 cm3 of aqueous ammonia(S.G. 0.88)),warm if necessary to dissolve the salt. Neutralize the solution the solution to litmus to litmus with 4M hydrochloric acid, and cool. Add sodium nitrite (1.5 g) and allow the solution to stand for 1-2 hours. Cool in
nk product, wash with ice-cold water, and alcohol. Dry at
(1) Measure the infrared spectra of the two isomers in the regin 2.5-15μm. Use a freshly
ice,filter off the salmon-piroom temperature. Record the yield.
Complementary work:
prepared sample in the case of the nitrito isomer. See p.164 a discussion of infrared spectroscopy and the treatment of samples. Compare the spectra and comment.
(2) Place a sample of the nitrito complex in an oven at 150oC for about 2 hours, note what happen ter this treatment. Comment on the results.
s and measure the infrared spectrum of the sample af
1. IONIC SALTS
Experiment 10. CYCLE OF COPPER REACTION
This is a short practical designed to:
1. Remind you of some simple experimental techniques all requiring care. 2. To show you some typical transition element behaviour.
u some chemical properties of copper. 3. To teach yoThe cycle involves the conversion of copper metal through a series of intermediate copper compounds back to copper metal. By weighing the copper at the start and the end of the experimental yield may be calculated.
Procedure Weigh(accurately) the copper wire supplied. In fume hood, using a 250cm3 beaker, prepare a solution of Cu(NO3)2 by adding 4.0cm3 concentrated nitric acid to the copper wire. After reaction is complete, add distilled water until the beaker is about half full. Add 18cm3 5.0M NaOH to precipitate Cu(OH)2. With stirring, heat just to boiling to convert Cu(OH)2 to insoluble black CuO. Allow CuO to settle, then decant supernatant liquid. Add about 200cm3 very hot distilled water, settle, and decant one more. All 18cm3 5.0M H2SO4 to convert CuO to CuSO4. In the fume cupboard, add (all at once) 2.0g 20-mesh zinc metal to precipitate copper metal, stir until the supernatant liquid is colorless. When evolution of hy5cm
drogen gas has become slow, decant. Still in hood, add
Repeat. Wash with about 5cm methanol, allow to settle, and decant. Finally, wash with bout 5cm3 acetone (keep away from flames!), allow to settle, and decant. Place dish
over steam bath(may be a beaker of boiling water) at least 5min to dry product. Transfer roduct to weighing paper, and after cooling, obtain product weight with a good balance.
If time permits, repeat washing and drying to see that product is at constant weight.
3 distilled water, then 10cm3 concentrated hydrochloric acid to react with the excess of zinc. When hydrogen evolution has become very slow, remove to your own bench, warm, but do not boil. When no further hydrogen avolution can be seen, cool, decant, and transfer solid to porcelain evaporation dish. Wash product with about 5cm3 distilled water, allow to settle, decatn.
3
a
p
Experiment11: Bioinorganic Chemistry: Synthesis and Study of an Oxygen-Carrying Cobalt Complex Which Models Hemoglobin.
