Chemistry 115B
Complexes and Coordination Compounds
of Transition Metal Elements
of Transition Metal Elements
Electronic Structures of Transition Metal Atoms and Ions
In the Periodic Table, there are three transition metal series: 1st, Sc-Zn; 2nd, Y-Cd; and 3rd, La-Hg. The chemical and physical properties of the transition metal atoms and elements do not change periodically across a period as much as the representative atoms and elements of the same period. Thus, there are similarities in properties not only in families (columns) but also across periods for the transition metal atoms and elements.
The transition metal atoms have a great variety of oxidation states. For example, manganese can form oxidation states of +2, +3, +4, +5, +6, and +7. Thus, the transition metals can form a great variety of ionic and covalent compounds.
Generally, the above experimental observations can be correlated to electronic configurations of the transition metal atoms. The transition metal atoms generally have the valence electronic configurations of ns2(n-1)dx where x = 1-10 and n = 4 for the first transition series, n = 5 for the second and n = 6 for the third. Because the inner (n-1) d-orbitals are being filled with electrons across a given series such as n = 4 period, the properties of the transition metal atoms and elements do not vary much. Also, the possible oxidation states relate to the number of ns-electrons and unpaired (n-1) d-electrons. All of the transition metal atoms form +2 ions due to losing the ns-electrons. Higher charged ions or higher oxidation states are formed formally by subsequently losing unpaired (n-1) d-electrons. For example, manganese (4s23dxy13dxz13dyz13dz213dx2-y21) has oxidation state of +2, +3, +4, +5, +6, and +7 while nickel (4s23dxy13dxz13dyz13dz213dx2-y21) has oxidation state of +2, +3, and +4. The higher oxidation states are found in species such as MnO4- and ReF7 containing covalent bonds to very electronegative atoms such as O and F.
Formation of Complexes and Coordination Compounds
Transition metal ions are Lewis Acids which tend to form stable Lewis Acid-Base adducts with Lewis Bases. These Lewis acid-base adducts are called complexes. Examples of these Lewis Bases, which are referred to as "ligands" in this context are :Cl:-, :CN:,- H2Ö: and :NH3. Some ligands contain more than one electron donating atom and are referred to as polydentate ligands. Examples of these are ethylenediamine, a bidentate ligand,
H2N-CH2CH2-NH2 and ethylene diaminetetracetate ion.
Examples of complexes are: Cu(NH3)42+, Ni(NH3)62+, CoCl42-
Color
Many transition metal complexes are colored. The color in these species is due to electronic transitions involving d-electrons, and the energies of these transitions are related to the ligands, the oxidation.state of the transition metal atom, and the n value in ns2(n-1)dx.
In the formation of complexes, differing numbers of ligands (coordination numbers) can associate with a central transition metal ion, and the resulting complexes can assume any of several possible geometric configurations. Usually two , four, or six ligands are found in complexes and typically they are arranged as shown below.
|
Number of Ligands |
Geometry |
Example |
|
two |
linear |
AgCl2- |
|
four |
tetrahedral |
CoCl42- |
|
six |
octahedral |
Ni(NH3)62+ |
1. Give the electronic configurations of
Co2+
Cr3+
Ni2+
Cu2+
Experiment:
1. Make 10 mL of a 0.1 M solution of Cu2+ from solid Cu(SO4)2 . 5 H2O (F.W. 249.68) and deionized water.
2. Transfer about 1 mL of the 0.1 M Cu2+ solution to a test tube and carefully add conc. HCl dropwise, with swirling, to this solution until a color change occurs. Gently heat this solution. Be sure to shake while heating to prevent "bumping".
3. Transfer about 1 mL of the solution from part 1 to a test tube and add 6M NH3(aq) dropwise with swirling. Continue the addition until the solid which was formed initially redissolves .
4. Carefully add conc. HCl dropwise to the solution in part 3 and observe any changes that occur.
5. Put about 1 mL of your 0.1 M Cu2+ solution from part 1 in an evaporating dish and gently heat over a flame to evaporate away the water. Observe the color of the anhydrous copper sulfate, CuSO4.
Exercise:
Copper forms an insoluble precipitate with hydroxide ion, Cu(OH)2 Ksp 2.2 x 10-20, and forms "complexes" with ammonia, Cu(NH3)42+, Kf 1.2 x 1012, chloride ion, CuCl42-, and water, Cu(H2O)62+.
1. Give equations to illustrate:
a) the reaction of aqueous copper ion in conc. HCl (exp. 2).
b) the observations you made when aqueous ammonia was added to a 0.1 M solution of copper ion (exp. 3).
c) the reactions that occurred when conc. HCl was added to the solution from experiment 3.
2. Assuming that the Cu(NH3)42+ complex has a square planar geometry, draw this complex.
3. Read sections 20.5-20.7 p. 917-926 in your lecture text for an explanation of the formation of color in coordination compounds and complexes. Note the spectrochemical series on the bottom of page 920 and note how it correlates with the colors you observed for the Cu(H2O6)2+, Cu(NH3)42+, and CuCl42- complexes.
