transition metal chemistry - chem hogchemrat.com/chemhog2/general...

TRANSITION METAL CHEMISTRYThe 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 4s239Y [Kr] 4d1 5s2

57La [Xe] 5d1 6s2

22Ti [Ar] 3d2 4s240Zr [Kr] 4d2 5s2

72Hf [Xe] 4f14 5d2 6s2

23V [Ar] 3d3 4s241Nb [Kr] 4d4 5s1

73Ta [Xe] 4f14 5d3 6s2

24Cr [Ar] 3d5 4s142Mo [Kr] 4d5 5s1

74W [Xe] 4f14 5d4 6s2

25Mn [Ar] 3d5 4s243Tc [Kr] 4d5 5s2

75Re [Xe] 4f14 5d5 6s2

26Fe [Ar] 3d6 4s244Ru [Kr] 4d7 5s1

76Os [Xe] 4f14 5d6 6s2

27Co [Ar] 3d7 4s245Rh [Kr] 4d8 5s1

77Ir [Xe] 4f14 5d7 6s2

28Ni [Ar] 3d8 4s246Pd [Kr] 4d10 5s0

78Pt [Xe] 4f14 5d9 6s1

29Cu [Ar] 3d10 4s147Ag [Kr] 4d10 5s1

79Au [Xe] 4f14 5d10 6s1

30Zn [Ar] 3d10 4s248Cd [Kr] 4d10 5s2

80Hg [Xe] 4f14 5d10 6s2

bb/CH306/Transitionmetals.doc 1

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).


METALS mp vs. Group Number

Cr

Cu

Ga

Ge

Zr

Mo

Sn

Ni

Ca

K

Zn

Ti Fe

CoMn

VSc Pd

Ru

RhTc

Nb

Sr

RbIn

Y

Ag

Cd

LaLa

Hg

Pt

Ir

Os

W

Ta

AuBa

CsPb

Re

Hf

-500

0

500

1000

1500

2000

2500

3000

3500

4000

0 2 4 6 8 10 12 14 16

Group Number

mp

(°C

)

Period 4

Period 5

Period 6

1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A

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 19.3Ir 22.4 Au 19.3Pt 21.5 Hg 13.5Re 20.8 Pb 11.3

Fe = 7.9 g/mL

2. 3rd transition series metals have unusually high ionization energies, i.e., are rather unreactive.

3. 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.


1A 2A 3B 4B 5B 6B 7B 8B 8B 8B 1B 2B 3A 4A

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

Basic Amphoteric Acidic

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 and HNO3 > 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 O2 4 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 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

bb/CH306/Transitionmetals.doc

Group 8B

Fe Co Ni

Ru Rh Pd

Os Ir Pt

7

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 2BSc 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 2BY 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 -2.08Ti Ti+2 + 2e- Ti -2.00


Thes

e m

etal

s (e

xcep

t Cu)

are

st

rong

er re

duci

ng a

gent

s

than

H2.

V V+2 + 2e- V -0.26Cr Cr+2 + 2e- Cr -0.56Mn Mn+2 + 2e- Mn -1.02Fe Fe+2 + 2e- Fe -0.41Co Co+2 + 2e- Co -0.28Ni Ni+2 + 2e- Ni -0.23Cu Cu+2 + 2e- Cu +0.52Zn 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), FeOCr2O3 (chromite), Mn2O3H2O (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’

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.

bb/CH306/Transitionmetals.doc

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)

11

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


Increasing ability to donate an electron pair to a metal ion STRONGLIGANDS

WEAKLIGANDS

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 dZ

2 and dX2-Y

2 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 dZ

2 and dX

2-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.


dXY dXZ dYZ dX2-Y

2 dZ2

dX2-Y

2 dZ2

dXY dXZ dYZ

Field Splitting Energy(Oct)

Free metal ion orbital energies

Metal ion orbital energies in an octahedral field.

eg

t2g

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.


300 400 500 600 700 800

max = 530 m

(m)

Abs

.

Ferromagnetism:All electrons spin and since all moving electric charges produce magnetic fields, electrons behave as tiny magnets. When electrons are paired in orbitals, their spins are opposite (lowest energy state) and so their magnetic fields cancel. Substances with only paired electrons are either nonmagnetic or ‘diamagnetic’. Diamagnetic objects, when placed in a magnetic field, are not attracted, in fact they are weakly repelled. Examples of diamagnetic elements include Zn and Cd.

