Transition Metal Chemistry

TRANSITION METAL CHEMISTRY

The term ‘transition elements’ originally was coined to denote elements in the middle of the periodic table that provided a ‘transition’ between the ‘base formers’ on the left (Groups 1A and 2A) and the ‘acid formers’ on the right (Groups 5A through 7A).

Recall that metal oxides typically form basic aqueous solutions whereas non metal oxides typically form acidic aqueous solutions.

The term ‘transition elements’ actually applies to both ‘d’ and ‘f’ transition elements, all of which are metals, but commonly is used to denote only ‘d’-transition metals. The f-transition metals are usually called ‘rare earths’ or ‘inner transition metals’.

Transition metals are located between Groups 2A and 3A. Strictly speaking, d-transition metals must have partially filled d-orbitals. Zn, Cd, and Hg (Group 2B) have completely filled d-orbitals and are actually ‘post transition metals’ but they are often referred to as transition metals because of similar properties.

Cu, Ag, and Au (Group 1B) and Pb (Group 8B) also have filled d-orbitals however their cations (except Ag+) have partially filled d-orbitals.

General Properties of Transition Metals:

All are metals

Most are harder, more brittle, have higher mp and bp and Hvap than non transition metals.

Their ions and compounds are often colored.

They form many coordination complexes

Most have multiple oxidation states

Many of the metals and their compounds are good catalysts.

Electron Configuration:

Period 4

/ Period 5 / Period 6
21Sc [Ar] 3d1 4s2 / 39Y [Kr] 4d1 5s2 / 57La [Xe] 5d1 6s2
22Ti [Ar] 3d2 4s2 / 40Zr [Kr] 4d2 5s2 / 72Hf [Xe] 4f14 5d2 6s2
23V [Ar] 3d3 4s2 / 41Nb [Kr] 4d4 5s1 / 73Ta [Xe] 4f14 5d3 6s2
24Cr [Ar] 3d5 4s1 / 42Mo [Kr] 4d5 5s1 / 74W [Xe] 4f14 5d4 6s2
25Mn [Ar] 3d5 4s2 / 43Tc [Kr] 4d5 5s2 / 75Re [Xe] 4f14 5d5 6s2
26Fe [Ar] 3d6 4s2 / 44Ru [Kr] 4d7 5s1 / 76Os [Xe] 4f14 5d6 6s2
27Co [Ar] 3d7 4s2 / 45Rh [Kr] 4d8 5s1 / 77Ir [Xe] 4f14 5d7 6s2
28Ni [Ar] 3d8 4s2 / 46Pd [Kr] 4d10 5s0 / 78Pt [Xe] 4f14 5d9 6s1
29Cu [Ar] 3d10 4s1 / 47Ag [Kr] 4d10 5s1 / 79Au [Xe] 4f14 5d10 6s1
30Zn [Ar] 3d10 4s2 / 48Cd [Kr] 4d10 5s2 / 80Hg [Xe] 4f14 5d10 6s2

Energies of 3d and 4s orbitals are nearly equal. Generally the 3d orbitals fill after the 4s-orbital is filled. Cr and Cu are exceptions to the filling order in the first transition series. They have only one electron in their 4s orbital and one ‘extra’ electron in their 3d orbital. This gives a more favorable configuration in accordance with Hund’s rule.

The 4d and 5s, and 5d and 6s orbitals are even closer in energy than the 3d and 4s orbtitals making electron configurations difficult to predict for the 2nd and 3rd transition series.

The properties of the transition metals can be correlated roughly with either the total number of d-electrons or the number of unpaired electrons.

Melting Points:

Alkali metals melt below 200 °C. Several post transition metals are low melting (Ga = 30 °C).
Transition metals typically melt above 1000 °C. Tungsten is highest melting, above 3400 °C. Of all the elements, carbon has the highest melting point, i.e., ca. 3800 °C. Hg is the exception to the rule. Hg, a liquid at room temperature, has the lowest mp of all metals (-39°C).

Atomic Size:

Atomic size decreases left to right across each transition metal series as net core charge increases but size then increases at the far right (Group 1B and 2B) as valence electrons repulsion increases as d-orbitals fill.

Atomic size, as expected, increases down all Groups in the periodic table as additional shells of electrons are added with increasing atomic mass. This holds for the 1st and 2nd row transition metals but the 3rd row transition metals are the same size as the 2nd row. This unexpected ‘shrinkage’ is called the ‘lanthanide contraction’.

