Co ordination compound

Coordination compound

The chemistry of Co-ordination compounds and chelating agents

Definitions of Some Terms

Addition or molecular compounds

When solutions of two or more salts (in simple molecular ratio) are mixed and allowed to evaporate, 
crystals of new compounds are obtained. This new compounds are called addition or molecular 
compound.

Molecular compounds are of two types as discussed below:

Double salts:

The addition compounds which are    (a) stable in solid state, but are broken into their individual 
compounds when dissolved in water or any other ionic solvent,  (b) have the same properties as 
those of individual components are called double salts.

Example:

1. Mohr’s salt – FeSO4.(NH4)2SO4.6H2O
2. Carnallite – KCl.MgCl2.6H2O
3. Potash alum – K2SO4.Al2(SO4)3.24H2O

Coordination/Complex compounds:

The addition compounds which are (a) stable in the solid as well as when dissolved in water or any 
other ionic solvent,  (b) have properties which are completely different from their components are 
called coordination/complex compounds.

Or in other words: Complex compounds are chemical substances formed by the combination of 
compounds which are capable of independent existence. Complex compounds are also known as 
coordination compounds.

Example:

When aqueous NH3 is added to a green solution of NiCl2 and the solution thus obtained is evaporated,
 purple crystals of a new compound having the composition, [Ni(NH3)6]Cl2 are obtained.

This new compound when dissolved in water, ionizes;  [Ni(NH3)6]Cl2 ⇌[Ni(NH3)6]2+ +2Cl–

This is thus a complex compound.                        Complex ion

Difference between double salts and complex compounds

Point	Double salts	Coordination/complex compounds
1. Stability	Stable in the solid state but dissolves into its component compounds in water or other ionic solvents	These compounds retain their identities in the solid as well as when dissolved in water or any other ionic solvents
2. Physical and chemical properties	Physical and chemical properties are same as those of its individual components	Their physical and chemical properties are completely different from those of their individual components.
3. Example	1. Mohr’s salt – FeSO4.(NH3)6.6H2O 2. Carnallite – KCl.MgCl2.6H2O 3. Potash alum – 2SO4.Al2(SO4)3.24H2O	1. [Ni(NH3)6]Cl2 2. [Fe(CN)6]4- etc.

Complex ion

A complex ion can be defined as an electrically charged radical which is obtained by the combination 
of a metal cation with one or more neutral molecules or simple anions.
Simply, a complex ion is an ion where central metal cation/atom is attached with one or more neutral 
molecules or simple anions.

Example: [Ni(NH3)6]2+, [Fe(CN)6]4─ etc.

Ligands or Coordinating groups

The neutral molecules or simple ions (usually anions) which are attached with the central (metal) 
ion (cation) or atom in the complex compounds/ions are called ligands or coordinating grounds.

In the above complex ion, [Fe(CN)6]3- the six CN ─ ions which are attached with the central Fe3+ ion act 
as ligands. In Lewis sense, in most of the complex compounds the ligands act as Lewis bases (electron
 pair donors) and central metal ion acts as a Lewis acid (electron pair acceptor).

The ligands are arranged round the central metal ion inside the first sphere of attraction in preferred geometries 
which are: linear, equilateral triangulai; tetrahedral, square planar, trigonal hipyramidal, square pyramidal and 
octahedral.

Classification of ligands:

The ligands are classified as follows:

Monodentate/Unidentate ligands Bidentate ligands Tridentate ligands

Tetradentate ligands Pentadentate ligands Hexadentate ligands

1. Monodentate/Unidentate ligands:

The ligands which are coordinated through one electron pair are called monodentate or unidentate 
ligands.

Example:

1. Neutral: Thiourea (abbreviated as tu), Pyridine (abbreviated as py),
   Ammine - *NH3, H2O*, *CO, *NO                         
2. Anionic: F-, Cl-, Br-, I-, CN -, NCS-, NO2-, NH2-, *OH-, CH3COO*- etc.

2. Bidentate ligands:

The ligands which are coordinated through two electron pair are called bidentate ligands.

Example:

1. Neutral: o-phenanthroline.

2. Anionic: Oxaleto (C2O42-) abbreviated as ox2-, Glycinato (gly─)

                                            

3. Tridentate ligands:

The ligands having three donor atom (coordinated through three electron pairs) are called tridentate ligands.