When an atom possesses one or more unpaired electrons, the substance is said to be ‘paramagnetic’. Paramagnetic substances contain many weak atomic magnets, but in the solid phase, the magnetic fields are not aligned; instead they are randomly oriented. When placed in a magnetic field, alignment of electron spins with the external magnetic field occurs to some extent throughout the various regions (domains) in the solid. As a result, such paramagnetic materials are weakly attracted to external magnetic fields. Examples of paramagnetic substances include Al and O2.

Some metals exhibit strong magnetic properties. These materials are called ‘ferromagnetic’. They are typically 106 times stronger than paramagnetic materials. Fe, Co, and Ni are the only three elements that are ferromagnetic. They have 4, 3, and 2 unpaired d-electrons, respectively. In ferromagnetic materials, after the external magnetic field is removed, the ferromagnetic material may remain magnetized as its electron spins remain aligned. It becomes a permanent magnet. Permanent magnets can become demagnetized by heating, pounding, or melting.

One might ask, Why is Mn, with 5 unpaired electrons, not ferromagnetic?

Ferromagnetism is a solid state property. When individual paramagnetic atoms are correctly spaced their magnetic fields readily align. In Mn, interatomic distance is incorrect to allow optimal magnetic field alignment. Adding proper amounts of Cu to Mn adjusts the interatomic spacing such that this alloy becomes ferromagnetic. Some naturally occurring metal oxides are ferromagnetic, e.g., Fe3O4 (magnetite) and CrO2 are both used in magnetic recording tapes.


unpairedelectrons

3d0 3d1 3d2 3d3 3d5 3d5 3d6 3d7 3d8 3d9 3d10

K+ Ti+3 V+3 V+2 Cr+2 Mn+2 Fe+2 Co+2 Ni+2 Cu+2 Cu+1 Ca+2 V+4 Cr+3 Mn+3 Fe+3 Zn+2 Ti+4 Ga+3

5

4

3

2

1

0

The Inner Transition Metals (f-Transition Metals):All inner transition elements, are metals in which f-orbitals are being filled. These f-orbitals are two shells inside the outermost occupied shell being filled.

In the Lanthanides (58Ce 71Lu), the 4f orbitals being filled are within the 6s and 5d orbitals.

In the Actinides (90Th 103Lr), the 5f orbitals being filled are within the 7s and 6d orbitals.

Note that the lanthanides do not include lanthanum and the actinides do not include actinium.

The chemical properties of these elements are governed by their valence electrons (in their outermost shells). Each rare earth element has two s-electrons in its outermost shell and either eight or nine electrons in the next shell inward (two s, six p, and zero or one d-electron). This accounts for the chemical similarity of these elements. Most form M+3 ions by loss of the two outer s-electrons and either the d-electron one shell inward or one of the f-electrons two shells inwards.

Electron Configuration of the Lanthanides:

Atomic Number

Name Symbol Electron Configuration

Formula of Chloride

Ox. State of Lanthanide

57* lanthanum* La [Xe] 5d1 6s2 LaCl3 +3

58 cerium Ce [Xe] 4f1 5d1 6s2 CeCl3, CeCl4 +3, +4

59 praseodymium Pr [Xe] 4f3 6s2 PrCl3, PrCl4 +3, +4

60 neodymium Nd [Xe] 4f4 6s2 NdCl3 +3

61 promethium Pm [Xe] 4f5 6s2 PmCl3 +3

62 samarium Sm [Xe] 4f6 6s2 SmCl2, SmCl3 +2, +3

63 europium Eu [Xe] 4f7 6s2 EuCl2, EuCl3 +2, +3

64 gadolinium Gd [Xe] 4f7 5d1 6s2 GdCl3 +3

65 terbium Tb [Xe] 4f9 6s2 TbCl3, TbCl4 +3, +4

66 dysprosium Dy [Xe] 4f10 6s2 DyCl3 +3

67 holmium Ho [Xe] 4f11 6s2 HoCl3 +3

68 erbium Er [Xe] 4f12 6s2 ErCl3 +3

69 thulium Tm [Xe] 4f13 6s2 TmCl3 +3

70 ytterbium Yb [Xe] 4f14 6s2 YbCl2, YbCl3 +2, +3

71 lutetium Lu [Xe] 4f14 5d1 6s2 LuCl3 +3

* La is not an f-transition element and is only shown here to emphasize the filling order

Empty (f0), half (f7) and full (f14) orbitals are favored electron configurations of the metals and their ions, e.g., Ce+4, Gd+3, Lu+3.

Whereas the d-transition metals show a strong tendency to form coordination compounds, the f-transition metals form few such compounds.


The Lanthanides (Rare Earths):The lanthanides are not as rare as the name ‘rare earths’ implies. Ce and Nd are more abundant than Pb. Tm is about as prevalent as iodine. Mixtures of lanthanides and actinides occur in minerals such as monazite and gadolinite. However, these elements are difficult to obtain in pure form because of their chemical similarities. They are usually separated chromatographically based on differences in attraction of these similar compounds in solution for ionic sites on insoluble resins.

The standard reduction potentials (E° values) of the lanthanides for M+3 (aq) + 3e- M (s) all fall between –2.52 V (La) and –2.25 V (Lu). Thus they are all active metals liberating H2 from hot water and dilute acids while undergoing oxidation to M+3 ions.

M(s) + 3H2O(l) M(OH)3(s) + 3/2 H2(g)

M(s) + 3HCl MCl3 + 3/2 H2(g)

The lanthanides react with chlorine forming salts of the general formula MCl3.

2M(s) + 3Cl2 2MCl3(s)

All lanthanides are silvery, high melting (800 to 1600°C) metals. They oxidize readily forming M2O3 oxides and MX3 halides. Their hydroxides are not very soluble and are basic. Their carbonates, phosphates and fluorides are only slightly soluble. In all these properties (except oxidation state) they resemble the alkaline earth metals (Group 2A).

Actinides:The actinides are all radioactive. Uranium (Z = 92) is the upper limit of naturally occurring elements. Of the 14 actinides, only U and Th are found in appreciable quantities in nature. All elements of higher atomic mass (the ‘transuranium’ elements) have only been synthesized in small amounts in particle accelerators. Macroscopic samples of mendelevium (Md), nobelium (No) and lawrencium (Lr) have never been seen.

Actinides are silvery and chemically reactive. Like the lanthanides, they form highly colored compounds.

Some Applications of Lanthanides and Actinides:A cerium salt, e.g., Ce(SO4)3, is a common oxidizing titrant in titrimetric analysis.

Several of the rare earths have some commercial importance. Praseodymium and neodymium are used to tint sunglasses and welding/glassblower’s goggles. Eu2O3, Y2O3, and Gd2O3 are phosphors used in the fluorescent screens of color televisions. Thorium nitrate has been used for over a century in gas mantles of lanterns and street lamps.

Uranium and Plutonium are used as fuel elements in nuclear reactors and nuclear weapons.

To extract uranium, its oxide ores are treated with nitric acid. The nitrate that forms is reduced to metal with Ca. To enrich uranium in its fissionable 235U isotope for use in nuclear power plants, the oxide is converted to UF4 with HF. The UF4 is oxidized to the volatile UF6 by F2. The vapor of UF6, which contains both 238U and 235U, is then subjected to repeated diffusion through porous barriers to concentrate the 235UF6 (Graham’s Law).

SmCo5 alloy is the strongest known permanent magnet.


SELF STUDY QUESTIONS FOR TRANSITION METALS1. Describe each of the following types of compounds as acidic, basic, or amphoteric

a) metal oxides

b) non metal oxides

c) oxides of metalloids

2. Collectively, the f-transition metals are also referred to by two other names. List both…………………………………………… and …………………………………………….

3. List 5 general (characteristic) properties of transition metals (aside from the fact that they are metals):

a) …………………………………………………………………………………..

b) …………………………………………………………………………………..

c) …………………………………………………………………………………..

d) …………………………………………………………………………………..

e) …………………………………………………………………………………..

4. The melting point of transition metals is usually not below ………………….. °C.

5. The highest melting transition metal is ……………………….

6. The highest melting elements of all is ……………………….. and its melting point is ……………………………. °C.

7. The highest melting element in all transition metal series occurs in Group ………. (see the graph on page 2).

8. The lowest melting transition metal is …………… .

9. Circle the correct answer. Atomic size increases or decreases left to right from Groups 3B to Group 8B across each transition metal series.