The lanthanide contraction of period 6 (the 3rd transition series) occurs because these elements contain an ‘additional’ 14 electrons in the 4f orbitals. The effective nuclear charge increases by 15 between La and Hf (the 1st and 2nd elements in the 3rd transition series).

Results of the Lanthanide Contraction:

1. 3rd transition series metals have the highest densities of all elements, for example:

 (g/mL) /  (g/mL)
Os / 22.6 / W / 19.3
Ir / 22.4 / Au / 19.3
Pt / 21.5 / Hg / 13.5
Re / 20.8 / Pb / 11.3

Fe = 7.9 g/mL

1. 3rd transition series metals have unusually high ionization energies, i.e., are rather unreactive.
2. 3rd transition series metals have much higher oxidation resistance. The platinum metals, i.e., Ru, Os, Rh, Ir, Pd and Pt (plus Au) do not form simple cations or even oxyanions. Au and Pt are especially useful in low voltage circuits where trace oxidation is problematic.

Acidity and Basicity of Transition Metals:

Recall that nonmetal hydroxides and oxyacids increase in acidity as the oxidation number of the central nonmetal increases. For example HClO4 (Cl = +7) is more acidic than HClO3 (Cl = +5).

Similarly, the acidity of the hydroxides and oxyacids of transition metals increases as the oxidation number of the central transition metal increases.

Oxid. # / +1 / +2 / +3 / +4 / +5 / +6 / +7

Mn

/ MnO / Mn2O3 / MnO2 / MnO3 / Mn2O7
 H2O /  H2O /  H2O /  H2O /  H2O
Mn(OH)2
manganese(II) hydroxide / Mn(OH)3
manganese(III) hydroxide / MnO(OH)2 (H2MnO3) manganous acid / MnO2(OH)2 (H2MnO4) manganic acid / MnO3(OH) (HMnO4) permanganic acid

BasicAmphotericAcidic

Note: The oxyacids of Cl are all acidic. They are shown here only to see an analogous oxidation pattern.

Oxid. # / +1 / +2 / +3 / +4 / +5 / +6 / +7
Cl / HClO
[ClOH]
hypochlorous acid / HClO2 [ClO(OH)]
chlorous acid / HClO3 [ClO2(OH)]
chloric acid / HClO4 [ClO3(OH)]
perchloric acid

Cr

/ CrO / Cr2O3 / CrO3
 H2O /  H2O /  H2O
Cr(OH)2
chromium(II) hydroxide / Cr(OH)3 chromium(III) hydroxide / CrO2(OH)2 (H2CrO4) (H2Cr2O7) chromic acid

General Trends in Acidity:

The more electronegative (EN) the central atom in an oxyacid, the greater the acidity. The following sequence of decreasing acidity (from left to right) with decreasing EN illustrates. Get the pKa values from the unit on ‘The Atom’ and compare EN values.
H2SO4 > H2SeO4 > H3PO4  H3AsO4 > H4GeO4

For a given central atom, the acid strength increases with the number of oxygens it holds.
H2SO4 > H2SO3 andHNO3 > HNO2

Problem 1: Write the formulas of all four oxyacids of bromine. Name them and number them in order of acidity where 1 is most acidic and 4 is least acidic. Explain why this trend in acidity occurs. Hint: Consider the oxidation state of bromine and consider the stability of the conjugate base that forms when the acid gives up an acidic hydrogen atom.

Transition Metals as Catalysts:

Transition metals and their compounds function as effective catalysts in both homogeneous (single phase) and heterogeneous (multiple phase) reactions.

Unreactive metals such as Pt, Pd, Ni and Au are sometimes used in a finely divided state to provide surfaces upon which heterogeneous reactions occur, e.g., hydrogenation of unsaturated organics.

Other transition metals act as homogeneous catalysts having d-orbital vacancies that can accept electrons from reactants to form intermediates that subsequently decompose.

Some typical reactions catalyzed by transition metals follow.