Experiment:
1. Place 1 mL of a 0.1 M acidic solution of sodium vanadate, NaVO3, in a small test tube (test tube 1) for use as a color standard.
2. Into each of two small test tubes (test tubes 2 & 3) place 2 mL of violet aqueous V2+ solution.
3. To one of these violet V2+ solutions* (test tube #2) add 0.075 M (0.10 N) sodium hypochlorite (bleach) dropwise from a disposable pipette while taping the test tube to promote mixing. Count the number of drops of NaOCl solution necessary to just change the violet V2+ solution to the bright yellow V5+ state.
4. To the other test tube (Test Tube 3) continuing the V2+ solution add 1/3 as many drops of NaOCl solution as were added to test tube #2 and note the new color of the solution. Next add another 1.3 aliquot of the NaOCl solution and again note the color. Add the final 1/3 aliquot to convert the vanadium solution to the final 5+ oxidation state.
Exercise:
1. Using the discussion on page 1 of this handout and a periodic table, predict the probable oxidation states of vanadium ions.
2. When sodium vanadate is dissolved in an acidic solution the oxidation state of the vanadium stays the same but the actual species in solution changes as shown in the equation given below.
a) Determine the oxidation state of the vanadium in each vanadium containing species in this equation.
b) Balance the equation.
c) State the general type of reaction that is occurring when sodium vanadate is dissolved in an acidic solution.
3. a) Balance each of the following equations and state the type of reaction that is occurring in equation.
1) V2+(aq) + OCl-(aq) = V3+(aq) + Cl-(aq)
2) V3+(aq) + OCl-(aq) = VO2+(aq) + Cl-(aq)
3) VO2+ + OCl-(aq) = VO2+(aq) + Cl-(aq)
4) For each of the following species, state the color (as determined in your experiment) and give the correct oxidation state of the vanadium.
|
Species |
Color |
Oxidation State |
|
V2+ |
|
|
|
V3+ |
|
|
|
VO2+ |
|
|
|
VO2+ |
|
|
Experiment
1. Make 10 mL of a 0.1 M solution of a Co2+ solution from solid Co(NO3)2.6H2O (F.W. 291) and deionized water.
2. Transfer about 1 mL of the Co(NO3)2 solution to a test tube and add conc. HCl dropwise until the solution is blue.
3. Add water dropwise to the test tube in experiment 2 and observe the color change.
4. Place 2 mL of 0.1 M Co2+ solution from part 1 in a test tube and add conc. HCl dropwise until an intermediate purple color is obtained. Transfer half of this solution to another test tube, place one of these tubes in boiling water and the other in ice water, and observe the colors.
Exercise:
1. Draw the aqueous complex of Co(H2O)62+.
2. Assuming a tetrahedral geometry, draw the cobalt tetrachloride complex ion, CoCl42-.
3. Given the results of experiment 4, state whether the
formation of the CoCl42- complex is an
endothermic (D
H is positive) or exothermic process.
Experiment:
1. Make 10 mL of a 0.5 M Ni2+ solution from solid NiSO4 (FW 155) and deionized water.
2. Transfer 1 mL portions of the 0.5M Ni2+ solution to each of three test tubes. To the first test tube add 6M aqueous ammonia dropwise with swirling and note any color change. To the second add 1M ethylenediamine, H2N-CH2CH2-NH2 dropwise noting the color at about 0.5 mL (10 drops), 1 mL (20 drops), and 2-4 mL. To the third test tube add 1% dimethylglyoxime dropwise and observe the change.
Exercises:
1. Using the Framework Molecular Model Kit demonstrate the octahedral geometry found in the Ni(NH3)62+ complex ion.
2. Both ethylenediamine and dimethylglyoxime are so-called bidentate ligands in that they have two Lewis Base sites within a single molecule. Polydentate ligands are also referred to as chelating agents (from the Greek word for claw). Using an octahedral model and the drawing supplied with this handout, study the formation of the Ni(H2O)4(en)2+, Ni(H2O)2(en)22+, and Ni (en)32+ complexes. Using your model and the drawing for the square planar complex of dimethylglyoxime and Ni2+ show the shape of this complex.
3. Ethylenediaminetetraacetate ion (EDTA) can function as a hexadentate ligand and as such is an extremely effective chelating agent. Using the drawing on page 908 in your lecture test and your model to show the structure of an EDTA M2+ complex.
4. Using your model and the drawing on page 913 of your lecture text distinguish between the "cis" and "trans" isomers of the Co(NH3)4Cl2+ "complexes" of Co3+.
5. Using your models and the discussion on p. 915 in your text, show the difference between the two possible mirror image isomers of cis Co(NH3)4(Cl)2+.