21. TRANSITION ELEMENTS AND THEIR COORDINATION COMPOUNDS

21.1 The Transition Metals: A Survey

Transition elements make up the d-block elements. With the exception of palladium, each element has either one or two electrons in the ns orbital of the outermost shell. They differ from each other in the number of electrons in (n - 1)d orbitals, where n being the valence shell. In contrast to the main group elements, which show chemical similarities only within the group, the transition metals show great similarities within a group as well as within a given period.

Most transition metals have high melting points and high densities. They are hard and strong; most have more than one oxidation states in their common compounds, some of which possess characteristic colors. Colored polyatomic anions usually contain transition metals - for example, MnO4

- (dark purple), CrO42-

(bright yellow), and C2O72- (orange). Most have unpaired electrons in at least one of the oxidation state,

giving their compounds paramagnetic property. One significant property of transition elements is the ability of their cations to form complex ions, such as Cu(NH3)4

2+, Co(H2O)63+, Co(NH3)6

3+, Ni(CN)42-, etc., many of

which are colored.

Electron Configurations of The First Series Transition Elements and Their Ions

Sc (Z = 21): [Ar] 4s2 3d1; Sc3+ : [Ar] ;

Ti (Z = 22): [Ar] 4s2 3d2; Ti2+ : [Ar] 3d2; __ __ ___ ___ ___

V (Z = 23): [Ar] 4s2 3d3; V2+ : [Ar] 3d3; __ __ __ ___ ___

*Cr (Z = 24): [Ar] 4s1 3d5; Cr3+ : [Ar] 3d3; __ __ __ ___ ___

Mn (Z = 25): [Ar] 4s2 3d5; Mn2+ : [Ar] 3d5; __ __ __ __ __

Fe (Z = 26): [Ar] 4s2 3d6; Fe3+ : [Ar] 3d5; __ __ __ __ __

Co (Z = 27): [Ar] 4s2 3d7; Co2+ : [Ar] 3d7; __ __ __ __ __

Co (Z = 27): [Ar] 4s2 3d7; Co3+ : [Ar] 3d6; __ __ __ __ __

Ni (Z = 28): [Ar] 4s2 3d8; Ni2+ : [Ar] 3d8; __ __ __ __ __

*Cu (Z = 29): [Ar] 4s1 3d10; Cu2+ : [Ar] 3d9; __ __ __ __ __

Zn (Z = 30): [Ar] 4s2 3d10; Zn2+ : [Ar] 3d10; __ __ __ __ __

In transition elements electrons are added into the (n - 1)d orbitals, singly at first. The electron configurations of the first row transition elements differ from one another in the number of electrons in the 3d sub-shells, except in Cr and Cu, where one of the 4s electrons jumps into the 3d orbitals to form the 3d 5

4s1 configuration in chromium and the 3d10 4s1 configuration in copper.

Transition metal ions are formed through the loss of ns electrons before the (n - 1)d electrons. Thus, the electron configuration of Fe2+ is [Ar] 3d6, NOT [Ar] 4s2 3d4.

Trends in Atomic and Physical Properties of The Transition Elements

Trends Within a GroupAtomic size generally increases from top to bottom down a group; however, same group elements in the

second and third transition series appear to have the same atomic radii. This is the consequence of higher effective nuclear charge exhibited by elements in the third transition series (6 th period). Transition metals in the 6th period come after the lanthanide series in which the 4f sub-shells are filled. Because (n-2)f electrons do not effectively shield valence electrons, atomic radii continue to decrease across the lanthanide series, a phenomenon known as the lanthanide contraction. As a consequent, transition metals in the 6th period are

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of about the same atomic sizes as the corresponding transition metals elements in the 5th period. Therefore, same group transition metals in the 5th and 6th periods exhibit quite similar chemical properties and they often occur together in nature. For example, zirconium and hafnium, but not titanium, very much resemble each other in their chemical properties that they are often found together in the same ores. Unlike the main group elements, where chemical reactivity increases from top to bottom down a group, the reactivity of transition metals decreases from top to bottom down a group.

Trends Across a PeriodThe atomic size of transition elements shows a decreasing trend towards the center, then gradually

increases across a period. Ionization energy and electronegativity generally increase across a given period. The enthalpy of atomization, melting points, and densities appear to approach maximum values near the middle of the period, which implies strongest metallic bonds in these elements.

Oxidation-Reduction ReactionsMost transition metals have more than one oxidation states, and elements in the middle of the series

usually acquire the highest oxidation states. For example, chromium has the highest oxidation state of +6 in CrO4

2- and Cr2O72-, while manganese has the highest oxidation state of +7 in MnO4

-. Ruthenium and Osmium, both in Group 8, attain the highest oxidation state of +8 in their compounds with oxygen and fluorine, such as RuF8 and OsO4. Transition metals in the highest oxidation states are strong oxidizing agents. They are strong reducing agents when in the lowest oxidation states, except for inactive metals such as palladium, silver, platinum, and gold.

Magnetic PropertiesMost transition metals contain unpaired electrons in their (n-1)d sub-shells and they exhibit

paramagnetic property, which increases as the number of unpaired electrons increases. Iron, cobalt and nickel, are ferromagnetic elements - they can be permanently magnetized. Transition metals located at the end of the series are normally diamagnetic because all electrons are paired. For example, zinc, cadmium and mercury are all diamagnetic. Palladium, which is expected to be paramagnetic, is found to be diamagnetic. This is because both 5s electrons jump to the 4d sub-shell, making the latter sub-shell completely filled with 10 electrons and leaving the 5s sub-shell empty.

Pd: [Kr] __ __ __ __ __ ____ 4d 5s

Exercise-1:1. Write the electron configuration and orbital box diagram of each of the following atom/ion. Indicate

whether each atom/ion is diamagnetic or paramagnetic.