1. Haber Process: [Fe2O3, 500 °C, 400 atm]
   N2 + 3 H2 2 NH3
2. Contact Process: [V2O5, 400 °C]
   2 SO2 + O2 2 SO3
3. Hydrogenation: [H2 /Ni @ 25 °C, 1 atm]
   CH2=CH-CH2CH3 CH3CH2CH2CH3
   
4. Ostwalt Process: [Pt, 850 °C]
   4 NH3 + 5 O24 NO + 6 H2O
   2 NO + O2  2 NO2
   NO2 + H2O  HNO3
5. Catalytic Converters: [Pd or Rh]
   CO + HC’s + O2 CO2 + H2O
   NO and NO2  N2 + O2
   
6. Linear Polyethylene: [TiCl4 or MoO3]
   CH2=CH2(CH2CH2) n
   
7. Redox Reactions: [Mon(PO4)m]
   AsF3 + H2O2 AsF5

Classification into Subgroups:

The transition metals including Zn, Cd and Hg (Group 2B) are divided into eight Groups designated as B-Group elements. The number designates the maximum oxidation number of the Group members. No simple ion of these elements possesses a charge greater than +3.

Elements in corresponding A and B-Groups form many compounds of identical stoichiometry as illustrated by the following examples.

1A / 2A / 3A / 4A / 5A / 6A / 7A
NaCl / MgBr2 / Al(NO3)3 / CCl4 / POCl3 / SO4-2 / Cl2O7
KNO3 / CaCl2 / Ga(OH)3 / PbO2 / PO4-3 / H2S2O7 / HClO4
1B / 2B / 3B / 4B / 5B / 6B / 7B
CuCl / ZnBr2 / Sc(NO3)3 / TiCl4 / VOCl3 / CrO4-2 / Mn2O7
AgNO3 / CdCl2 / Y(OH)3 / ZrO2 / VO4-3 / H2Cr2O7 / HMnO4

Despite similar stoichiometry in compounds, the chemical properties of A-Group and B-Group elements are dissimilar.

Group 1B and 2B metals have filled d-orbitals, and d- and s-orbitals, respectively (pseudo-noble gas configurations) are unusually stable. In contrast, Groups 1A and 2A metals are very reactive and are never found unreacted in the native state.

Group 8B

Fe / Co / Ni
Ru / Rh / Pd
Os / Ir / Pt

Group 8B consists of three columns of three metals each. Each horizontal row is called ‘triad’ and is named after the best known metal of the row, i.e., the iron triad, the palladium triad, and the platinum triad. Greater horizontal similarities than vertical similarities are found in Group 8B.

The iron triad metals (Fe, Co, and Ni) are the only elements in the periodic table that exhibit ‘ferromagnetism’ in the uncombined state.

Problem 2: Without looking this up, write the formula of

a)a common oxide of tungsten

b)a naturally occuring oxide of osmium

Oxidation States:

Characteristically, transition metals can exhibit more than one oxidation state. In most cases, the maximum oxidation state found in a group is the same as the group number, but this is usually not the most common oxidation state. The tables below list common oxidation states.

1st Transition Series

3B

/ 4B / 5B / 6B / 7B / 8B / 1B / 2B
Sc /

Ti

/ V / Cr / Mn / Fe / Co / Ni / Cu / Zn
+7
+6 / +6
+5
+4 / +4 / +4
+3 / +3 / +3 / +3 / +3 / +3 / +3 / +3
+2 / +2 / +2 / +2 / +2 / +2 / +2 / +2 / +2
+1

The most common oxidation states are in bold. Not all oxidation states are shown.

2nd Transition Series

3B

/ 4B / 5B / 6B / 7B / 8B / 1B / 2B
Y /

Zr

/ Nb / Mo / Tc / Ru / Rh / Pd / Ag / Cd
+8
+7
+6 / +6
+5 / +5
+4 / +4 / +4 / +4 / +4
+3 / +3 / +3 / +3 / +3
+2 / +2 / +2 / +2 / +2
+1

s-electrons are outside the d-electrons and are removed first.

Grp. 2B & 3B elements have fixed oxidation states, +2 and +3, respectively (except Hg)

1st transition series metals form ionic compounds which dissolve in water giving solvated cations (Cr+3, Fe+3, Co+2) whereas 2nd and 3rd transition series metals form only water soluble oxyanions, e.g., MoO4-2, WO4-2, etc.