10. Rank 1st transition metal series, 2nd transition metal series and 3rd transition metal series in order of

a) increasing melting point (1= lowest) (see graph on page 2)

b) increasing density (1 = lowest) (see graph on page 3)

11. Explain what is meant by the Lanthanide contraction.

12. List 3 results of the Lanthanide contraction.

13. List the following densities to the nearest whole number (g/mL): Os ……….. , Hg …………, Pb ……….., Fe ………., Li (to 1 decimal point) ……….. . (Check a periodic table for the values not listed in the notes.)

14. Given the formulas of a series of oxides of transition metals, rank the oxides in order of increasing acidity, where 1 = least acidic. For example, rank the following in order of increasing acidity: MnO2 Mn2O7 Mn2O3 MnO MnO3

15. What is the principle behind the trend in question 14? Make a general statement that describes the trend and that could be applied to other series of oxides of a transition metal.

16. Given a group of analogous oxyacids, circle the most acidic and draw a square around the least acidic, e.g.,

a) H3PO4 H2SeO4 H2SO4 H4GeO4 (same number of oxygens)

b) H3PO4 H3PO2 H3PO3 (different number of oxygens)

c) HClO HClO4 HClO2 HClO3 (different number of oxygens)


17. What are the two principles behind the trends in question 16? Make two general statements, one for each of the two trends. These trends should be applicable to other groups of analogous oxyacids.

18. Transition metals are catalysts in many diverse types of chemical reactions (see page 5). What characteristic of transition metals makes this possible? (Hint: read page 7).

19. Distinguish between homogenous and heterogeneous catalysis and give an example of each.

20. List 4 transition metals that are often used as catalysts for hydrogenation of unsaturated organics.

21. State the relationship between B-Group number and oxidation state of metals in the group.

22. No simple (uncomplexed) transition metal possess a charge greater than ………. .

23. Predict the formulas of transition metal compounds. (Hint: use group numbers and oxidation numbers, see page 6). For example, give the formula of:

a) the chloride of cadmium

b) the phosphate of silver

c) the sulfate of scandium

d) the carbonate of zirconium

e) permanganic acid (you will note the similarity to Grp 7A oxyacids)

f) chromic acid (you will note the similarity to Grp 6A oxyacids)

24. Name the 3 horizontal triads of Group 8B metals

25. List 3 elements that are ferromagnetic.

26. State the difference between the water soluble ions of the 1st transition series metals compared to those of the 2nd and 3rd transition series and give one example of each to illustrate. (see page 7).

27. Circle the correct answer. Most transition metals are readily oxidized, readily reduced, or unreactive in dilute mineral acids such as HCl, H2SO4, etc.

28. List the formula of 4 mineral acids. (These are not listed in these notes, however, mineral refers to the fact that they are derived from minerals, i.e., the earth, not organic)

29. List the chemical symbols of 5 noble metals.

30. Describe the chemical reactivity of noble metals.

31. You will need to memorize the mnemonic on the top of page 9, i.e., LiKe CaBaNa, MAZINCTL H. CHAPA and then answer questions like those in problem 6 on page 8.

32. Circle the correct answer: A negative or positive reduction potential (voltage) indicates that a reaction is spontaneous as written, compared to hydrogen.

33. Most A-group elements hydroxides are insoluble. List the formulas of all A-group element hydroxides that are water soluble.

34. Write the formula of an amphoteric hydroxide of an A-group element.

35. All transition metal hydroxides from four of the B-groups dissolve in aqueous ammonia. List the symbols of the metals at the top of each of the four groups.

36. Write an balanced equation for the reaction of

a) zinc hydroxide with 4 moles NH3

b) cobalt(III) hydroxide with 6 moles of NH3

c) etc. (see page 10).

37. The five degenerate d-orbitals are split into two energy levels in octahedral complexes. Draw a diagram showing this splitting and indicate which orbitals are each level.


38. Explain what is meant by the term ‘field splitting energy’ in an octahedral complex.

39. Draw the electron configuration of Cu+ and Cu+2 in an octahedral complex and explain why Cu+2 is more strongly colored than Cu+. The same rationale applies to other transition metal ions so you should be able to do this for other transition metal ions as well.

40. The magnitude of the field splitting energy in a transition metal complex depends upon what characteristic of the ligands?