(a) Mn (d) Mn2+

(b) Fe (e) Fe3+

(c) Ni (f) Ni2+

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21.2 The First Row of Transition Metals

Scandium exists in compounds mainly in the +3 oxidation state, such as in ScCl3, Sc2O3, and Sc2(SO4)3. Most of its compounds are colorless and diamagnetic. The metal can be prepared by electrolysis of molten ScCl3, but because of its rarity, the metal is not of commercial importance.

Titanium is widely distributed in the Earth’s crust (0.6% by mass). The metal has a relatively low density, but of high tensile strength, which makes it an excellent structural material, especially in jet engines. Nearly 5000 kg of titanium alloys is used in each engine of a Boeing 747 jetliner. Titanium is also quite resistance to chemical attack.

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The most important compound of titanium is TiO2, which is used to make white pigment in paper, paint, linoleum, plastics, synthetic fibers, and cosmetics. About 700,000 tons of TiO2 is used annually in these products. Titanium(IV) oxide is widespread in nature, but the main ores are rutile (impure TiO 2) and ilmenite (FeTiO3). Rutile is processed by treatment with chlorine to form volatile TiCl4, which separated from the impurities and then burned to form TiO2:

TiO4(s, impure) + 2 Cl2(g) TiCl4(g) + 2 O2(g)

TiCl4(g) + O2(g) TiO2(s) + Cl2(g)

Ilmenite is treated with sulfuric acid to form a soluble sulfate:

FeTiO3(s) + 2 H2SO4(aq) Fe2+(aq) + TiO2+

(aq) + 2 SO42-

(aq) + 2 H2O(l)

Upon standing under vacuum, solid FeSO4.7H2O separates out. The saturated solution is then heated to produce TiO2.H2O, which is further heated to form pure TiO2:

TiO2.H2O(s) + heat TiO2(s) + H2O(g)

The common oxidation state of titanium in most of its compounds is the +4 state. Titanium(III) compounds can be formed by reduction of the +4 state. Aqueous solution of titanium(III) compounds contains the purple complex ion, Ti(H2O)6

3+, which is slowly oxidized to titanium(IV) by air. Titanium(II) is not stable in solution, but occurs in the solid state in compounds such as TiO and the halides, TiX2.

Vanadium is widespread throughout the earth’s crust, but at very low concentration (0.02% by mass). Pure vanadium can be produced by electrolysis of molten VCl2. It is used mainly in alloys with iron and titanium. About 80% of vanadium is used in steel. Like titanium, the metal is corrosion resistant. Ferrovanadium is produced by the reduction of a mixture of V2O5 and Fe2O3 with aluminum, and it is added to iron to form vanadium steel, hard steel used for engine parts and axles. The most important compound of vanadium is vanadium(V) oxide (V2O5, mp = 650oC), which is used as an industrial catalyst in the production of chemicals such as sulfuric acid. V2O5 is orange and contains vanadium in its principal oxidation state of +5. However, VF5 (mp = 19.5oC), in which vanadium has an oxidation state of +5) is colorless.

The oxidation state from +5 to +2 all exist in aqueous solution. Only V2+(aq) and V3+

(aq) exist in the hydrated forms to give a violet and blue-green solution, respectively. The oxidation states +4 and +5 only occur in oxycations, VO2+ (blue) and VO2

+ (yellow), respectively.

Chromium occurs mainly in the ore called chromite (FeCr2O4), which can be reduced by carbon to give ferrochrome

FeCr2O4(s) + 4 C(s) (Fe + 2Cr)(s) + 4 CO(g)

The material ferrochrome (Fe+Cr) can be added directly to iron in the steel making process. Chromium metal, which is often used to make stainless steel, is a hard and brittle metal, which maintains a bright surface due to a tough shiny oxide coating.

The common oxidation states for chromium in its compounds are +2, +3, and +6, for example in CrCl2, Cr2O3, and CrO3, respectively. The Cr2+ ion is a powerful reducing agent in aqueous solution. Traces of O2

can be removed by bubbling through a Cr2+ solution:

4 Cr2+(aq) + O2(g) + 4 H+

(aq) 4 Cr3+(aq) + 2 H2O(l)

The chromium(VI) species are strong oxidizing agents, especially in acidic solution, where chromium(VI) in dichromate, Cr2O7

2-, is reduced to Cr3+ ion:

C2O72-

(aq) + 14 H+(aq) + 6 e- 2 Cr3+

(aq) + 7 H2O(l), Eo = 1.33 V

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A mixture of chromium(VI) oxide and concentrated H2SO4 is a powerful oxidant that can oxidize and remove organic materials from analytical glassware. In basic solution, chromium(VI) occurs as the chromate ion, CrO4

2-, which is not a good oxidizing agent:

CrO42-

(aq) + 4 H2O(l) + 3 e- Cr(OH)3(s) + 5 OH-(aq), Eo = -0.13 V

Compounds of chromium exhibit a variety of colors. For example, Cr2O3 is green, Cr(OH)3 is blue-green, compounds containing chromate ion (CrO4

2-) are bright yellow and those containing dichromate (Cr2O72-)

such as K2Cr2O7 are bright orange, and CrO3 is red. Solutions containing Cr3+(aq) are violet-blue, and

solutions containing the complex ion Cr(OH)4- are bright green.

The oxides of chromium exhibit a range of acid-base properties; for example, CrO is basic, Cr2O3 is amphoteric, and CrO3 is acidic.

CrO(s) + 2 H+(aq) Cr2+

(aq) + H2O(l)

Cr2O3(s) + 6 H+(aq) 2 Cr3+

(aq) + 3 H2O(l)

Cr2O3(s) + 2 OH-(aq) + 3 H2O(l) 2 Cr(OH)4

-(aq)

CrO3(s) + H2O(l) H2CrO4(aq)

Manganese is relatively abundant in the earth’s crust (0.1% by mass). The most common use of the metal is in the production of especially hard steel used for rock crushers, bank vaults, and armor plate. An abundant source of manganese occurs in the form of manganese nodules found on the ocean floor, which contain mixtures of manganese and iron oxides as well as small amounts of cobalt, nickel, and copper.

Manganese can exist in all oxidation states from +2 to +7, with the +2 and +7 states being the most common. In aqueous solution Mn2+ occurs as Mn(H2O)6

2+, which gives the solution a very light pink color. Manganese(VII) is found in the dark purple permanganate ion (MnO4

-), which is a strong oxidizing agent in acidic solution where manganese(VII) is reduced to Mn2+:

MnO4-(aq) + 8 H+

(aq) + 5 e- Mn2+(aq) + 4 H2O(l), Eo = 1.51 V.

Like chromium, compounds of manganese are invariably colored. For example, MnCl2 is light pink, MnSO4

is reddish, MnO2 is brown, KMnO4 is dark purple, and solutions containing MnO42- is green. MnO4

2- is unstable - it quickly disproportionate to MnO4

- and MnO2 in basic solution:

3 MnO42-

(aq) + 2 H2O MnO2(s) + 2 MnO4-(aq) + 4 OH-

(aq)

Iron is the most abundant heavy metal (4.7% of the earth’s crust) and the most important to our civilization. It occurs is ores such as hematite (mostly as Fe2O3), pyrite (FeS2), and calcopyrite (CuFeS2). The metal is quite reactive and in moist air it is rapidly oxidized by oxygen in the atmosphere to form rust, which is a mixture of iron oxides. The most important oxidation states for iron are the +2 and +3 states. Examples of compounds containing the Fe2+ ion are FeO (black), FeS (brownish-black), and FeSO4.7H2O (light green). Solutions containing Fe2+ ion is slowly oxidized to Fe3+ upon standing. Examples of compounds containing Fe3+ ion are Fe2O3 (reddish-brown), FeCl3 (brownish-black), Fe(SCN)3 (red), K3Fe(CN)6 (red). Magnetite (Fe3O4) contains a mixture of Fe2+ and Fe3+.

Cobalt is found in ores such as smaltite (CoAs2) and cobaltite (CoAsS). The metal is mainly used in alloys such as stainless steel and stellite, which is an alloy of iron, copper, and tungsten that is normally used in surgical instruments. The most common oxidation states of cobalt are the +2 and +3 oxidation states found in compounds such as CoO (greenish-brown), CoS (black), CoSO4 (dark blue), CoCl2.6H2O (pink), Co(NO3)2.6H2O (red), CoF3 (brown), Co2O3 (charcoal), K3{Co(CN)6] (yellow), and CoCl3.6NH3 (yellow).

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Aqueous solutions of cobalt((II) salts contain the hydrated Co(H2O)62+ ion, which has a characteristic rose

color.

Nickel is found in ores where it is combined mainly with arsenic, antimony, and sulfur. Nickel is silvery white metal with high electrical and thermal conductivities; it is quite resistant to corrosion and is widely used in the production of alloys such as stainless steel. Common nickel compounds contain almost exclusively Ni2+ ion. Aqueous solutions of nickel(II) salts contain the Ni(H2O)6

2+ species, which has a characteristic emerald green color. Nickel(II) ion forms a complex with dimethylglyoxime (CH3C=NOH)2 to give a characteristic pink color that is often used to test the presence of Ni2+ ion is solution. Nickel(III) species occurs in compound such as NiO(OH), which is an oxidant in the nickel-cadmium (NiCad) battery.

Copper is widely distributed in nature in ores containing sulfides, arsenides, chlorides, and carbonates. It is valued for its high electrical conductivity and its resistance to corrosion and its resistant acid attack (except nitric acid). It is widely used for plumbing, and 50% of all copper produced annually is used for electrical applications. Copper is an important constituent in several well-known alloys, such as brass and bronze. Copper is slowly oxidized by moist atmospheric oxygen to form a characteristic green patina consisting of basic copper sulfate (Cu3(OH)4SO4):

3 Cu(s) + 2 H2O(l) + 2 O2(g) + SO2(g) Cu3(OH)4SO4(s)

The chemistry of copper involves the +1 and +2 oxidation states. Copper(I) compounds are mostly insoluble, although it may form a soluble complex ion, such as CuCl2

- and Cu(CN)2-. Copper(II) compounds

are more common, and aqueous solutions of copper(II) salts have a characteristic bright blue color due to the presence of Cu(H2O)6

2+ species. Some examples of copper compounds are: Cu2O (red), Cu2S (black), CuCl (white), CuO (black), CuSO4.5H2O (bright blue), CuCl2.2H2O (green), and [Cu(H2O)6](NO3)2 (blue).

Zinc is widely spread in the earth’s crust in ores mainly as ZnS (zinc blend). Zinc is a white lustrous metal, which is more reactive than iron and copper. It is an excellent reducing agent. About 90% of the zinc produced is used for galvanizing steel to prevent corrosion. In all of its compounds zinc has an oxidation state of +2 and all compounds of zinc are colorless or white.

21.3 COORDINATION COMPOUNDS

A coordination compound is a compound consisting of at least one complex ion . Complex ion contains a metal ion covalently bonded to ligands (molecules or anions). Ligands are Lewis bases, species that contains at least one pair of nonbonding electrons, which can be shared to form a covalent bond with the central metal ion. This type of covalent bond, in which the shared electrons are donated by one of the atom involved, is originally called coordination covalent bond.

Most compounds containing transition metals have characteristic colors. These colors are due to the complex ions, which in turn depend on the type of transition metal that forms the central ion, its oxidation number, and the type of ligands bonded to the central ion. For examples, Co(H2O)6

2+ is pink; Ni(H2O)62+ is

green; Cu(H2O)62+ is blue, and Cu(NH3)4

2+ is deep blue, but Zn(NH3)42+ and Ag(NH3)2+ are colorless.

Structures of Complex Ions: Coordination Numbers, Geometry, and LigandsA complex ion is characterized by the metal ion and the number and types of ligands covalently bonded

to it. Its structure is described in terms of the coordination number, the geometry, and the number of donor atoms per ligand:

Coordination number is the number of ligand atoms bonded directly to the central ion; it is specific for a given metal ion in a particular oxidation state and compound.For example, the following complex ions have the coordination number 6:

[Co(NH3)6]3+, [Ti(H2O)6]3+, [Fe(CN)6]4-, [Pt(NH3)3Cl3]+

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The following complex ions have the coordination number 4:

[Ni(CN)4]2-, [Cu(NH3)4]2+, [Cu(CN)4]2-, [Zn(NH3)4]2+

These complex ions have the coordination number 2:

[Ag(NH3)2]+, [CuCl2]-, [AuCl2]-

The geometry or shape of a complex ion depends on the coordination number, the nature of the metal ion and its oxidation state, and the type of ligands. For example, complex ions with coordination numbers 2 and 6 are linear and octahedral, respectively. Thus, [Ag(NH 3)2]+, [CuCl2]-, and [AuCl2]- are linear; while [Co(NH3)6]3+, [Ti(H2O)6]3+, [Fe(CN)6]4-, and [Pt(NH3)3Cl3]+ are octahedral. However, complex ions with coordination number 4 are either square planar or tetrahedral. For example, [Ni(CN)4]2- and [Cu(NH3)4]2+ are square planar, but [Cu(CN)4]2- and [Zn(NH3)4]2+ are tetrahedral.

Donor atoms per ligand. Ligands are molecules or anions with one or more donor atoms that each a lone (nonbonding) pair electrons to the metal ion to form a covalent bond. These donor atoms often come from Group 5A(15), 6A(16), or 7A(17).

Ligands are generally classified as monodentate, bidentate or didentate, and polydentate, depending on whether they are bonded to the central metal ion through a single or more donor atoms in the complex.

Monodentate, Bidentate, and Polydentate ligands

Species like H2O, NH3, CO, X-, CN-, NO2-, etc, are called monodentate ligands because each is attached

to the central (metal) ion through a single donor atom. Ethylenediamine (H2NCH2CH2NH2 = en) and oxalate ion (C2O4

2-) are bidentate ligands - they are attached through two donor atoms. Those with more than two donor atoms, such as diethylenetriamine, (H2NCH2CH2NHCH2CH2NH2) and ethylenediamine tetraacetate (EDTA) are generally called polydentate. Polydentate ligands are often called chelate ("crab claw").

-OOCCH2 CH2COO-

NCH2CH2N-OOCCH2 CH2COO-

Compounds, which contain polydentate ligands, are called chelate complexes.

Equilibria involving complexes containing monodentate and polydentate ligands normally favor the latter. For example, the following equilibrium shifts to the right.

[Ni(NH3)6]2+(aq) + 3 en ⇄ [Ni(en)3]2+

(aq) + 6NH3(aq)

where en = H2NCH2CH2NH2

Formulas and Names of Coordination Compounds

There are three important rules for writing the formulas of coordination compounds:

1. The cation is written before anion.2. The charge of cation(s) is balanced by the charge of the anion(s).3. In the complex ion, neutral ligands are written before anionic ligands, and the whole ion is placed in

brackets (with the charge of the complex ion placed outside the bracket).6

The complex ion can be a cation or anion. When writing the formula of a coordination compound, the complex ion (cation or anion) is inside a square bracket, [ ], and its counter ion outside the bracket. For example, in [Cu(NH3)4]SO4, the complex ion is a cation, [Cu(NH3)4]2+, and SO4

2- is the counter anion. In K4[Fe(CN)6], the complex ion [Fe(CN)6]4- is the anion, and K+ is the counter cation. In order to determine the charge of a complex ion, you must be familiar with the charge of its counter ion. For example, in [Cu(NH3)4]SO4, the charge of the complex ion, [Cu(NH3)4]2+, is “+2”. In K4[Fe(CN)6], the charge of the complex ion, [Fe(CN)6]4- is “-4”.

The charge or oxidation state of the central metal ion in the complex can be determined from the net charge of the complex ion and the total charges of the ligands. For example, to determine the oxidation state of Fe in [Fe(CN)6]4-,

Oxidation state of Fe = (charge of complex ion) – (total charge of ligands)

= (-4) – {(6 x (-1)} = +2

The oxidation state of Pt in [Pt(NH3)4Cl2]2+ = (+2) - [(4 x 0) + {2 x (-1)} = +4

Coordination compounds are named systematically using a set of rules:

1. The cation is named before the anion 2. Within the complex ion, the ligands are named, in alphabetical order, before the metal ion;3. Ligands are named in specific ways (refer to naming of ligands)4. A numerical prefix (such as, di, tri, etc.) is used to denote the number of a particular ligand.

(prefix mono- is never used)5. The oxidation state of the central metal ion is indicated by a Roman numeral (in parentheses)6. If the complex ion is an anion, the ending of the metal name is modified to end with "-ate".

Ligands' Names

The following are names used for ligands:

NH3 - ammine; F- - fluoro; C2O42- - oxalato;

H2O - aqua; Cl- - chloro; CN- - cyano;CO - carbonyl; Br- - bromo; OH- - hydroxo; en - ethylenediammine; I- - iodo; NO2

- - nitro;SCN- - thiocyanato; SO4

2- - sulfato; ONO- - nitrito;

The following examples are formulas and systematic names of some coordination compounds in which the complex ions are cation:

[Co(NH3)6]Cl3 - Hexaamminecobalt(III) chloride;

[Pt(NH3)4(NH3)2]Cl2 - tetraamminedichloroplatinum(IV) chloride;

[Ni(NH3)2(en)2]SO4 - diamminebisethylenediaminenickel(II) sulfate;

For coordination compounds in which the complex ions are anions, the ending “-ate” is used:

K2[TiF6] - potassium hexafluorotitanate(IV);

K2[Zn(CN)4] - potassium tetracyanozincate(II);

Na3[CoF6] - sodium hexafluorocobaltate(III)

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Latin names are used for metal ions with Latin names, but apply only to anionic complexes:

Na[AgCl2] - sodium dichloroargentate;

K4[Fe(CN)6] - potassium hexacyanoferrate(II)

Na[Cu(CN)4] - sodium tetracyanocuprate(II)

Na[AuCl4] - sodium tetrachloroaurate(III)

Exercise-2: 1. Name the following compounds:

(a) Na3[FeCO(CN)5] : (d) K[Au(CN)4] :

(b) [Co(en)3]Cl3: (e) Na2[PtCl4]:

(c) [Cu(H2O)2(en)2](NO3)2:

The Werner's Formula for Coordination Compounds

Determination of Coordination Number

(a) Reactivity of Chlorides

A Swiss chemist Alfred Werner began studying coordination compounds in the 1890s. He studied the formula and coordination numbers of the central metal ions by determining the reactivity of chlorides in the compounds and the conductivity of solutions prepared from coordination compounds. For example, coordination compounds of cobalt with the general formula CoCl3

.xNH3, with x varies from 3 to 6 studied by Werner were found to exhibit different chloride reactivity when reacted with excess AgNO 3(aq) as well as different conductivity.

For example, aqueous solution of CoCl3.3NH3 did not produce AgCl precipitate when reacted with

AgNO3(aq). While reactions of CoCl3.4NH3, CoCl3

.5NH3, and CoCl3.6NH3 with excess AgNO3(aq) yielded

1, 2 , and 3 moles of AgCl(s), respectively, per mole of the compounds. For example,

CoCl3.4NH3(aq) + excess AgNO3(aq) AgCl(s) + CoCl2(NO3).4NH3(aq)

CoCl3.5NH3(aq) + excess AgNO3(aq) 2AgCl(s) + CoCl(NO3)2

.5NH3(aq)

CoCl3.6NH3(aq) + excess AgNO3(aq) 3AgCl(s) + Co(NO3)3

.6NH3(aq)

The formation of an AgCl precipitate with AgNO3(aq) is a test for the presence of free Cl- ion in solution. The fact that CoCl3

.3NH3(aq) did not form AgCl(s) implies that all of the three Cl- ions in the compounds are firmly bonded to the Co3+ ion. CoCl3.4NH3 yields a mole of AgCl imply that it must contain only one free Cl- per formula unit; the other two Cl- ions in CoCl3

.4NH3(aq) are firmly bonded to the Co3+ ion. Therefore, CoCl3

.5NH3 and CoCl3.6NH3 yield 2 and 3 free Cl- ions, respectively, per formula unit.

(b) Conductivity of Solution of Coordination Compounds

Werner also determined the conductivity of aqueous solution of CoCl3.xNH3 compounds. It was found

that, CoCl3.3NH3 did not exhibit any conductivity, implying that no free ion is present in solution. While

aqueous solutions of CoCl3.4NH3, CoCl3

.5NH3, and CoCl3.6NH3, respectively, exhibited a conductivity that

indicated the presence of 2, 3, and 4 free ions per formula unit of each compound. These analyses also indicate that in each case the total number of firmly bound Cl- and NH3 equals 6.

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Based on these observations, Werner proposed that Co3+ ion has the coordination number 6, and he proposed a new way of writing formulas for the coordination compounds CoCl3

.nNH3:

No. of Ions No. of free No of CovalentlyOld Formula per Formula Cl-/Formula bonded Cl- Werner's FormulaCoCl3.3NH3 0 0 3 [Co(NH3)3Cl3]

CoCl3.4NH3 2 1 2 [Co(NH3)4Cl2]Cl

CoCl3.5NH3 3 2 1 [Co(NH3)5Cl]Cl2

CoCl3.6NH3 4 3 0 [Co(NH3)6]Cl3

In aqueous solutions, these compounds dissociate as follows:

[Co(NH3)3Cl3] none (does not dissociate);(# of ions = 0)

[Co(NH3)4Cl2]Cl [Co(NH3)4Cl2]+ + Cl-; (# of ions = 2)

[Co(NH3)5Cl]Cl2 [Co(NH3)5Cl]2+ + 2 Cl-; (# of ions = 3)

[Co(NH3)6]Cl3 [Co(NH3)6]3+ + 3 Cl-; (# of ions = 4)

Exercise-3:1. A coordination compound is found to have the empirical formula PtCl4

.4NH3. 100 mL of 0.100 M aqueous solution of this compound is reacted with excess AgNO3(aq), which yields 2.87 g of AgCl. A 0.10 M solution of PtCl4

.4NH3 is found to exhibit conductivity that is equivalent to 0.15 M of NaCl. Deduce the Werner's formula for this coordination compound and give its systematic name. What is the oxidation state and coordination number of Pt in this compound?

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21.4 Isomerism in Coordination Compounds

Isomers are compounds with the same chemical formula but different structure and properties.

Constitutional (or Structural) IsomersThese are compounds with the same molecular formula, but the atoms are connected differently.

Coordination compounds exhibit two types of constitutional isomers - one involves changes in the composition of the complex ion, the other in the donor atom of the ligand.

1. Coordination isomers occur when the composition of the complex ion changes but not that of the compound. This type of isomerism occurs when ligand and counter ion exchange positions, such as in [Pt(NH3)4Cl2](NO2)2 and [Pt(NH3)4(NO2)2]Cl2. This type of isomerism also occurs in compounds where both cation and anion are complex ions, such as in [Cr(NH3)6][Co(CN)6] and [Co(NH3)6][Cr(CN)6], and in [Cu(NH3)4][PtCl4] (violet) and [Pt(NH3)4][CuCl4] (green).

2. Linkage isomers occur when the composition of the complex ion remains the same but the bonding with the ligand donor atom changes. Some ligands can bind to the metal ion through either of two donor atoms:

O N: [:O═C═N:]- [:S═C═N:]- O cyanate thiocyanate

nitrite

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For example, the nitrite ion can bind through either the N atom (nitro-, O2N:) or through one of the O atoms (nitrito-, ONO:) to give linkage isomers. For example, the yellow compound pentaamminenitrocobalt(III) chloride, [Co(NH3)5(NO2)]Cl2, and red pentaamminenitritocobalt(III) chloride, [Co(NH3)5(NO2)]Cl2, are linkage isomers. The cyanate ion can also attach through the O atom (cyanato-, NCO:) or the N atom (isocyanato-, OCN:). While the thiocyanate ion may linked through the S atom (thiocyanato-, NCS:), or the N atom (isothiocyanato-, SCN:)

StereoisomerismStereomers are compounds that have the same atomic connections but different spatial arrangements of the atoms. The two types of stereoisomers are called geometric isomers and optical isomers.

1. Geometric isomers occur when atoms or groups of atoms are arranged differently in space relative to the central ion. For example, in the square planar complex [Pt(NH3)2Cl2] have two geometric isomers:

H3N Cl H3N Cl

Pt Pt

H3N Cl Cl NH3

cis-[Pt(NH3)2Cl2] trans-[Pt(NH3)2Cl2]

The trans-[Pt(NH3)2Cl2] is nonpolar, whereas the isomer cis-[Pt(NH3)2Cl2] is polar. Only the square planar [Pt(NH3)2Cl2] exhibits geometric isomerism with polar cis-isomer and nonpolar trans-isomer. A tetrahedral [Pt(NH3)2Cl2] would not exhibit geometric isomerism, while a square pyramidal shape would also give a polar trans-isomer.

2. Optical isomers (also called enantiomers) are isomers that have the ability to rotate the plane of polarized light passing through their solutions. These isomers occur when a molecule and its mirror image cannot be superimposed. Optical isomers are physically identical in all manners except the direction in which they rotate the plane of polarized light. For example, cis-dichlorobis(ethylenediamine)cobalt(III) ion, [Co(en)2Cl2]+, and its mirror image are not superimposible and they are enantiomers. One isomer is designated d-[Co(en)2Cl2]+ and the other l-[Co(en)2Cl2]+, depending on whether it rotates polarized light to the right (d for dextro-) or to the left (l for levo-). In contrast, trans-dichlorobis(ethylenediamine)cobalt(III) ion does not have optical isomers.

Exercise-4: 1. How many geometric isomers does the compound triamminetrichlorocobalt(III), [Co(NH3)3Cl3] have?

Draw these isomers. Do any of these geometric isomers have enantiomers?

2. Does cis- or trans-square planar diamminedichloroplatinum(II) have optical isomers? Explain.

_____________________________________________________________________________

21.5 Bonding in Complex Ions

The Application of Valence Bond Model to ComplexesIn the formation of complex ion, a filled ligand orbital overlaps with the empty metal-ion orbital. This

orbital overlap concept in bond formation is called the valence bond method. The ligand (Lewis base) donates a pair of electrons and the metal-ion (Lewis acid) accepts the electron pair to form a covalent bond in the complex ion (Lewis adduct). This type of bond, in which one atom contributes both electrons is called coordination covalent bond.

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The valence bond theory for complexes also proposes orbital hybridization of the metal-ion orbitals. The number and type of metal-ion hybrid orbitals that overlap with filled ligand orbitals determine the geometry (shape) of the complex ions.

Bonding in Octahedral Complexes

Consider the complex ion hexaamminechromium(III) ion, [Cr(NH3)6]3+, the which the chromium ion has the electron configuration [Ar] 3d3. The six lowest energy empty orbitals on Cr3+ are two 3d and one 4s, and three 4p:

Cr3+: [Ar] ___ ___ ____ ___ ___ ___ 3d 4s 4p

Hybridization of the empty two 3d and one 4s, and three 4p orbitals yields to six d2sp3 hybrid orbitals. Each of these hybrid orbitals then overlaps with the lone pair orbital on the nitrogen atom of NH3 (as ligand), forming a complex ion with an octahedral geometry:

[Cr(NH3)6]3+ : [Ar] _

3d d2sp3

Bond in Square Planar Complexes

Metal ions with a d8 configuration usually form square planar complexes. For example, in [Ni(CN) 4]2-

ion, the model proposes that empty one 3d, one 4s, and two 4p orbitals of Ni2+ are hybridized to form four dsp2 hybrid orbitals, which point to the corners of a square. These hybrid orbitals accept an electron pair from each of the four CN- ligands. Note that the free Ni2+ has the configuration,

Ni2+: [Ar] ____ ____ ____ ____ 3d8 4s 4p

For Ni2+ ion to have one empty 3d orbital, the two electrons in the half-filled orbital pair and leave one 3d orbital empty. This is consistent with the fact that the complex ion is diamagnetic.

Pairing Ni2+: [Ar] ____ ____ ____ ____ ____ 3d 4s 4p

Hybridization involving one empty 3d, a 4s, and two of the 4p orbitals, followed by orbital overlap between each of the hybridized orbitals with the lone-pair orbitals of ligands lead to the following configuration for the complex ion:

[Ni(CN)4]2-: [Ar] ____ 3d dsp2 4p

The energy gained by using a 3d orbital for bonding in the hybrid orbital is large enough to compensate for the repulsion energy from pairing the electrons.

Bonding in Tetrahedral ComplexesMetal ions with a filled d sublevel, such as Zn2+ and Cd2+, often form tetrahedral complexes. For

example, valence bond method proposes that in the complex ion [Zn(OH)4]2- four lowest empty orbitals - one 4s and three 4p - mix to form four sp3 hybrid orbitals that have the tetrahedral orientation. When these hybrid orbitals are occupied by a lone pair from each of the four OH- ligands, a tetrahedral complex ion is obtained.

[Zn(OH)4]2-: [Ar]

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3d sp3

21.6 Crystal Field ModelWhile the valence bond theory provides an excellent explanation regarding bonding and the shape of

complexes, it does not explain the various color of coordination compounds and sometimes predicts their magnetic properties incorrectly. In contrast, crystal field theory, while providing little information about metal-ligand bonding, appears to explain the color and magnetism clearly. This theory highlights the effects of ligands approach on the metal d-orbital energy levels.

The d OrbitalsThe d orbitals contain lobes that are directed along specific directions in three dimensional space. There

are five d orbitals, each has four lobes, except for the d z2, which has two lobes along the z-axis and a torus (a

donut-shape orbital) in the xy-plane. The four lobes in each of the other d-orbitals lie along the coordinate axes or in the plane between these coordinate axes. For example, the dx

2-y2 orbital has the four lobes along

the x- and y-axes, the dxy orbital has the four lobes in the xy-plane, the dxz has the lobes in the xz-plane, and the dyz orbital has the lobes in the yz-plane. In the atom or free ion, all five d orbitals are degenerate (have the same energy), but in complexes they are split into two or more groups of different energy levels, depending on the coordination number of the complexes.

Splitting of d-Orbital Energy Levels by LigandsAs ligands approach the metal ion it causes the splitting of d-orbitals into two or more energy groups.

The manner in which d-orbitals are split depends on: (1) the identity and oxidation state of the central atom; (2) the number and identity of ligands, and (3) the geometric arrangement of ligands around the central ion. For example, octahedral splitting is different from that of tetrahedral or square planar splitting.

Crystal Field Splitting in Octahedral ComplexesIn an isolated ion, the five d orbital are degenerate - they have the same energy. In an octahedral

complex, the six ligands are arranged around the central metal ion along the x-, y-, z- axes. Repulsion between electron pairs of ligands and electrons in d orbitals on the metal ion causes an overall increase in the energy of d-orbitals. However, orbitals dx

2-y2 and dz

2, which lie on the axes, are more affected than the orbitals dxy, dyz, and dxz. As a results, in octahedral complex, the d-orbitals of the transition metal are split into two sets of degenerate levels, as shown in the diagram:

___ ___ dx2-y

2 and dz2

Splitting by an octahedral fieldforms two sets of d-orbitals - two orbitals with higher energy and three with lower energy

___ ___ ____ dxy, dyz, and dxz

split d-orbitals in octahedral complexes ___ ___ ___ ___ ____ degenerate d-orbitals in isolated metal ions

The difference in energy between the two sets of d-orbitals is called the crystal field splitting (O, where the subscript “O” stands for octahedral).

When assigning electrons to d-orbitals, the three lower d-orbitals must be filled singly first. If the crystal field splitting is weak and O is small, the two upper d-orbitals are filled next to minimize electron-electron repulsion. In a weak-field octrahedral complex with five d-electrons, all of the d-orbitals will be occupied by unpaired electrons and a high spin state is obtained, as shown bellow:

_

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_

A high spin state resulting from a weak field splitting

If the crystal field splitting is strong and O is large, the fourth and fifth electrons will double-up with any two of the electrons in the lower set of d-orbitals. The electron configuration yields a low spin state as shown below.

____ ____

_

A low spin state resulting from a strong field splitting

The paramagnetic property of a complex depends on the number of unpaired d-electrons on the metal. From the paramagnetic property the number of unpaired electrons in the complex can be deduced and it can be determined whether the complex is a low-spin or a high-spin type. A low spin state is the result of a strong crystal field, which causes a large energy splitting. A high spin states is the result of a weak crystal field, which causes small energy splitting.

The crystal field splitting, O, is of the same order of magnitude as the energy of photons of the visible light, in the range of 160 - 280 kJ/mol. When white light strikes a complex or its solution, only certain region of the visible light is absorbed by the complex, while region that does not interact with the complex is transmitted or reflected and seen by the viewers. The transmitted or reflected light, hence the compound, would appear colored. Thus, interactions between visible light and electrons of d-orbitals in complexes yields various colored transition-metal complexes.

The magnitude of O can be determined by measuring the wavelength of light absorbed by the compounds. For example, the complex ion [Ti(H2O)6]3+ contains one d electron (3d1) in octahedral field.

___ ___

O = 243 kJ/mol; l = 493 nm; n = 6.09 x 1014 Hz ___ ___

The single d electron absorbs visible light with a wavelength, = 493 nm, which corresponds to the blue-green region of the visible spectrum. This wavelength also corresponds to a splitting energy, O = 4.04 x 10-

19 J/e- (or 243 kJ/mol e-). Absorption at blue-green region leaves red and orange color intact; consequently, aqueous solution of [Ti(H2O)6]3+ appears reddish.

The table below gives an approximate range of wavelengths and observed colors.Wavelength Color Color Wavelength,Absorbed (nm) Absorbed Observed (nm)________

400 - 430 Violet Yellow 560 - 590430 - 490 Blue Orange 590 - 610490 - 560 Green Red 610 - 750560 - 590 Yellow Violet 400 – 430590 - 610 Orange Blue 430 - 490610 - 750 Red Green 490 - 560

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The magnitude of the crystal field splitting depends on the identity and oxidation state of metal and the type of ligands bonded to the central ion. Metal ions with greater charges attract ligands more strongly, which leads to a greater splitting of the d-orbital energy levels. For a given pair of metal and ligands, the higher the charge on the metal ion, the greater the d-orbital splitting, and a low spin complex is obtained. For example, most octahedral complexes of Co(II) are of weak field, high spin states, while those of Co(III) are of strong field, low spin states. Effective nuclear charge general increases as one goes down a group of transition elements, which leads to a greater d-orbital energy splitting. Most complexes of the second and third series of transition elements are strong field, low spin states.

Spectrochemical SeriesIn complexes containing identical central ion, the magnitude of the crystal field splitting (O) and

spectral properties of complexes depend on the type of ligands attached to the central ion. For example, the paramagnetic property indicates that the complex ion hexaamminechromium(II), [Cr(NH3)6]2+, has four unpaired electrons, whereas hexacyanochromium(II), [Cr(CN)6]4-, has only two unpaired electrons:

___ ___ ___

_

[Cr(NH3)6]2+ is [Cr(CN)6]4- is a high-spin complex low-spin complex.

Crystal field splitting by cyanide ions is greater than that by ammonia molecules. In general, the magnitude of crystal field splitting increases with ligands arranged along the following spectrochemical series:

I- < Br- < Cl- < F- < OH- < C2O42- < H2O < NH3 < en < NO2

- < CN- < COweak-field ligands in termediate strength strong-field ligands

The wavelength of light absorbed by complexes containing these ligands shifts toward the blue region of the visible spectrum from the weakest-field ligands to the strong-field ligands. The color of complexes shifts toward the red region of the spectrum. For example, [CuBr4]2- is violet, [Cu(H2O)6]2+ is blue, and [Cu(en)3]2+

is green. Ligands that have the same atom acting as the electron-pair donor, such as OH -, C2O42-, and H2O

are close together in the spectrochemical series, and complexes containing these ligands usually have the same colors.

Crystal Field Splitting in Tetrahedral Complexes

In tetrahedral complexes the splitting of d-orbital energy is generally smaller than those octahedral complexes. The three d-orbitals that lie between axes increases more than the two d-orbitals that lie along the axes. That is, the energies of dxy, dxz, and dyz are raised more than energies of d x

2-y2 and dz

2. Tetrahedral complexes are generally of high spin types, due to weak field splitting, regardless of the type of metal ions and/or ligands.

___ ___ ___ dxy, dxz, and dyz Splitting of d-orbital in tetrahedralfield yields two set of d-orbitals - lower energy levels consisting of dx

2-y2 and dz

2 and a high set of energy levels consisting of dxy, dxz, and dyz.. ___ ___ dx

2-y2 and dz

2 split d-orbitals in tetrahedral complexes

___ ___ ___ ___ ___d-orbital energy levelsin isolated ion

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Crystal Field Splitting in Square Planar Complexes

In square planar complexes, four ligands are assumed to approach along the x- and y- axes. This causes the greatest repulsion between the ligand electron-pairs and dx

2-y2 orbitals which energy level is raised to a

maximum. Square planar field also raises the energy of the other d-orbitals in the order indicated in the following energy diagram:

___ dx2- y

2

___ dxy

___ dz2

___ ___ dxz and dyz

Splitting of d orbitals in square-planar complexes

The d-orbital energy splitting in square planar field is large and often results in strong-field, low-spin type complexes.

Exercise-5:1. The complex ion [CoF6]3- is paramagnetic. Draw the d-orbital diagram of [CoF6]3- based on the crystal

field theory.

2. Hexaaquacobalt(II) ion, [Co(H2O)6]2+, forms a high-spin complex. Use the crystal field diagram to show the distribution of d electrons in the complex ion. What would be the distribution of d electrons in octahedral hexaamminecobalt(III) ion, [Co(NH3)6]3+, if it is a strong-field, low-spin type? Are these complex ions paramagnetic or diamagnetic? How many unpaired electrons are there in each ion?

3. Tetraamminenickel(II) ion, [Ni(NH3)4]2+, is paramagnetic, while tetracyanonickelate(II) ion, [Ni(CN)4]2-, is diamagnetic. Do both complex ions have the same or different geometrical shape? Explain. What possible geometry does each complex have. Draw the d-electron distributions of each complex according to the crystal field theory.

Uses of Transition-Metal ComplexesSome transition-metal complexes are responsible for the natural color of many biological systems, such as the red

color of hemoglobin. Some synthetic dyes and pigments contain transition-metal complexes. For example, Prussian blue used in paint and water or oil-based colors contains the compound Fe4[Fe(CN)6]3.nH2O. Some transition-metal complexes are used as catalysts for important industrial processes. For example, Co 2(CO)8 is a catalyst for the reaction between hydrogen, carbon monoxide and an alkene to form aldehyde:

Co2(CO)8

RCH═CH2 + H2(g) + CO(g) RCH2CH2CHO

Certain transition metal complexes are used in the extraction of gold and silver from their ores,

4 Au(s) + 8 NaCN(aq) + O2(g) + 2 H2O(l) 4 Na[Au(CN)2](aq) + 4 NaOH(aq)

Nickel can be purified by converting it first into a complex with carbon monoxide, which takes place at low temperature:

80oCNi(s) + 4 CO(g) > Ni(CO)4(g)

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The Ni(CO)4 complex which boils at 43oC is separated from impurities by distillation and then decomposed to pure nickel at high temperatures:

230 oCNi(CO)4(g) Ni(s) + 4CO(g)

This process yields 99.95 % pure nickel; the carbon monoxide can be re-cycled.

21.7 The Biological Importance of Coordination Complexes

Certain transition metals have important biological functions and their activity depends on the ability to form complexes with biological molecules. For example, iron are important components of hemoglobin. They are responsible in oxygen transport. Cobalt is an important component of vitamin B12, which takes part in the development of red blood cells. Other biologically important transition metals are listed in the following table:

Transition Metals in Biomolecules:Element Biomolecule Containing Metal Function of Biomolecules______

Chromium Glucose tolerance factor Glucose utilization

Manganese isocitrate dehydrogenase Cell respiration

Iron Hemoglobin and myoglobin Oxygen transportCytrochrome c Cell respiration; ATP formationCatalase Decomposition of H2O2

Cobalt Cobalamin (vitamin B12) development of read blood cells

Copper Ceruloplasmin Hemoglobin synthesisCytochrome oxidase Cell respiration; ATP formation

Zinc carbonic anhydrase Elimination of CO2 Carboxypeptidase Protein digestion

Alcohol dehydrogenase Metabolism of ethanol_________

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