7
COMPLEX FORMATION
7.1 WHAT IS A METAL COMPLEX?
What is your picture of copper (II) nitrate dissolved in water? Most likely, it would be that the
ions are floating around individually at random. This is not the case for many salts: your
perception of ionic species in solution has been simplified in the past.
In fact, each of the copper ions has four water molecules linked to it in a definite way: if
you imagine a copper atom at the centre of a square, the water molecules are at each corner.
If ammonia is added to the solution, the solution goes a deep blue, as the water molecules are
replaced by ammonia molecules. The result is an ion that is more than just the metal ion,
since it also contains four neutral ammonia molecules. The product is a species which has
new properties, which are different to those of the individual components. For example, if the
complex ion was crystallised from solution (with sulfate anion), it would not smell of
ammonia. The structures and reaction are shown in Figure 7.1.
H2O
OH2
Cu2+
H2O
H3N
Cu2+
+ 4NH3
OH2
NH3
H3N
+ 4H2O
NH3
FIGURE 7.1 Copper complexes
Such an ion is known a metal complex: a species (can be neutral, cationic or anionic)
containing a metal ion bonded to a set number of neutral or anionic species arranged in a
particular geometry.
The name given in general to the species that links to the metal ion is ligand. Thus, water
and ammonia act as ligands in the example above. Ligands are characterised by having at
least one pair of non-bonded electrons. They may be neutral molecules, such as water
and ammonia, or anions, such as chloride, fluoride and cyanide.
Not all metals ions form complexes: the greater the charge and the larger the ion, the
more likely that a metal will form a complex. One of the significant properties of the
transition metals is their ability to form complexes. On the other hand, alkali metal ions do
not form complexes at all.
The number of metal-ligand bonds formed in a complex is known as the
coordination number (abbreviated as CN). In Figure 7.1, the coordination number for the
two copper complexes is four in each case. The coordination number is determined by the
size and charge of each partner in the complex. Iron (III) forma a complex with six fluoride
ions, but only four chloride ions. This is due to the larger size of the chloride ion. Some metal
ions have constant coordination numbers, e.g. chromium (III) six and platinum (II) four.
Some ligands have more than one bonding site. They are known as chelating agents.
Ligands may classified by the number of bonding sites, using the Latin word for teeth:
dentate. Thus, ammonia is monodentate, while the organic nitrogen compound commonly
known as ethylenediamine is bidentate. A very important ligand in analytical chemistry is
known as EDTA (EthyleneDiamineTetraAcetic acid), which is hexadentate: it has six
bonding sites. The structures of these are given below in Figure 7.2.
7. Complex Formation
CH2CH2
H2N
CH2COOH
HOOCCH2
NCH2CH2N
NH2
HOOCCH2
(a)
FIGURE 7.2
CH2COOH
(b)
Important multidentate ligands (a) ethylenediamine (b) EDTA; bonding sites shown by arrows
Ligands, such as ethylenediamine and EDTA, when complexed to a
metal, form a type of ring structure, since the ligand loops around
to bind to the metal in at least two places. This type of complex is
shown in Figure 7.3, and is known as a chelate.
FIGURE 7.3
Chelate complex of calcium ion with EDTA (solid lines indicate carbon chains
linking bonding atoms, dotted lines are the metal-ligand bonds)
Charges of complexes
The charges held by the individual components of the complex - the cation and the ligand(s) determine the overall charge on the complex. It is a simple matter if adding up these charges.
EXAMPLE
What are the charges on the following complexes: (a) copper (II)-ammonia, (b) copper (II)EDTA, (c) copper (II)-fluoride (CN 4)?
(a)
(b)
(c)
Ammonia is a neutral ligand, and therefore, doesn’t affect the 2+ charge of copper. The
charge of the complex is 2+.
When EDTA forms a complex, it loses its four acid protons, thus attaining a 4- charge.
When combined with the 2+ charge of copper, the complex gains a 2- charge.
Four fluoride ions have a total charge of 4-, so the complex has a charge of 2-.
The structural formula for the complex ion is often shown in large square brackets, with the
overall charge outside the brackets, rather than next to the cation, as shown in Figure 7.4.
3F
F
Fe
F
F
F
F
FIGURE 7.4 Structural formula for FeF63-
Chemistry 2
7.2
7. Complex Formation
PRACTICE QUESTIONS
1.
Explain why ethane could not act as a ligand.
2.
Draw the structure of a complex of copper (II) with (a) cyanide ion and (b)
ethylenediamine as ligands. The coordination number in each case is 4. What is the
charge on the complex ion in each case?
3.
What is the coordination number of the complex in Figure 7.3?
7.2 HOW DO THE BONDS FORM?
The exact way that a ligand bonds to a metal is quite complex (pun intended), and is not
entirely understood by chemists. One theory suggests it is similar to a very polar covalent
bond, but where one atom (in the ligand) contributes both electrons to the bond. The other
theory, which better explains experimental results, proposes a bond more like ionic
attraction.
It is not critical for you to understand either of these theories: a picture of the metalligand bond as being somewhere between the two is all that is necessary. It is most important
that you realise that the bond is not as strong as a covalent bond, but does not break apart in
water (like an ionic bond).
7.3 FACTORS AFFECTING COMPLEX FORMATION
The reaction of a metal ion (M) with a ligand (L) is an equilibrium reaction, as shown in
Equation 7.1, where n is the number of ligands involved in the complex.
M + nL  MLn
Eqn 7.1
An equilibrium constant can be written for the reaction, as shown in Equation 7.2.
[ML]
K = 
[M][L]n
Eqn 7.2
This equilibrium constant is given a special name to indicate that it refers to the formation of
a complex: hence the name, formation constant (Kf). It has also been called the stability
constant (Ks).
Being formed in an equilibrium process, the complex can be destroyed - the
equilibrium shifted towards the separate metal ion and free ligand. The value of the
formation constant is a guide to the strength (or stability) of the complex: the larger the
number, the less likely that the complex will break up.
Table 7.1 lists the formation constants of some complexes. Look at the data, and note
the effect of the different ions and ligands on the value of Kf.
The formation constants in Table 7.1 indicate that both the nature of the metal ion and
the ligand affect the strength of the complex. In general, the strength of a complex increases
with:

increasing charge on the metal ion (3+ ions form stronger complexes than 1+ ions),

increasing size of the metal ion,

more bonding sites per ligand molecule (multidentate ligands form stronger complexes
than monodentates).
Chemistry 2
7.3
7. Complex Formation
TABLE 7.1 Formation constants for selected complexes
Metal
Ligand
Coordination Number
Formation Constant
Cu2+
NH3
4
2 x 1013
Cd2+
NH3
4
1 x 107
Cd2+
CN-
4
6 x 1018
Ag+
EDTA
6
2 x 107
Ca2+
EDTA
6
5 x 1010
Fe2+
EDTA
6
2 x 1014
Al3+
EDTA
6
1 x 1016
Cd2+
EDTA
6
3 x 1016
Cu2+
EDTA
6
6 x 1018
Fe3+
EDTA
6
1 x 1025
Some metal ions have particular attraction for a single ligand, which cause exceptions to the
generalisations above. For example, aluminium and fluoride form a complex - AlF63- - with a
stability constant of 5 x 1019. This is stronger than the complex formed by aluminium and
EDTA.
What do Kf values tell us?
In any aqueous solution of a metal ion (other than the alkali metals), there must be a
concentration of the balancing anions (which are potential ligands), and lots of water - which
is a weak ligand. What does this mean? Any reaction of the metal ion with an added ligand
involves the breaking up of one complex (metal-water or metal-anion) and the formation of a
new, more stable one. Equation 11.3 summarises this process, where S is the solvent or anion
which forms a weak complex with the metal ion, M.
MS + L  ML + S
Eqn 11.3
This also extends to stronger complexes when another ligand is added, which has the capacity
to form a more stable complex with the metal. To summarise: a ligand whose complex
has a higher Kf will cause an existing complex to become unstable. The difference
in Kf values will determine how complete the exchange of ligands would be.
EXAMPLE
A copper (II) solution is mixed with ammonia to form the complex. EDTA is added to this
solution? What happens?
The Kf value for Cu-EDTA is significantly higher (300,000 times greater) than that of copperammonia. So, in the presence of EDTA, Cu(NH3)42+ would fall apart, and the copper would be
complexed with the EDTA.
If two different ligands are added to a solution of a metal ion, most of the cations will bond to
the ligand which forms the more stable complex. The proportion of the two complexes
formed is determined by the difference in the two Kf values. Of course, this is also affected by
the relative concentrations of the ligands.
Chemistry 2
7.4
7. Complex Formation
EXAMPLE
What will happen to a solution of copper (II) ions if EDTA and ammonia are both added?
The copper-EDTA complex is far more stable than copper-ammonia, so the great majority of
the copper ions will bond to EDTA.
PRACTICE QUESTIONS
4.
Explain why the effects of a cyanide spill could be minimised by the addition of Fe3+.
5.
Predict what would happen if:
(a) ammonia was added to a solution of cadmium nitrate
(b) EDTA was added to a solution of cadmium-ammonia complex
(c) ammonia was added to a solution of cadmium-EDTA complex
(d) a small concentration of ammonia was added to a solution containing cadmium
and copper (II) ions
7.4 IMPORTANCE OF COMPLEXES
Metal complexes are extremely important across a variety of areas, from our blood to medical
treatment of certain illnesses to chemical analysis of water. Haemoglobin, which is the red
part of our blood, is a complex between iron (III) and large organic ligands. It is this complex,
which rearranges itself to incorporate oxygen when we breathe, and to release it to the cells.
A large hexadentate ligand - known as desferrioxamine B - is used to treat patients who
cannot excrete iron from their bodies after blood transfusions cause a build up of the metal
ion in the body. The ligand forms a very stable chelate complex with iron, and helps to
remove it.
The ligand EDTA is a most important titrant in chemical analysis, since it has the
capacity to react with many different metals. One of its most significant applications is the
determination of water hardness (caused by Ca2+ and Mg2+), which is a problem in natural
waters, causing soaps not to work, and boilers to be clogged up with ionic salts. EDTA is used
as a „last resort‟ first aid treatment for someone who has swallowed some heavy metal ions.
The EDTA will „mop‟ them up, but also runs the risk of stripping important biochemical
complexes (e.g. co-enzymes) of their metal ions.
Soils contain a number of important ligands called humic acids, which are complex
organic compounds, the structures of which aren‟t fully known. However, it is known that
they are significant in controlling the transport of many metals through the soil, in solution
as a soluble complex or precipitated out as an insoluble one. Iron and aluminium are very
strongly bound, while magnesium is not. These are natural processes, but can be disrupted by
the presence of “unnatural” ligands, which may release certain metals from their humic
complexes.
One major problem associated with “unnatural” complexes is the ability of ligands to
remove heavy metals from insoluble forms, and therefore, allow their transport in solution.
PbX (s) + L (aq)  PbL (aq) + X (aq)
Once mobilised, the heavy metals can then be ingested by organisms at various levels in the
food chain, potentially causing health problems.
Chemistry 2
7.5
7. Complex Formation
WHAT YOU NEED TO BE ABLE TO DO






define terminology associated with complex formation
list the structural feature required for a ligand
briefly describe the formation of the metal-ligand bond
draw the structures of important multidentate ligands, and their complexes with metal
ions
outline factors that affect the stability of a complex
describe some significant applications of complexes
Chemistry 2
7.6