Example:

1. Neutral: Triamino propane. CH 2 (NH2 )—CH(NH2 )—CH2 (NH2 )

                        ⃰         ⃰           ⃰   

2. Anionic: Anion of aspartic acid (Asp2-), Anion of diamino prpinoic acid etc.

              –OOC—CH2—CH(NH2 )—COO–        H2C(NH2 )—CH(NH2 )—COO–

           ⃰                     ⃰             ⃰                ⃰           ⃰            ⃰

                                          

4. Tetradentate ligands:

The ligands having four donor atoms (coordinated through four electron pairs) are called tetradentate ligands.

Example:

1. Neutral: Triethylentetramine (trien).

2. Anionic: Nitrilotriacetate (nta3-).

5. Pentadentate ligands:

The ligands having five donor atoms (coordinated through five electron pairs) are called pentadentate ligands.

Example:

1. Neutral: Tetraethylenepentamine (tetraen).

2. Anionic: Ethylenediaminetriacetato.

6. Hexadentate ligands:

The ligands having six donor atoms (coordinated through six electron pairs) are called hexadentate ligands.

Example: Ethylenediamine tetracetato ligand

Anion of EDTA (ethylene diamine tetraacetato)

   

Coordination number

The number of ligands which are directly attached with the central metal atom/ion is known as the 
 coordination number of that metal atom/ion.

Example: In [FeCl4]─ , Fe3+ ion is directly attached with four Cl─ ligands, therefore in this compound 
the coordination number of Fe3+ is 4.The co-ordination number of Ag+ ion in [Ag(NH3)]2+ is two.

Chelates and chelation

When all the donor atoms of a polydentate ligand get coordinated with the same metal ion a complex
 ion with one or rings in its structure. This is called a chelated or cyclic complex or a chelate.

The process of the formation of a chelated complex is called chelation.

Example: the figure below shows a chelate complex given by ethylene diamine tetraacetato Ligand 
with metal atom, M.

        

Polydentate ligands are called chelating ligands where monodentate ligands do not form any chelate 
complex.

Application of Chelation

Application of chelates can be studied under the following heads:

1. Formation of chelates in analytical chemistry:

1. Estimation and identification of Ni2+ ions by dimethyl glyoxime: Dimethyl glyoxime solution 
reacts with Ni2+ ions in ammoniacal medium and forms a red colored precipitate of bis- (dimethyl 
glyoximato) nickel (II). Thus formation of this precipitate can be used for the identification, estimation
 of Ni2+ and also for the separation of Ni2+ ions from Co2+ ions.

2. Estimation of Mg2+ and Ca2+ ions by EDTA: Firstly the metal ions (Mg2+ or Ca2+) are buffered to 
pH = 10, a few drops of the Erio-chrome Black T (abbreviation. H3D) are added and the solution is 
titrated with a standard solution of EDTAH2Na2. In the titration, indicator ions (D3- ions) form red
 metal-indicator complex with M2+ (M2+ = Mg2+ or Ca2+) ions.

Metal indicator complex reacts with EDTA4- ion and forms metal-EDTA complex.

Thus we see that free indicator is obtained at the end point. The release of the end point is marked by 
a change of red color complex to blue color.

Criteria of indicator:

1. The indicator must form a colored complex with the metal ion at the same pH at which 
   the EDTA4- ion forms complex with the same metal ion.
2. The color of the metal-indicator complex should be different from that of the free 
   indicator ion.
3. The stability of the metal-indicator complex should be low compared to that of 
   metal-EDTA complex.

2. Formation of chelates in medicine:

1. In the removal of poisonous metals from the body: An injection of CaNa2EDTA is given to the 
patient for this purpose. It reacts with poisonous metal (Pb) and forms metal-EDTA complex which is
 excreted in urine.

2. Disorder of liver: the disease regarding the disorder of liver leads to the malfunctioning of the 
CNS. This disease is due to the fact that free Cu2+ is accumulated in the liver. This excess Cu2+ is
 removed by giving penicillamine (I) which forms a soluble complex with Cu2+.

    C. Formation of chelates in living systems:

1. Chlorophyll is a square planar complex of Mg2+. Chlorophyll is responsible for photosynthesis in 
plants.

2. Hemoglobin, a conjugated protein is a porphyrin complex of Fe2+ ion. Hemoglobin imparts red 
color to the human blood. In human body at lungs blood is saturated with air and the hemoglobin is 
completely changed to the oxyhemoglobin form.

As the blood runs through the arteries to the tissues, the oxygen pressure dwindles and the bound 
oxygen gets liberated. Such way hemoglobin serves as the carrier of oxygen in the body.

Heme is a non-protein chelated complex compound.

Hemoglobin

It has also been found that oxygen carriers in the blood of various invertebrates are the chelates of Cu, 
Mn and V.

3. At present it is believed that all the essential trace metals in human body functions through process
 involving chelate formation.

Theories of coordination bond

1. Werner’s coordination theory:

Werner’s coordination theory was introduced in 1893 by Werner to explain the formation complex 
compounds. It explains the formation of compounds like amines of Pt (IV) 
(e.g. PtCl4.6NH3, PtCl4.5NH3, PtCl4.3NH3, PtCl4.2NH3) and Co (III) 
(e.g. CoCl3.6NH3, CoCl3.5NH3, CoCl3.4NH3, CoCl3.3NH3).

Features of Werner’s coordination theory:

a) Every metal atom or ion exhibits two types of valencies. They are the “primary” and “secondary” 
valency.

Primary valency

• It is satisfied by negative ions
• It corresponds to the oxidation state of the metal atom/ion.
• It is ionizable.

Secondary valency

• It is satisfied either by negative ions or neutral molecules.
• It corresponds to the coordination number of the metal atom/ion.
• It is non-ionizable.

b) Every metal atom or ion wants to satisfy both its primary and secondary valencies. In order to 
meet this requirement a negative ion may perform dual function.

c) Every metal atom /ion has a fixed number of secondary valencies.

d) The ligands which satisfy the secondary valencies are always directed towards fixed positions in 
space which determines the geometry of the complex.

For instance if

Secondary valency = 4, complex is tetrahedral/square/planar

Secondary valency = 6, complex is octahedral

2. Sidgewick’s electronic concept of coordination bond in complex compound:

In 1916 lewis introduced a concept of two-electron covalent bond between two atoms in a molecule. 
Sidgewick extended this concept in 1927 and introduced a new concept of coordinate bond.

Features of Sidgewick’s electronic concept:

a). The ligands donate the electron pair to the central metal ion and thus form a coordinate bond.

This indicates that the ligand, L has donated an electron pair to the metal ion, M.

b). All ligands contain at least one lone pair of electrons.

EAN (Effective Atomic Number):

The total number of electrons on the central atom including those gained from the ligands in the 
bonding is called the Effective Atomic Number (EAN) of the central metal atom/ion.

EAN of a central metal atom or ion in a given complex is calculated by the following formula.

Where,

    Z = Atomic number of the central metal atom

    x = positive oxidation state of the central metal atom

    n = number of ligands

    y = number of electrons donated by one ligand

For example in [Fe2+(CN)6]4- the EAN of Fe = 26 ─2 + 6X2 = 36

Limitation of sidgewick’s concept:

If each ligand donates one pair of electrons to the central metal atom/ion to form L→M coordinate 
bond, a negative charge on the central atom is accumulated which is most unlikely. For example, in 
[Co3+(NH3)6]3- six NH3 molecules donate half share of 12 electrons (i.e. 6 electrons) to Co3+ ion and 
thus +3 charge on cobalt atom should be reduced to –3 (+3 – 6 = –3). Such accumulation of negative 
charge on the central atom is unlikely.

3. Valence bond theory (VBT):

The theory is mainly due to Pauling. It deals with the electronic structure of the central metal ion in
 its ground state, kind of bonding, geometry and magnetic properties of the complexes.

Features of Valence bond theory:

i) The central atom/ion makes available a number of empty s, p and d atomic orbitals equal to its 
coordination number. These vacant orbitals hybridize together to form hybrid orbitals. These hybrid 
orbitals are vacant, equivalent in energy and have defined geometry.

ii) The ligands have at least one σ-orbital containing a lone pair of electrons.

iii)  Hybridized vacant orbitals of the metal atom/ion overlap with the filled σ-orbitals of the ligands
 to form ligand → metal σ-bond (M→L bond).

iv) The non-bonding electrons of the metal atom or ion are then rearranged in the metal orbitals 
(pure d, s or p) which do not participate in the forming of the hybrid orbitals. This rearrangement 
takes place according to Hund’s rule.

Limitations of VBT:

1. VBT cannot explain the relative stabilities for different shapes and different coordination numbers 
in metal complexes.

2. VBT cannot explain the relative rates of reactions of analogous metal complexes.

3. This theory can’t explain as to why Cu(+2) forms only distorted octahedral complexes even when 
all six ligands are identical.

4. It classifies metal complexes on the basis of their magnetic behaviour into covalent and ionic
 complexes. It is not satisfactory and often misleading.

5. VBT fails to explain the finer details of magnetic properties of complexes including the magnitude 
of the orbital contribution to the magnetic moments.

6. VBT cannot interpret the spectra (color) of the complexes.

7. This theory does not predict or explain the magnetic behaviour of the complexes, this theory only 
predicts only predicts the number of unpaired electrons. Its prediction even for the number of unpaired electrons and their correlation with stereochemistry may be misleading sometimes.

8. VBT can’t explain the order of reactivities of the inner-orbital inert complexes of d3, d4, d5 and d6 
 ions and also can’t explain the observed differences in the energies of activation in a series of similar 
complexes.

9. The magnetic moment values of the complexes of certain ions (e.g. Co2+, Ni2+ etc.) are much higher
 than those expected by spin only formula. VBT cannot explain the enhanced values of the magnetic 
moment.

Isomerism

Coordination compounds show two types of isomerism namely structural isomerism and stereo 
isomerism or space isomerism both of which can further be subdivided.

1. structural isomrerism: This isomerism is of the following types:

a. Conformation isomerism: In this type of isomerism two isomers have different geometries but 
otherwise identical.

Example: [Ni2+(P.Et.Ph2)2Br2] has two conformation isomers where one is green, paramagnetic and 
tetrahedral while the other is brown, diamagnetic and square planar.

b. Ionisation isomerism: Complexes which have the same empirical formula but they are produced 
by the interchange of the position of ligands inside the complex zone and the anions outside the 
complex zone are called ionisation isomers and this type of isomerism is regarded as ionisation 
isomers.

Example: [Co3+(NH3)5Br]SO4 and [Co3+(NH3)5(SO4)]Br are two ionisation isomers and they produce 
different ions in solution on ionisation.

Fig: Types of isomerism in complex compound

c. Hydrate isomerism: This type of isomerism is due to different disposition of water molecules 
inside and outside the coordination sphere (i.e. complex zone).

Example: [Cr(H2O)6]Cl3 and [Cr(H2O)5Cl]Cl.2H2O are two hydrate isomers. The first one is violet, 
doesn’t lose water over H2SO4 and all Cl─ ions are precipitated by Ag+ ions, but the second one is 
green, loses two water molecule over H2SO4 and only one Cl─ ion is precipitated by Ag+ ions.

d. Ligand isomerism: this type of isomerism is due to the isomerism of the ligands themselves.

Example: [Co(pn)2Cl2]+ and [Co(tn)2Cl2]+ ions. Here pn = 1, 2-diamino propane and tn = 1, 3 diamino 
propane.

e. Linkage isomerism: Ambidentate ligands (possesses two different donor atoms like N & O) can
 coordinate to the metal ion/atom through any of the two atoms. This gives two linkage isomers and 
this is called linkage isomerism.

Example: NO2─ is an ambidentate ligand which can attach to the metal through N or O atoms. In the 
first case it is called nitro-N or nitro isomer and in second case it is called nitro-O or nitrite isomers.

f. Coordination position isomerism: In some poly-nuclear complexes (more than one metal atom/ion 
present) an interchange between ligands is possible. As a result isomers are created and this type of
 isomerism is called coordination position isomerism.

Example: The following two compounds are coordination position isomers.

g. Coordination isomerism: if both the anion and cation of a complex compound are complex then it 
is possible for a exchange of ligands between the ions. Isomerism obtained in such way is called 
coordination isomerism.

Example: [Cr(NH3)6]3+[Cr(SCN)6]3- and [Cr(NH3)4(SCN)2]+[Cr(NH3)2(SCN)4]─ are coordinate isomers to each other.

h. Polymerisation isomerism: This type of isomerism is found in complex compounds which are 
polymers of simple complex compounds.

Example:

    (i) [Co(NH3)3(NO2)3]

    (ii) [Co(NH3)6][Co(NO2)6]

    (iii) [Co(NH3)4(NO2)2][Co(NH3)2(NO2)4]

    (iv) [Co(NH3)5(NO2)]3[Co(NO2)6]2

(ii) and (iii) complexes are dimers of (i) complex and (iv) complex is pentamer of (i) complex.

2. Stereo/space isomerism: When two compounds contain the same ligands coordinated to the 
same central metal atom, but the arrangement in space is different, the two compounds are said to 
be stereo-isomers of each other and this type of isomerism is called stereo/space isomerism.
This is following types.

a. Geometrical isomerism:

In this isomerism, the ligands occupy different positions round the central metal ion/atom.
If the two identical ligands occupy adjacent positions the isomer is called a cis-isomer.
On the other hand when two identical ligands are placed opposite to each other, the isomer is called 
a trans-isomer.

Example: Cis- and trans-isomers of Pt2+(NH3)2Cl2

b. Optical isomerism:

When the solutions of certain complex compounds are placed in the path of a plane-polarised light 
they rotate its plane through a certain angle which may be either to the left or to the right.
This is called optical activity and complexes showing optical activity can exist in two forms, d-form 
and l-form. They are called optical isomer of each other this is called optical isomerism.

d-form/dextro-rotatory: Rotates the plane to the right (i.e. in clockwise direction).

l-form/levo-rotatory: Rotates the plane to the left (i.e. in anticlockwise direction).

Conditions for a molecule to show optical isomerism:

1. The compound must be an asymmetric molecule which has no plane of symmetry.
   Only asymmetric molecule is optically active.

2. It should not be superimposable on its mirror image.

Characteristics of d-form and l-form:

1. These two forms have exactly identical physical and chemical properties. Only difference is 
   under polarized light.
2. d-form , l-forms are mirror images to each other, thus they can be superimposed on each other. 
   Due to this property they are called enantiomers.
3. The substance composed of 50% d-form and 50% l-form is called racemer and it has n optical 
   activity.

Naming coordination compounds

Examples:

        [NiCl4]2−     → tetrachloronickelate(II) ion

        [CuNH3Cl5]3−    → amminepentachlorocuprate(II) ion

        [Cd(en)2(CN)2]→dicyanobis(ethylenediamine)cadmium(II)

        [Co(NH3)5Cl]SO4 → pentaamminechlorocobalt(III) sulfate

The coordination compounds are named in the following way.

A. To name a coordination compound, no matter whether the complex ion is the cation or the anion, 
the cation is always named before anion. (This is just like naming an ionic compound.)

B. In naming the complex ion:

1. The ligands are named first, in alphabetical order, then the metal atom or ion.
Note: The metal atom or ion is written before the ligands in the chemical formula.

2. The names of some common ligands are listed in Table 1.

• For anionic ligands end in "-o"; for anions that end in "-ide"(e.g. chloride), "-ate" (e.g. sulfate,
   nitrate), and "-ite" (e.g. nirite), change the endings as follows:

-ide →-o; -ate → -ato; -ite → -ito.

• For neutral ligands, the common name of the molecule is used e.g. H2NCH2CH2NH2 
  (ethylenediamine).

Important exceptions: water is called ‘aqua’, ammonia is called ‘ammine’, carbon monoxide is called 
‘carbonyl’, and the N2 and O2 are called ‘dinitrogen’ and ‘dioxygen’.

Table 1. Names of Some Common Ligands

Anionic Ligands	Names		Neutral Ligands	Names
Br–	bromo		NH3	ammine
F–	fluoro		H2O	aqua
O2–	oxo		NO	Nitrosyl
OH–	Hydroxo		CO	Carbonyl
CN–	cyano		O2	dioxygen
C2O42–	oxalato		N2	dinitrogen
CO32–	carbonato		C5H5N	pyridine
CH3COO–	acetato		H2NCH2CH2NH2	ethylenediamine

3. Greek prefixes are used to designate the number of each type of ligand in the complex ion,
e.g. di-, tri- and tetra-. If the ligand already contains a Greek prefix (e.g. ethylenediamine) or if it is 
polydentate ligands (ie. can attach at more than one binding site) the prefixes
bis-, tris-, tetrakis-, pentakis-, are used instead. (See examples 3 and 4.)

The numerical prefixes are listed in Table 2.

Table 2. Numerical Prefixes

 

Number	Prefix	Number	Prefix	Number	Prefix
1	mono	5	penta (pentakis)	9	nona (ennea)
2	di (bis)	6	hexa (hexakis)	10	deca
3	tri (tris)	7	hepta	11	undeca
4	tetra (tetrakis)	8	octa	12	dodeca

4. After naming the ligands, the central metal is named. If the complex ion is a cation, the metal is 
named same as the element.

For example, Co in a complex cation is call cobalt and Pt is called platinum. (See examples 1-4). If the
 complex ion is an anion, the name of the metal ends with the suffix –ate. (See examples 5 and 6.).

For example, Co in a complex anion is called cobaltate and Pt is called platinate. For some metals, the
 Latin names are used in the complex anions e.g. Fe is called ferrate (not ironate).

Table 3: Name of Metals in Anionic Complexes

Name of Metal	Name in an Anionic Complex
Iron	Ferrate
Copper	Cuprate
Lead	Plumbate
Silver	Argenate
Gold	Aurate
Tin	Stannate

5. Following the name of the metal, the oxidation state of the metal in the complex is given as a 
Roman numeral in parentheses.

C. To name a neutral complex molecule, the rules of naming a complex cation is followed. 
Remember: The (possibly complex) cation is named BEFORE the (possibly complex) anion.
See examples 7 and 8.

For historic reasons, some coordination compounds are called by their common names. For example,
 Fe(CN)63- and Fe(CN)64- are named ferricyanide and ferrocyanide respectively, and Fe(CO)5 is called 
iron carbonyl.

Examples Give the systematic names for the following coordination compounds:

1. [Cr(NH3)3(H2O)3]Cl3

Answer: triamminetriaquachromium(III) chloride

Solution: The complex ion is inside the parentheses, which is a cation.

The ammine ligands are named before the aqua ligands according to alphabetical order.

Since there are three chlorides binding with the complex ion, the charge on the complex ion must be +3 
(since the compound is electrically neutral).

From the charge on the complex ion and the charge on the ligands, we can calculate the oxidation number of 
the metal. In this example, all the ligands are neutral molecules. Therefore, the oxidation number of chromium 
must be same as the charge of the complex ion, +3.

2. [Pt(NH3)5Cl]Br3

Answer: pentaamminechloroplatinum(IV) bromide

Solution: The complex ion is a cation; the counter anion is the 3 bromides.

The charge of the complex ion must be +3 since it bonds with 3 bromides.

The NH3 are neutral molecules while the chloride carries -1 charge. Therefore, the oxidation number of 
platinum must be +4.

3. [Pt(H2NCH2CH2NH2)2Cl2]Cl2

Answer: dichlorobis(ethylenediamine)platinum(IV) chloride

Solution: ethylenediamine is a bidentate ligand, the bis- prefix is used instead of di-

4. [Co(H2NCH2CH2NH2)3]2(SO4)3

Answer: tris(ethylenediamine)cobalt(III) sulfate

Solution: The sulfate is the counter anion in this molecule. Since it takes 3 sulfates to bond with two complex 
cations, the charge on each complex cation must be +3.

Since ethylenediamine is a neutral molecule, the oxidation number of cobalt in the complex ion
must be +3.

Again, remember that you never have to indicate the number of cations and anions in the name of an ionic 
compound.

5. K4[Fe(CN)6]

Answer: potassium hexacyanoferrate(II)

Solution: potassium is the cation and the complex ion is the anion.

Since there are 4 K+ binding with a complex ion, the charge on the complex ion must be - 4.

Since each ligand carries –1 charge, the oxidation number of Fe must be +2.

The common name of this compound is potassium ferrocyanide.

6. Na2[NiCl4]

Answer: sodium tetrachloronickelate(II)

Solution: The complex ion is the anion so we have to add the suffix –ate in the name of the metal.

7. Pt(NH3)2Cl4

Answer: diamminetetrachloroplatinum(IV)

Solution: This is a neutral molecule because the charge on Pt+4 equals the negative charges on the four chloro 
ligands.

If the compound is [Pt(NH3)2Cl2]Cl2, eventhough the number of ions and atoms in the molecule are identical to 
the example, it should be named: diamminedichloroplatinum(II) chloride, a big difference.

8. Fe(CO)5

Answer: pentacarbonyliron(0)

Solution: Since it is a neutral complex, it is named in the same way as a complex cation. The common name of 
this compound, iron carbonyl, is used more often.

9. (NH4)2[Ni(C2O4)2(H2O)2]

Answer: ammonium diaquabis(oxalato)nickelate(II)

Solution: The oxalate ion is a bidentate ligand.

10. [Ag(NH3)2][Ag(CN)2]

Answer: diamminesilver(I) dicyanoargentate(I)

Solution: There can be a compound where both the cation and the anion are complex ions.
Notice how the name of the metal differs even though they are the same metal ions.

H E G