41. Write the formula of one ligand that is a weak donor (………….), one that is a strong donor (……..)

42. CuSO45H2O is bright blue in appearance. What color of light is absorbed by this compound?

43. K2Cr2O7 is a dark orange crystalline solid. What color of light is absorbed by this compound?

44. KMnO4 forms dark purple solutions. What color of light is absorbed by this compound? You are not expected to memorize all the complementary colors. These three should be enough to illustrate the concept.

45. Explain what is meant by the term max in an absorption spectrum.

46. Explain why K+ and Sr+2 are usually colorless.

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

48. Answer question 11 on page 13. From our previous study of Coordination Complexes, it is assumed that you recall that NH3 is a strong donor ligand, followed by H2O. F is a weak donor ligand. It is also assumed that you recall the visible spectrum runs from 400 to 800 m, 400 being higher energy (violet light) and 800+ being lower energy (red light).

49. Describe how the following materials behave when placed in a magnetic field…

a) diamagnetic

b) paramagnetic

c) ferromagnetic

50. Write the formulas of 3 ferromagnetic elements and 2 ferromagnetic metal oxides

51. For the f-transition metals, list the most common oxidation state ………….. (see table p. 15)

52. How many actinides are there?

53. The standard reduction potential of all lanthanides fall between …….. and ……. V (answer to 1 decimal place and include the sign, + or -)

54. Write a balanced chemical equation for the reaction of a lanthanide in hot water. (hint recall the most common oxidation state of lanthanides).

55. How many actinides are there?

56. How many of the actinides are radioactive?

57. Give the symbol and atomic number of the heaviest (largest) naturally occurring element.

58. Uranium and plutonium are used primarily for what purposes?


CHEMISTRY AND APPLICATIONS OF SELECTED TRANSITION METALS

Group 6B: Chromium, Molybdenum and Tungsten:Chromium: [Ar] 3d5 4s1 The chief ore is chromite, FeCr2O3. The element is industrially produced by reduction of this oxide with Al. Chromium is a silvery, hard, brittle metal (mp = 1860°C). It is extremely resistant to ordinary corrosive agents, which accounts for its extensive use as an electroplated protective coating. The metal dissolves in nonoxidizing mineral acids such as hydrochloric or sulfuric, but not in aqua regia (a 3:1 mixture of HCl:HNO3) or concentrated or dilute cold nitric acid. These later reagents ‘passivate’ the metal making it more resistant to corrosion.

Chromium is mixed with iron to make various stainless steels (10 to 25% Cr). Nickel is sometimes also used. Stainless steels are hard, dense, corrosion resistant steels.

Chromium’s most common oxidation states are +2, +3 and +6. The +3 state is most stable.

Cr+2 + 2e- Cr E° = -0.56

Cr+3 + e- Cr+2 E° = -0.41

Chromium dissolves slowly in dilute HCl or dilute H2SO4 liberating H2 and forming a blue solution of chromous ion.

Cr(s) + 2H+(aq) Cr+2(aq) + H2(g)

Normally, the blue color is not observed because the Cr+2 ion is rapidly oxidized to violet Cr+3.

4Cr+2(aq) + O2(g) + 4H+(aq) 4Cr+3(aq) + 2H2O(l)

When HCl solution is used, the solution appears green as a result of the formation of complex ions containing chloride coordinated to chromium, such as Cr(H2O)4Cl2+. Note that all Cr+3 complexes are hexacoordinate.

Chromium(III) hydroxide is amphoteric.

Cr(OH)3(s) + 3H+ Cr+3 + 3H2O

Cr(OH)3(s) + OH- Cr(OH)4-

Chromium is also encountered in aqueous solution in the +6 oxidation state. In basic solution, the yellow chromate ion, CrO4

-2, is the most stable form. In acidic solution, the dichromate ion, Cr2O7

-2 is formed.

CrO4-2(aq) + H+(aq) HCrO4

-(aq)

2HCrO4-(aq) Cr2O7

-2(aq) + H2O(l)

The equilibrium between these ions is easily observable because CrO4-2 is bright yellow and

Cr2O7-2 is deep orange.

The dichromate ion in acidic solution is a strong oxidant but chromate ion in basic solution is a weak oxidant.