Higher oxidation states are more common in the 2nd and 3rd transition series

Element

/ Half-Reaction / Reduction Potential, E°, (V)
Sc / Sc+3 + 3e-  Sc / -2.08
Ti / Ti+2 + 2e-  Ti / -2.00
V / V+2 + 2e-  V / -0.26
Cr / Cr+2 + 2e-  Cr / -0.56
Mn / Mn+2 + 2e-  Mn / -1.02
Fe / Fe+2 + 2e-  Fe / -0.41
Co / Co+2 + 2e-  Co / -0.28
Ni / Ni+2 + 2e-  Ni / -0.23
Cu / Cu+2 + 2e-  Cu / +0.52
Zn / Zn+2 + 2e-  Zn / -0.76

Elements with oxidation states below the most common for a particular metal usually act as reducing agents. Elements with oxidation states above the most common for a particular metal usually act as oxidizing agents.

Most transition metals are moderately strong reducing agents that react with dilute mineral acids (HCl, H2SO4, HNO3) liberating H2 gas. The previous table shows that the reduction potential of most transition metal ions is unfavorable (negative values, i.e., not spontaneous). Rather the metals are readily oxidized to their ions. For this reason, few are found as pure metals in the native state. Most are readily oxidized and thus occur naturally as oxides. Examples include TiO2 (rutile), MoS2 (molybdenite), WO3 (wolframite), Fe2O3 (hematite and rust), FeOCr2O3 (chromite), Mn2O3H2O (manganite), etc.

Cu and the ‘noble metals’ ( Hg, Ag, Pt, Au, Pd, Ir, Ru, Re and Os) are the exceptions. They are not readily oxidized and can only be dissolved in strong mineral acids or mixtures of these acids. A number of these metals occur naturally in the elemental state, e.g., Ag, Au, Cu, etc.

Study the Electrochemical Series table on the next page and memorize the mnemonic
‘LiKe CaBaNa MAZICNTL H. CHAPA’. This is not a complete list but it is useful to the chemist and is easily remembered.

Problem 3: On a periodic table, mark all the ‘noble metals’ to learn where they are located.

Problem 4: Write the formula of

a)the most common chloride salt of lanthanum

b)the only sulfide of zinc

c)the most common phosphate of mercury.

Problem 5: Write out the electrochemical series in order of decreasing reactivity as metals.

Problem 6: State whether the following reactions will occur spontaneously at room temperature. If the reaction does occur, write a balanced equation for the reaction. If the reaction is not spontaneous, state ‘no reaction’.

a)Zn reacts with water to evolve H2 gas.

b)Hg reacts with cold dilute HNO3

c)Ba reacts with water

d)Ca reacts with concentrated HCl

e)Cd reacts with dilute H2SO4

f)Sc reacts with Cr+2 in solution

g)Ti+2 reacts with V in solution

h)Cu+2 reacts with Ni in solution

ACITIVITY SERIES OF METALS

The reduction potentials of various metals ions vs. the H2 half cell are listed. The order of this series can be easily remembered using the mnemonic ‘LiKe CaBaNa MAZICNTL H. CHAPA’

These metals are powerful reducing agents which react even with HOH to liberate H2 gas

These metals are moderate reducing reagents which react with dilute mineral acids (HCl, H2SO4, HNO3, etc.) and liberate H2 gas.

These ‘noble’ metals are poor reducing agents and do not react with dilute mineral acids. They are weaker reducing agents than H2 (g)

Ion / Metal / Reduction Potential (V)
Li+ (aq) / + / 1 e- /  / Li (s) / - 3.0
K+ (aq) / + / 1 e- /  / K (s) / -2.92
Ca+2 (aq) / + / 2 e- /  / Ca (s) / -2.90
Ba+2 (aq) / + / 2 e- /  / Ba (s) / -2.87
Na+ (aq) / + / 1 e- /  / Na (s) / -2.71
Mg+2 (aq) / + / 2 e- /  / Mg (s) / -2.37
Al+3 (aq) / + / 3 e- /  / Al (s) / -1.66
Zn+2 (aq) / + / 2 e- /  / Zn (s) / -0.763
Fe+2 (aq) / + / 2 e- /  / Fe (s) / -0.440
Cd+2 (aq) / + / 2 e- /  / Cd (s) / -0.403
Ni+2 (aq) / + / 2 e- /  / Ni (s) / -0.28
Sn+2 (aq) / + / 2 e- /  / Sn (s) / -0.136
Pb+2 (aq) / + / 2 e- /  / Pb (s) / -0.126
2 H+ (aq) / + / 2 e- /  / H2 (g) / 0.00
Cu+2 (aq) / + / 2 e- /  / Cu (s) / +0.337
Hg+2 (aq) / + / 2 e- /  / Hg (s) / +0.778
Ag+1 (aq) / + / 1 e- /  / Ag (s) / +0.799
Pt+2 (aq) / + / 2 e- /  / Pt (s) / +1.2
Au+1 (aq) / + / 1 e- /  / Au (s) / +1.68