bjective: In this project, we will synthesize O a coordination complex of cobalt, and
r ill-characterized, metal-
semble those found in nature (e.g., as part of a peptide chain). The multidentate ligand used in this experiment is tetradentate N,N'-bis(salicylaldehyde)-ethylenediimine (abbreviated salenH2). It is obtained by the condensation reaction shown below:
demonstrate how it displays chemistry similar to that of metal-containing biological systems. A simple apparatus will be used to quantitatively monitor the gas-uptake/release reaction. Background: Bioinorganic chemistry is a growing field of inorganic chemistry that addresses a number of significant biological issues. Much of the work in this area involves attempts to model the activity of a complicated, ocontaining system with more simple coordination complexes that can be prepared in the laboratory. The coordination chemistry relevant to biological systems is reviewed in most inorganic textbooks, for example Chapter 30 in Cotton and Wilkinson (5th Ed) and Chapter 18 in Huheey (3rd Ed). This experiment involves use of a cobalt(II) complex, which serves as a model for oxygen transport systems. Oxygen transport and storage are accomplished in higher animals by the iron containing materials hemoglobin and myoglobin. The challenge in developing suitable model systems to study this chemistry is in designing appropriate metal coordination complexes which react with dioxygen simply by binding it intact, rather than by much more commonly encountered redox processes, which can result in LnM=O (with a terminal oxo ligand) or LnM-O-MLn (with a bridging μ-oxo ligand). Complexes of the latter type are usually inactive toward release of O2. Cobalt(II) complexes are known to form two types of adducts with dioxygen: a 1:1 adduct LnCo-O2, and a 2:1 peroxo structure LnCo-O-O-CoLn. Biologically active metal centers are very often found in a coordination environment that includes multidentate ligands. Therefore the synthesis of bioinorganic "model complexes" often makes use of multidentate ligands that re
Addition of a Co(II) salt to salenH2 results in formation of the Co(II) salen com plex:
This four-coordinate Co(salen) complex is reactive toward addition of donor ligands in forming five- or six-coordinate complexes. In the solid state it exists in two different forms (I, a dark red isomer inactive toward oxygen binding, and II, the active isomer).
The experiments below will be conducted in chloroform and dimethyl sulfoxide, Me2S=O(DMSO) solvents. Chloroform (CHCl3) is not a good ligand for transition metals, whereas DMSO is. Briefly, we will prepare salenH2, and its complex with Co(II), Co(salen), as a mixture f the brown and dark red forms. This mixture is heated to accomplish complete
his material is dissolved in DMSO in a closed system ygen uptake is quantitatively
is then dissolved in
. F. Floriani and F. Calderazzo J. Chem. Soc. (A) 1969, 946-953. R. H. Bailes and M. Calvin J. Am. Chem. Soc. 1947, 69, 1886-1893.
oconversion to the dark red form. Tcontaining pure oxygen to generate an active form. The oxmeasured at constant temperature and pressure. The oxygen adduct chloroform to release the O2. Additional literature references: E. I. Ochiai, J. Inorg. Nucl. Chem. 1973, 35, 1727, 3375. C
W. L. Jolly, The Synthesis and Characterization of Inorganic Compounds, p. 466. H. Diehl and C. C. Hach Inorg. Synth. 1950, 3, 196. Experiment taken from: Appleton, T. G. J. Chem. Educ. 1977, 443. Safety precautions: Co(salen) is reported to be toxic, therefore avoid inhalation of
quipment and Materials:
el (w/ 24/40 joint), vacuum desiccator with CaCl2, O2 tank with gulator and hose, O2-uptake apparatus, centrifuge, test tube (1 x 7.5 cm)
thanol (in an Erlenmeyer flask) add 1.0 g (1.08 mL) of ethylenediamine in
3.5 min, and leave the solution to cool in an ice bath. Filter the right yellow flaky crystals using house vacuum, wash with a small volume of ethanol,
funnel. Continue heating and stirring for about an hour, during which time a red
particles of this material. While DMSO is not itself very toxic, it can transport dissolved compounds through your skin. Therefore avoid skin contact with DMSO. EBuchner funnel. melting point apparatus, 100 mL 3-necked flask (24/40), magnetic stir-bar, addition funnre salicylaldehyde, 95% ethanol, ethylenediamine, Co(O2CMe)2.4H2O, diethyl ether, DMSO, CHCl3 Procedure (2 periods): Preparation of salenH2. To a solution of 3.9 g (3.4 mL) of salicylaldehyde in 40 mL of boiling 95% eportions with a pipette (otherwise, the reaction mixture may froth out of the flask). Stir the reaction mixture for band air-dry. Record the melting point and yield of this product. Record an IR spectrum (vibrational spectrum; KBr pellet). Preparation of Co(salen). Weigh 2.0 g of salenH2 into a 100 mL 3-necked flask fitted with a magnetic stirring bar, an addition funnel, and a condenser capped with a nitrogen inlet. Add 60 mL of 95% ethanol. Stir using the magnetic stir bar, and flush the apparatus with nitrogen (the synthesis requires the exclusion of oxygen, but extreme rigor is not necessary). Adjust the nitrogen flow to a steady rate (ca. one bubble/sec) and provide a steady flow of cooling water through the condenser. Immerse the flask in a water bath maintained at 70-80 °C. Cover the hot part of the system with Al foil. Dissolve 1.86 g of Co(O2CMe)2.4H2O in 9 mL of hot water and put it in the addition funnel. When the salenH2 has all dissolved, add the cobalt acetate solution from the
precipitate should form. Cool the flask by immersing it in cold water. Discontinue the nitrogen flow and filter off the solid in the air. Wash three times with 5 mL of water,
thanol. Dry the solid on the funnel and wash it two times with mL of diethyl ether. If further drying is needed use the vacuum desiccator (over
d an IR
s oxygen is absorbed, the height of the movable arm must be changed to keep the water level the same on both sides, which maintains a constant pressure on both sides. The difference in initial and final volume readings gives the amount of O2 absorbed.