Cr2O7-2 + 14H+ + 6e- 2Cr+3 + 7H2O E° = +1.33 V

CrO4- + 4H2O + 3e- Cr(OH)3 + 5OH- E° = -0.12 V


Dehydration of chromate or dichromate salts with concentrated H2SO4 produces chromium trioxide (CrO3), also called chromia. Chromium(VI) oxide is a strong oxidant. A powerful ‘cleaning solution’ used for removing greasy stains and coatings from laboratory glassware is made by adding concentrated H2SO4 to a solution of K2Cr2O7. The active ingredients are CrO3, an oxidant, and H2SO4, an excellent solvent. This cleaning solution should be handled with caution because it is corrosive and chromium(VI) compounds are toxic and carcinogenic.

Chromium(VI) is commonly used in electroplating operations. Removal of chromium from industrial wastewater is required by law. A common method is reduction of Cr+6 with Na2SO3 (sodium sulfite) or NaHSO3 (sodium bisulfite) in acidic solution. The Cr+3 is then precipitated as the insoluble hydroxide [Cr(OH)3], (Ksp = 6.7 10-31) by raising the pH to the point where the hydroxide is least soluble (pH 8 to 9). Recall that Cr(OH)3 is amphoteric and its solubility increases in strongly basic solution.

Problem: Draw the Lewis structures of chromate and dichromate ions. Note that Lewis structures show all valence (outer shell) electrons. For Lewis structures of B Group elements the outermost s-electrons and d-electrons are considered as valence electrons and are shown in Lewis structures.

Problem: Calculate the residual concentration of Cr+3 (mg/L as Cr) remaining after an acidic 2.0 M solution of Cr+3 is neutralized with NaOH. Cr(OH)3 Ksp = 6.7 10-31.

Molybdenum and Tungsten:Other members of Group 6B are Mo ([Kr] 4d5 5s1) and W ([Xe] 4f14 5d4 6s2).

Molybdenum occurs chiefly as molybdenite, MoS2 but also as molybdates such as PbMoO4 and MgMoO4.

Tungsten is found almost exclusively as tungstates, i.e., wolframite (FeWO4+ MnWO4), scheelite (CaWO4) and stozite (PbWO4).

The elemental forms of both are silvery and extremely ‘refractory’ (high melting); Mo melts at 2620°C and W melts at 3420°C. Neither metal is attacked even by concentrated acids. Concentrated nitric acid passivates the surface of Mo.

The chief uses of both metals are in the production of alloy steels; even small amounts cause tremendous increases in hardness and strength. ‘High speed’ steels, containing W and Cr, are used to make cutting tools that remain hard even at high heat. Tungsten is also extensively used as filaments for incandescent light bulbs.

Both Mo and W exist in multiple oxidation states (+2, +3, +4, +5, and +6). The +6 state is most stable.

MoO3 and WO3 are more thermally stable, less acidic and weaker oxidants than CrO3.

Both Mo+6 and W+6 give tetrahedral oxo anions, MO4-2, whose alkali metal salts are water

soluble, e.g., Na2MoO4 and (NH4)2WO4. These anions polymerize on acidification giving rise to a very complicated series of isopolyanions and in the presence of other elements heteropolyanions. Some examples follow.


Isopolymolybdates:

When a basic solution of MoO4-2 is acidified, the molybdate ions condense in definite steps to

form a series of polymolybdate ions.

7MoO4-2 + 8H+ 4H2O + Mo7O24

-6 (‘paramolybdate’ ion)

In more acidic solution octamolybdate ion [Mo8O26]-4 is formed.

Isopolytungstates:

Like Mo, as the pH is lowered, tungstate (WO4-2) in solution polymerizes to various

polytungstates, the most important being ‘paratungstate’, W12O41-10.

These isopolyanions combine with other oxo anions, particularly those of orthophosphate (PO4

-3) and orthosilicate (SiO4-4) forming brightly colored heteropolyanions.

Heteropolymolybdates

One important use of heteropolymolybdates is in the spectrophotometric determination of phosphates and silicates in waste water and in industrial cleaning and electroplating solutions.

Phosphorus (as orthophosphate) can be quantitatively measured at concentrations below 1ppm. Ammonium molybdate, (NH4)6Mo7O244H2O is added to a sample containing phosphate. When the pH is lowered to ca. 0.7, yellow 12-molybdophosphoric acid polymer forms (H3PMo12O40). This is reduced with SnCl2, for example, to a blue complex called ‘molybdenum blue’ and read on a spectrophotometer at 660 m.

Silica (as orthosilicate) can be quantitatively measured at concentrations near 1ppm. Ammonium molybdate, (NH4)6Mo7O244H2O is added to a sample containing silicate. When the pH is lowered to ca. 1.2, yellow 12-molybdosilicic acid polymer forms (H4SiMo12O40). This is read spectrophotometrically at ca. 300 m.

Alternately, the yellow 12-molybdosilicic acid can be reduced to a ‘molybdenum blue’ complex and measured down to 0.2ppm at 815m.


Group 7B (Mn, Tc and Re)Manganese:Mn [Ar] 3d5 4s2

Manganese is relatively abundant, constituting about 0.085% of the earth’s crust. Among the heavy metals, only Fe is more abundant. It occurs naturally as oxides or carbonates. It also occurs over wide expanses of the Pacific seabed as nodules (combined with Ni, Cu and Co) ranging in size from a few millimeters to a few meters in diameter. Billions of tons are present but mining them presents a major technical and political challenge.

The metal is obtained from the oxide by reduction with Al.

Elemental manganese is hard and shiny and, like vanadium and chromium, is used mostly to make steel alloys. A small amount of Mn (< 1%) makes steel easier to roll, forge, and weld. Steel blades made with 12% Mn is tough enough to be used for naval armor and bulldozer blades. Small amounts of Mn are added to aluminum beverage cans and bronze alloys to make them stiffer and tougher.

The chemistry of Mn resembles that of Cr in some respects. The free metal is reactive and readily displaces H2 from acids to form pale-pink Mn+2 ion:

Mn(s) + 2H+(aq) Mn+2(aq) + H2(g) E° = + 1.1.8 V

Like Cr, Mn can use all its valence electrons in its compounds and exhibits all oxidation states from +1 through +7, with +2, +4. The +7 state is most common but only occurs in oxo compounds.

As the oxidation state of Mn rises, its electronegativity rises and its oxides change from basic to acidic. Manganese(II) oxide (MnO) is basic, manganese (III) oxide (Mn2O3) is amphoteric. Manganese(IV) oxide (MnO2) is insoluble, so it shows no acid-base properties; it is used in dry-cell and alkaline batteries as the oxidizing agent in a redox reaction with Zn. Manganese(VII) oxide (Mn2O7) reacts with water to form permanganic acid (HMnO4), which is as strong as perchloric acid (HClO4).

All manganese species with oxidation states greater than +2 act as oxidants, but the purple permanganate ion is particularly powerful. As with Cr in its highest oxidation state, MnO4

- is a much stronger oxidant in acidic than in basic solution.

MnO4- (aq) + 2H2O (aq) + 3e- MnO2(s) + 4OH- E° = + 0.58 V

MnO4- (aq) + 4H+(aq) + 3e- MnO2(s) + 2H2O E° = + 1.68 V

MnO4- (aq) + 8H+(aq) + 5e- Mn+2(aq) + 4H2O E° = + 1.49 V

KMnO4 is useful in the laboratory as a strong oxidant, which in strongly acidic media is reduced to colorless Mn+2 ion and in moderately acidic or basic solution is reduced to a brown precipitate (MnO2). The color change and formation of the precipitate are good indicators of the occurrence of a chemical reaction in qualitative testing.

Unlike Cr+2 and Fe+2, the Mn+2 ion resists oxidation in air. Removing an electron from the Mn+2 ion is difficult because it disrupts the stable d5 configuration.

Problem: Draw Lewis structures of MnO4- and dimanganese heptoxide (Mn2O7).


Technetium (Tc) and Rhenium (Re):Tc [Kr] 4d5 5s2 and Re [Xe] 4f145d5 6s2 resemble each other closely in chemistry but differ noticeably from manganese, although there are similarities in stoichiometries, e.g., MnO4

- and ReO4

-.

There are no stable isotopes of Tc. It occurs naturally in trace quantities as short-lived products of radioactive decay of uranium. Most of the Tc in use is produced synthetically in fission reactors.

Rhenium is in low abundance in the earth’s crust (ca. 10-90%). It is recovered in flue dusts (soot) during the refining of molybdenum sulfide ores. Pt-Re thermocouples (electric thermometers) and rhenium catalysts for hydrogenation are two of the few commercial uses of rhenium. Owing to its scarcity Re is five times more expensive than gold.


transition metal chemistry - chem hogchemrat.com/chemhog2/general...

Documents