• A ‘+’ voltage indicates that a reaction is favorable (compared to hydrogen) in the direction shown.
  A ‘-’ voltage indicates that a reaction is unfavorable but its reverse reaction is favorable.
• The voltages are standard potentials (Eo) which would be measured at standard conditions, i.e.,
  25 C, 1 M concentration for aq. ions, and 1 atm. pressure for gases (as per the Nernst equation).
• These reactions are referred to as ‘half-cells’ because each is only half of a reaction. Two half cells must be combined in order for a chemical reaction to occur. As reduction occurs, oxidation must also be occurring simultaneously. As one chemical gains electrons, another must be providing (losing) them.
• Add two half cell potentials (one for oxidation and one for reduction). If the combined voltage is a ‘+’ value, we can expect the reaction to proceed as written.

Ammine Complexes:

Most representative metal hydroxides and transition metal hydroxides are insoluble in water.

e.g.,Mg+2 + 2 OH-  Mg(OH)2

Exceptions are the Group 1A metals and Sr(OH)2 and Ba(OH)2.

Some insoluble metal hydroxides are amphoteric; i.e., in addition to acting as bases, they can act as acids dissolving in an excess of strong base.

e.g., Al(OH)3 + NaOH  [Al(OH)4]- (a soluble hydroxo complex).

Ammonia is a weak base (pKb = 4.8) and would not be expected to be able to produce a high enough OH- concentration to dissolve insoluble amphoteric metal hydroxides and form soluble hydroxo complexes. However, several metal hydroxides do dissolve in an excess of aqueous ammonia to form ammine complexes.

e.g.,Cu(OH)2 + 4 NH3  [Cu(NH3)4]+2 + 2 OH-

e.g., Co(OH)2 + 6 NH3  [Cu(NH3)6]+2 + 2 OH-

Interestingly, all metal hydroxides that exhibit this behavior are derived from the 12 metals of the Co, Ni, Cu, and Zn groups. All common cations of these metals form soluble complexes in the presence of aqueous ammonia.

Co(OH)2 + 6 NH3  [Co(NH3)6]+2 + 2 OH-

Co(OH)3 + 6 NH3  [Co(NH3)6]+3 + 3 OH-

Ni(OH)2 + 6 NH3  [Ni(NH3)6]+2 + 2 OH-

CuOH  + 2 NH3  [Cu(NH3)2]+1 + OH-

Cu(OH)2 + 4 NH3  [Cu(NH3)4]+2 + 2 OH-

AgOH  + 2 NH3  [Ag(NH3)2]+1 + OH-

Zn(OH)2 + 4 NH3  [Zn(NH3)4]+2 + 2 OH-

Cd(OH)2 + 4 NH3  [Cd(NH3)4]+2 + 2 OH-

Hg(OH)2 + 4 NH3  [Hg(NH3)4]+2 + 2 OH-

CuOH and AgOH are unstable compounds and decompose to Cu2O and Ag2O, but in the presence of aqueous NH3 they dissolve as shown above.

Problem 7: Write a balanced equation for the reaction of 6 moles NH3 with each of the following

a)Au(OH)3

b)Pt(OH)4

Colors of Transition Metal Complexes:

Unlike compounds of the representative elements, compounds of transition metals are colored. Why?

Colors of transition metal complexes arise from absorption of visible light as electrons undergo transitions from one d-orbital to another. In their elemental forms, all five of the d-orbitals of transition metals are degenerate (of the same energy) but in their complexes this changes.

Review the shapes and orientations of the five d-orbitals. In octahedral complexes, for example, electrons in the dZ2 and dX2-Y2 orbitals are repelled by the donor electrons of the six ligands. d-orbital electrons resist occupying the same region of space as the ligands’ donor electrons. As a result, the dZ2 and dX2-Y2 orbitals are higher in energy than the dXY, dXZ and dYZ orbitals.

The d-orbitals are split into sets of orbitals separated by energies corresponding to wavelengths of electromagnetic radiation in the visible region (400 to 800 m).

d-orbitals are split in other ways in other complexes such as tetrahedral, square planar complexes.