then with 5 mL of 95% e5drierite). Record your yield. Recorspectrum (KBr pellet). Oxygen Uptake by Co(salen). Oxygen uptake will be measured with the apparatus diagrammed below. Stopper the side-arm test tube with a #49 suba seal when necessary. The movable arm is used to keep the pressure constant in the system. A
Weigh out 0.05-0.1 g of ground, dry Co(salen) into a side-arm test tube (1.5 x 15 cm). Place approximately 5 mL of DMSO in a beaker and bubble oxygen through it for a few
re within the apparatus is
seconds to saturate it with O2. Transfer the DMSO into a small test tube (1 x 7.5 cm) until it is filled to about 2 cm from the rim. Tilt the side arm test tube and carefully lower the small test tube into the side-arm tube without spillage. Flush the side-arm tube with a gentle stream of O2. During flushing, the movable arm can be moved up and down to aid in the replacement of air with pure oxygen. Insert a tight-fitting rubber stopper in the mouth of the tube. Adjust the movable arm to make the water levels the same in both tubes (i.e., pressuatmospheric). Note the water level in the graduated tube. Carefully invert the side-arm
tube (holding near the stopper to minimize heating by the hand, but being careful not to push the stopper furhter into the tube causing a change in pressure) so that the DMSO is introduced onto the Co(salen) without spilling any on the tygon connecting tube. Gently shake the tube. As oxygen is absorbed the water level in the graduated tube should begin to rise. Note the changes that occur. The tube can be
nd read off the new level in the graduated tube.
2 absorbed per mole of Co(salen) can be calculated. You should do 2-3 riments to test reproducibility.
much as possible of the dark-brown spension into a centrifuge tube. Centrifuge until the precipitate has settled to the
pernatant DMSO. To the residue in the tube
O2 adduct. Report what happens pon dissolution of the O2 adduct in chloroform, and explain this result. Why is the
active toward O2, whereas the red form is not? Why does
similar activation? Explain your answer.
tipped to the side to increase the surface area of the solution and increase the rate of oxygen absorption. Continue shaking until no further change in water level occurs (5-10 min). Adjust the movable arm so that the water levels in the two tubes are again equal, aFrom the decrease in volume at room temperature and atmospheric pressure, the number of moles of OO2-uptake expe Now test the reactivity of Co(salen) toward O2 in chloroform, using the above procedure. Do this once. Behavior of the oxygen adduct in chloroform. Remove the stopper from the side-arm tube in the DMSO reaction, and remove assubottom of the tube. Carefully remove the su(drying is not necessary) add 5-10 mL of chloroform without stirring. Observe the result and note your observations in your notebook. Some points to address in your report: Report synthetic yields, giving amounts of starting materials used (weights and moles). Report the colors of the two products. Interpret the results of the O2 experiments, including a discussion of possible structures for the ubrown form of Co(salen)addition of DMSO to the red form activate it? Would addition of chloroform to the red form cause Additional questions: 1) When salenH2 reacts to form the cobalt(II) complex, 2 protons are lost. Where do they go?
2) What are the pKa’ of phenol and acetic acid? What happens when 1 equivalent of phenol is added to sodium acetate?
ded to salenH2.
n complexes. ) Discuss the O2 bond-order in metal dioxygen complexes (an MO diagram of O2 is elpful!) ) What are the oxidation state, spin state, and electron configuration of the cobalt in the 2 adduct?
3) Given your answer to 2, explain why a cobalt(II) salen complex forms when cobalt acetate is ad4) Give examples for different dioxygen binding modes in mononuclear and dinuclear metaldioxyge5h6O
Experiment 12: The Synthesis and Characterization of YBa2Cu3Ox - the "1-2-3" Superconductor.
Background: The discovery of high temperature copper oxide superconductors (specifically the compound YBa2Cu3O7, often referred to as 1-2-3 superconductors) created an extensive burst in the study of materials science and superconductivity theory which continues unabated even today. The resulting interest in chemical aspects of superconductivity and the ready availability of a refrigerant, liquid N2 (bp = 77 K), necessary to demonstrate superconducting phenomena created a secondary increase in the incorporation of these topics in chemical education. The traditional method of ceramic synthesis involves the grinding of metal powder oxides in the proper stoichiometry, followed by heating in a furnace. These "shake and bake" methods have several disadvantages: the repeated mixing and grinding of the oxide powders needs to be conducted in a fume hood to minimize exposure to dust; the mixing itself might be incomplete due to the relatively large grain
as been carried out in the past five years to improve the nds in order to achieve a greater uniformity in their
ed by the method of homogeneous co-recipitation. An aqueous urea/oxalic acid solution of the metal salts in the proper 1-2-3
stoichiometry is prepare tween 80 and 100°C hydrolysis of urea takes place with the simultaneous evolution of CO2 and NH3
sizes of commercially available oxides, resulting in failure to achieve the intimate mixing necessary to create a homogeneous composition; and these methods do not reflect the various ways ceramic materials synthesis has been transformed through the application of chemical procedures. In fact, much research hsynthesis of these compoucomposition. The homogeneous coprecipitation synthetic method produces a reproducibly consistent ceramic precursor powder which can be processed into a reliable superconductor. Synthesis of YBa2Cu3Ox. YBa2Cu3 oxide pre-ceramic powder is preparp
d and heated. At a temperature be
according to the following equation:
CO(NH2)2 + H2O → 2NH3 + CO2
As the urea hydrolyzes, the pH of the solution gradually rises. Increasing the pH causes the metal ions to precipitate out as their hydroxide, oxalate or carbonate salts. The final stoichiometry of the powder is determined by the final pH of the solution which in turn is determined by the initial molarity of the urea solution. An initially high urea concentration favors the proper 1-2-3 stoichiometry in the pre-ceramic powder. The precipitate is separated, washed and dried. It is then heated at 900°C for a least 16 hours in air to burn out all the residual carbon. The powder obtained at this stage is pressed into a pellet and sintered at 900°C for 4 hours, followed by annealing in oxygen at 500°C for another 16 hours. This results in a material with a composition of
Ba2Cu3Ox(6.5<x<7.0). The pre-ceramic powder synthesis part of the experiment can e 3-4 hour laboratory period. The calcining, pellet pressing,
al work, but the total working time
7, 64, 836-853. iu, R.S.; Chang, C. T.; Wu, P. T. Inorg. Chem. 1989, 28, 154.
. M. Acc. Chem. Res. 1988, 21, 8.
ions in a ood wearing gloves. Barium is quite toxic and inhalation of it poses the most
n this experiment. The starch indicator solution is
Yeasily be carried out in onand final sintering and annealing necessitate additionin the laboratory for these steps is short (~1 hour). Literature References: Several articles in J. Chem. Ed. 198LHolland, G. F; Stacy, AMüller, K. A.; Bednorz, J. G. Science 1987, 237, 1133. Pool, R. Science 1988, 241, 655. Safety precautions: In addition to the usual precautions, conduct all grinding and transfer operathsignificant chemical hazard ipreserved with mercurous ion as a biocide. This is also quite toxic. It should be handled with care, with gloves, and if a spill oc Equipment and Materials: High Temperature Tube Furnace, Combustion Boats, Heat-proof gloves, AsbC
Procedure (2 periods): 1) In a 250 mL flask, prepare 100 ml of solution that is 12 to 14 M in urea and 0.5 M in
ht
e with a spatula. Add this mixture to the urea solution and stir to dissolve.
to cover the urea solution. Heat the solution until is at 90-100°C and then continue heating for another hour. Evolution of CO2 and NH3
emperature
followed with ethanol. Leave the suction on for several
t maker similar to the kind used for
oxalic acid (C2O4H2.2H2O). This typically involves 80 g of urea, 6.3 g of oxalic acid and 35 mL of deionized (DI) water. 2) Warning - both the yttrium and copper nitrates are extremely hygroscopic (i.e., they absorb moisture from the air and become nonstoichiometric). In a 50 ml beaker, weig2.70 g of Y(NO3)3.5H2O, 5.10 g of Cu(NO3)2.2.5H2O, and 3.80 g of Ba(NO3)2 and stir the mixtur3) Place a 600 mL beaker (for water bath) onto a hot plate and clamp the flask inside the beaker. Add enough water to the beaker so asitshould be observed during heating. 4) After one hour of heating cool a small portion of the solution to room tand measure its pH with pH paper. If the pH is 7, allow the solution to cool to room temperature, otherwise, continue heating until the pH reaches a value of 7. 5) Add ~100 ml of DI water to the solution and stir it for 10 minutes before filtering. This will help dissolve any unreacted urea from the copper oxide precipitate. Filter the pale blue precipitate in a large (10 cm) Buchner funnel using a suction apparatus. If the mixture filters slowly, stir the precipitate with 100 mL more DI water and refilter. Wash the precipitate with waterminutes to dry the powder then place the mixture in a beaker and let it dry in an oven at 140°C for several hours. 6) Place the powder in a porcelain crucible and calcine it in air by heating at 900°C for 16 hours in a tube furnace. 7) Examine the powder after it has been removed from the oven and cooled. YBa2Cu3Ox is black. A green coloration in the powder indicates the formation of Y2BaCuO5 and CuO caused by a non-1-2-3 stoichiometry in the starting preceramic powder. The presence of the green phase usually does not interfere with the superconductivity. At this stage the powder can be molded into a 3 mm thick pellet through the use of a hydraulic press and an IR pellemaking KBr pellets. The pellet should be pressed @ 15,000 psi for about 3 - 5 mins. The pellet must be removed with extreme care to avoid crumbling it. Consult your TA about use of the press. If you have enough material, press two pellets.
8) The pellets should be sintered under nitrogen to give them structural integrity. Place the pellets in a porcelain crucible and heat the material under nitrogen at 900°C for 4 hours followed by annealing at 500°C for 18-24 hours under oxygen. Turn off the
rnace, and allow it to cool to room temperature while maintaining the oxygen flow.
best to use samarium-cobalt or neodynium-iron-
often drift off the perconductor and into the liquid nitrogen. Care must be observed when handling the
temperature can cause frostbite.
reservatives. This solution is stable for a few eeks, but should be standardized within a week or two of its use.) Standardize this
solution against a standard Cu
2Cu2+ + 5I- 2CuI + I3-
y HNO2 and other oxides of nitrogen. Cool to room
fuThe material obtained after this procedure should be black. Meissner Effect. The determination as to whether or not the YBa2Cu3Ox made is a high Tc superconductor is most easily done by observing the Meissner effect. Cooling a pellet of the material in liquid nitrogen and levitating a ferromagnet over its surface is a positive test for superconductivity. It isboron ferromagnets (which should be about one-third or less the size of the pellet) as they possess stronger magnetic fields which allow the magnet to levitate higher above the superconductor. The cut off bottom of a styrofoam coffee cup makes a suitable reservoir for the liquid nitrogen and pellet. A pair of plastic tweezers is useful for handling the magnets as they suliquid nitrogen as its extremely low Characterization (1-2 periods): Regrind the lump of sample and use it to analyze the copper content of the material. Iodometric Analysis of Copper Oxidation States in YBa2Cu3Ox. Prepare a 0.03 M Na2S2O3 solution by dissolving 2.36 g of Na2S2O3 plus 0.05 g of Na2CO3 in 500 ml of deionized water. Add three drops of chloroform and store it in an amber bottle. (The sodium carbonate and chloroform act as pw
solution.
→I3
- + 2S2O32- → 3I- + S4O6
2-
Prepare a standard copper solution by weighing accurately 0.5 - 0.6 g of reagent Cu wire into a 100 mL beaker. In the hood, add 6 mL distilled water and 3 mL of 70% nitric acid. Cover with a clean watch glass, and boil the mixture gently on a hot plate until the copper dissolves. Add 10 mL of distilled water and boil gently. Add 1.0 g urea and boil for 1 min to destro
temperature, transfer to a 100 mL volumetric flask, and dilute to 100 mL with 1 M HCl (be sure to rinse the beaker). The ease of oxidation of I- by O2 in acid requires that the titrations be conducted rapidly and under a brisk flow of N2. Using one of the tall-form beakers with a magnetic stirrer, loosely fit a two hole stopper to the top of the beaker. Through one hole, place a 50 mL buret. Through the other hole, pipet in 10.00 mL of your standard Cu solution, then insert a disposable pipet connected to a nitrogen tank and flush with N2. Remove the cork just long enough to add a solution of 1.0 -1.5 g of KI in 10 mL distilled water (freshly prepared). With stirring, titrate the solution with the Na2S2O3 solution using the buret. Add 5-8 drops of starch indicator just before the last trace of I2 color (light brown) disappears and the solution takes on a pale pink color. If the starch is added too soon, the I2 may bind to the starch and obscure the endpoint. After the addition of starch, the endpoint is reached when the solution appears lavender or pale violet in color.
etermine the molarity of your thiosulfate solution. The standardization should be
topper, flush veral minutes with N2, add 10 ml of distilled water containing 1.0 - 1.5 g of KI, and
tion that is 0.7 M in KI and stir for one minute. Add 10 mL of distilled water and titrate ess counts here). Do these
Cu3+ + 4I- → CuI + I3-
ach mole of Cu3+ produces one mole of I3- and each mole of Cu2+ produces half a mole
titration data. Report the total copper content and the amounts of Cu2+
Dperformed at least three times. Titration A: To determine the total Cu content of your superconductor, accurately weigh 150 -200 mg of the ground sample and dissolve it in 10 ml of 1.0 M HClO4 in a titration beaker in the hood. Boil the covered mixture gently in the hood to ensure complete destruction of Cu3+. Cool to room temperature, attach it to the ssetitrate with stirring as before. Do these titrations at least three times each. Titration B: To determine the amount of Cu3+ present, accurately weigh 150 - 200 mg of sample into a titration beaker and attach the stopper. Begin the N2 flow, and add 10 mL of a 1.0 M HClO4 solu
as described above (quickntitrations at least three times each.
Eof I3 Some Points to Address in Your Report: Report all of your and Cu3+ in your sample. Determine the standard deviations in your copper determinations.
Determine the value of x in the formula of the compound and determine the uncertainty in your value of x.
te x b using: x = 7/2 + 3/2(2+p)
b = volume of thiosulfate for b;
eport if you were successful in observing the Meissner effect with your sample. What ur titration beaker?
why there are differences in the Meissner effect from sample to sample.
terials? What leads to their
) What is the purpose of the various furnace procedures? 5) What is the importance of the O stoichiometry in the superconductor?
calcula y
where
Titration a = total Cu titration Titration b = Cu3+ titration Va = volume of thiosulfate for a; Ma = mass of material for a; VMb = mass of material for b. Why does these equations give you x? Ris the solid that forms in yo Additional Questions: 1) Discuss 2) Discuss how the pH at which the reaction is terminated affects the nature of the final product. 3) What are the colors of the Y, Ba, and Cu starting macolors, or lack thereof? 4