The magnitude of the field splitting energy depends upon the electron donating ability (field strength) of the ligands. The frequency and wavelength of light absorbed (i.e., the color) are related to the size of the field splitting energy, which in turn depends upon the electron donating ability (field strength) of the ligands. Recall the relative field strength of various ligands.

I- < Br- < Cl- < F- < OH- < C2O4-2 < H2O < SCN- NH3 < en < NO2- < CN- < CO

A Cu+1 ion has a d10 electron configuration and all d-orbitals are filled. In order for absorption of electromagnetic radiation to occur, a d-electron must be promoted to a 4p orbital (the next highest orbital with room). Because the energy level of the 4p orbitals is much higher that of the 3d orbitals, photons of a very high energy are needed, i.e., ultraviolet radiation. No visible light is absorbed so Cu+1 complexes, such as Cu(CN)2-, are colorless.

A Cu+2 ion has a d9 configuration with a vacancy in one of its eg orbitals. Visible light has sufficient energy to excite a t2g electron to the vacant position in the eg orbitals. Many Cu+2 complexes are colored. Some examples of colored Cu+2 complexes follow.

Complex

/  of Light
Absorbed / Color of Light
Absorbed / Color of Light Transmitted
[Cu(H2O)4]+2 / 610 m / orange / blue
[Cu(en)2]+2 / 430 m / purple / green
[CuBr4]-2 / 530 m / green / purple

When certain visible wavelengths are absorbed from incoming white light, the light not absorbed is transmitted or reflected and has the complementary color of the light absorbed. The complementary color is the color seen by an observer.

Complementary Colors:

Very simply, blue and yellow are complementary; red and cyan (blue-green) are complementary; green and magenta (blue-red) are complementary. A more complete list follows.

 (m) / Spectral
Color / Complementary Color
410 / violet / lemon yellow
430 / indigo / yellow
480 / blue / orange
500 / blue-green / red
530 / green / purple
560 / lemon yellow / violet
580 / yellow / indigo
610 / orange / blue
680 / red / blue-green

The color of aqueous solutions of some transition metal nitrates compared to those of representative metal nitrates are listed.

Representative Metal Nitrates / Color of Aqueous Solution / Transition Metal Nitrates / Color of Aqueous Solution
Na+ / colorless / Cr+3 / deep blue
Ca+2 / colorless / Mn+2 / pale pink
Mg+2 / colorless / Fe+2 / pale green
Al+3 / colorless / Fe+3 / orchid (rust)
Sn+2 / colorless / Co+2 / pink
Sn+4 / colorless / Ni+2 / green
Pb+2 / colorless / Cu+2 / blue

Absorption Spectrum:

The amount of light absorbed by a sample as a function of wavelength is known as its absorption spectrum. The visible absorption spectrum of clear (not colorless) samples can be determined by scanning with a visible spectrophotometer. This is useful in the analysis of transition metal complexes. The wavelength of maximum absorption max is related to the field splitting energy of a complex.

In the example below, the absorption spectrum of [Ti(H2O)6]+3 shows a max in the visible region at 530 m (green) and so the complex appears purple to the eye.

Problem 8: The presence of partially filled d-orbitals is usually necessary for color.

Explain why K+ and Sr+2 ions are usually colorless

Explain why Al+3 and Sn+4 ions are usually colorless.

Problem 9: Aqueous solutions of CrCl2 are blue, CrCl3 are green, K2CrO4 are yellow and K2Cr2O7 are orange. Explain why these solutions are different colors.

Problem 10: In the discussion of colorless Cu+1 ion complexes on page 9, no mention was made of the fact that these complexes have empty 4s orbitals that could accommodate excited electrons from the 3d orbitals. This being the case, propose a reason for that fact that such transitions do not give rise to absorption in the visible spectrum. What type of the electromagnetic radiation would you expect to be absorbed in such a transition?

Problem 11: Answer the following questions with respect to the three transition metal complexes:
[Co(H2O)6]+3, [CoF6]-3, [Co(NH3)6]+3

a)List the complexes in order of increasing field splitting energy (Oct).

b)The wavelengths of light absorbed by these are 475, 600, and 900 m. Predict which wavelength corresponds to each compound.

c)Predict the apparent (observed) color of each complex.

Ferromagnetism: