Lecture 8
Complex compounds.
Structure of complex
compounds.
Associate prof. Yu. B. Dmukhalska
Outline
1. Solubility. The mechanism of dissolving.
2. Solubility of gases in liquids. The Henry’s law.
3. Colligative properties:
a) osmosis. The vant’-Hoff’s law. Hemolysis and plasmolysis;
b) vapor-pressure lowering of solution. A Raoult’s law;
c) boiling-point elevation;
d) freezing-point depression.
4. Concept of complex compounds and complexing process.
Nomenclature of complex compounds. Types of complexes.
5. Structure of complex compounds. Isomerism of complex
compounds. Chemical bonds in complex compounds molecule.
6. Stability of complexes and influence of different factors on it.
7. Biological role of complex compounds. Usage of complexing
in chemistry.
Water –
main solvent
The water - main component of organisms
and medium, in which lives the person. The
main properties of water that in water can
solubility a different matters.
In the human, animal, plant organisms the
water is main part, a constituent solvent and
it participates in exchange reactions of
matters (hydrolysis, hydration, swelling,
digestion).
In a human organism are about 70 - 80 % of
water.
The mechanism of
dissolving
(- +)Polar molecule
(-)
Negative ion
(+) Positive ion
(-+)Water dipole
The dissolving depends primarily on the
relative strengths of three attractive
forces:
1) the forces between the particles of the
solute before it has dissolved {solutesolute forces),
2) the forces between solvent particles
before dissolution has taken place
(solvent-solvent forces),
3) the forces that are formed between
solute and solvent particles during the
dissolving
process
(solute-solvent
forces).
Type of solution
A saturated solution is one that is in equilibrium with excess
undissolved solute, or would be in equilibrium if excess
solute were present. The term saturated denotes the highest
concentration of solute which a solution can have and be in
equilibrium with any undissolved solute with which it is
placed in contact.
An unsaturated solution is one in which the concentration of
solute is less than its concentration in a saturated solution.
A supersaturated solution is one in which the concentration of
solute is greater than its concentration in a saturated
solution.
A supersaturated solution is unstable and its solute tends
eventually to crystallize out of solution, much as a super
cooled liquid tends eventually to crystallize.
 Gas solution is not possible to prepare a heterogeneous
mixture of two gases because all gases mix uniformly with
each other in all proportions. Gaseous solutions have the
structure that is typical of all gases. Air, the gaseous
solution with which we come in closest contact, is
composed primarily of N2 (78 % by volume), O2 (21 %),
and Ar (1 %), with smaller concentrations of CO2, H2O,
Ne, He, and dozens of other substances at very low
levels.
 Liquid solutions have the internal structure that is typical
of pure liquids: closely spaced particles arranged with little
order. Unlike a pure liquid, however, a liquid solution is
composed of different particles. Much of this chapter is
devoted to the properties of liquid solutions, and special
emphasis is given to aqueous solutions, in which the
major component is water.
 Two kinds of solid solutions are common. The first, the
substitutional solid solution, exhibits a crystal lattice that
has structural regularity but in which there is a random
Disperse systems
Disperse
systems
are
called
systems, which consist of two
phases, one of which is scattered or
dispersed in other.
The disperse phase - phase which is
scattered (dispersed) in medium.
The disperse medium - phase in
which dispersion done.
Classification disperse
systems
• By stat of dispersed phase and dispersed medium
By size of dispersed phase
 Colloidal solutions are
disperse systems, which
have dispersed phase
particle, which size
-9
-7
between 10 to 10 m or 1
nm to 100 nm.
Concentration units of
solution
 Mass fraction (i) of solute in solution is the ratio of
the mass solute (mi) to the mass of solution mi +ms; msmass of a solvent:
 Percentage by weight (mass) or mass percent, is the
quantity of one component of a solution expressed as a
percentage of the total mass:
 where m - percent by mass,
 mA, mB, mC - mass of components in the solution.


Mass concentration, titer (T) is number grams of solute (m) per one milliliter
of solution (V). Or it is the ratio of the quantity grams of solute and volume
solution:
T= m.
V
Molarity (CM), or molar concentration, is the number of moles of solute
dissolved per liter of solution.
CM =
γ
V
where: CM
= m .
MV
- molarity (by mole of solute per liter of a solution);
γ - number mole solute;
m - mass solute, grams;
M - molar mass solute, in grams/mole;
V - volume of the solution;

Molality is defined as the number of moles (γ) of solute dissolved per
kilogram of solvent. Thus, the molality of solute in a solution is
Cm =
γ
msolvent
=
msolute
;
M solute msolvent
when Cm – molality (by mole of solute per kilogram of solvent);
γ - number of moles of solute;
m – mass of solvent.
 In measure analysis for the characteristic the composition of
solution will use molar mass of an equivalent (equivalent mass)
 Molar mass of an equivalent of element is the mass of the element
which combines with or displaces 1.008 parts by mass of hydrogen
or 8 part by mass of oxygen or 35.5 parts by mass of chlorine: E =
fequivalence · MB
 The factor of equivalence (fequiv) - number, which is demonstrated
which part of matter (equivalent) can react with one atom of
Hydrogen, or one electron in reduction reactions.
 Molar concentration of an equivalent (normal concentration),
normality is quantity gram-equivalent of solute per one liter of
solution (V):
 Ceq=
γeq
= m .
V
EV
where: CM - molarity (by mole of solute per liter of a solution);
γeq - number mole-equivalent of solute;
m - mass solute, grams;
E - molar mass of an equivalent solute (equivalent mass of solute);
V - volume of the solution;
Henry's Law:
The solubility of a gas dissolved in a liquid is
proportional to the partial pressure of the gas
above the liquid.
This is a statement of Henry's law, which can be
written
X = KP
X is the equilibrium mole fraction of the gas in
solution (its solubility)
P is its partial pressure in the gas phase
K - constant of proportionality or Henry's-law
constant.
 The partial pressure is a part of common
pressure, which one is a share of each gas in
gas mixture.
Properties of a solution which
depend only on the concentration
of the solute and not upon its
identity are known as colligative
properties.
 vapor-pressure lowering
 boiling-point elevation
 freezing-point depression
 osmotic pressure.
The spontaneous mixing of the particles of the solute
(present in the solution) and the solvent (present
above the solution) to form а homogeneous mixture
is called diffusion, just as the term is used for the
spontaneous mixing of gases to form homogeneous
mixtures.
A semi-permeable membrane - а membrane which
allows the solvent molecules to pass through but not
the solute particles.
The net spontaneous flow of the solvent molecules from
the solvent to the solution or from a less concentrated
solution to а more concentrated solution through а
semi-permeable membrane is called osmosis
(Greek: push).
 The osmotic pressure of а solution may thus
be defined as the equivalent of excess
pressure which must be applied, to the
solution in order to prevent the passage of
the solvent into it through а semi-permeable
membrane separating the two, i.e. the
solution and the pure solvent.
 Osmotic pressure may be defined as the
equilibrium hydrostatic pressure of the
column set up as а result of osmosis.
 Р С; Т; Р СТ or P=RCT
 PV= nRT – van’t Hoff equation for dilute
solutions
Laws of osmotic pressure These are the same as gas
laws and apply to dilute
solutions which occur in the
living body
The effect of hypertonic and hypotonic solutions on
animal cells.
(а) Hypertonic solutions cause cells to shrink
(crenation);
(b) hypotonic solutions cause cell rupture;
(c) isotonic solutions cause no changes in cell
volume.
The partial vapor pressure of a component in liquid
solution is proportional to the mole fraction of that
component, the constant of proportionality being the
vapor pressure of the pure component.
Raoult's law can be written as
P1 = X1 P10
where P1 and P10 are the vapor pressure of the solution
and that of the pure solvent, respectively,
X1 is the mole fraction of the solvent in the solution.
P1 is the total vapor pressure of the solution.
X2 = 1 - X2,
P1 = (1- X2)P10
P10 - P1 is the vapor-pressure lowering
P10 - P1
---------- = X2 fractional vapor-pressure lowering
P10
which can be seen to be equal to the mole fraction of the
solute - X2.
The relationship between boiling-point elevation
and solute concentration: it can be shown that in
dilute solutions the boiling-point elevation is
proportional to the molality of the solute
particles.
if Tb, represents the boiling-point elevation:
Tboiling =Tboiling (solution) - Tboiling (solvent),
Tb = KbCm
Cm = molality, number of mole of solute per one
kilogram of solvent
Where: Cm - molality of the solute in solution
Kb- proportionality constant known as the molal
boiling-point elevation constant.
The relationship between freezing-point
depression and molality in dilute
solutions is a direct proportionality
Tf = Tfreezing(solvent) - Tfreezing(solution) freezing-point depression
Tfreezing= KfCm
Where: Cm - molality of solute;
Kf - molal freezing-point depression
constant
Complex compounds. Classification of
complex compounds.
Complexes are multiple objects, which are formed of
more simple objects (ions, molecules), capable to
independent existence in solutions.
Coordination compounds are the compounds in which
the central metal atom is linked to а number of ions
or neutral molecules by coordinate bonds i.е. by
donation of lone pairs of electrons by these ions or
neutral molecules to the central metal atom.
Complexing – it is a process of complex compounds
formation from more simple objects.
 The term complex in chemistry is usually used to
describe molecules or ensembles formed by the
combination of ligands and metal ions.
 The molecules or ions that surround the central
metal ion in a coordination compound are called
ligands, and the atoms that are attached directly
to the metal are called ligand donor atoms.
 The number of ligand donor atoms that surround
a central metal ion in a complex is called the
coordination number of the metal
 Originally, a complex implied a reversible
association of molecules, atoms, or ions through
weak chemical bonds.
 Some important characteristics of chelates.
 (i) Chelating ligands form more stable
complexes than the monodentate analogs. This
is called chelating effect.
 (ii) Chelating ligands, which do not contain
double bonds e.g. ethylenediamine form five
membered stable rings. The chelating ligands
such as acetylacetone form six membered
stable ring complexes.
 (iii) Ligands with large groups form unstable
rings than the ligands with smaller groups due to
steric hindrance.
 Coordination number. The total number
of monodentate ligands (plus double the
number of bi dentate ligands if any)
attached to the central metal ion through
coordinate bonds is called the
coordination number of the metal ion.
 [Ag(СN)2]-, [Cu(NН3)4]2+ and [Cr(Н2О)6]3
Coordination sphere.
 The central atom and the ligands which
are directly attached to it are enclosed in
square brackets and are collectively
termed as the coordination sphere.
Oxidation number or oxidation state.
 It is а number that represents an electric charge
which an atom or ion actually has or appears to
have when combined with other atoms,
 oxidation number of copper in [Cu(NH3)4]2+ is +2
but coordination number is 4.
 oxidation number of Fe in [Fe(СN)6]3- is + 3 but
the coordination number is 6.
 (i) [Cu (NНЗ)4]SO4.
 (ii) Fe in [Fe (СN)6]3 (iii)К3[Fe(С2О4)3].
 (iv) [Ni(CO)4].
Charge on the complex ion.
 The charge carried by а complex ion is the
algebraic sum of the charges carried by
central metal ion and the ligands
coordinated to the central metal ion.
 [Ag (CN)2] [Cu (NH3)4]2+
Co-ordination Werner’s theory
Charge
+1
+2
+3
+4
coordination number
example
of the metal ion
2
Ag+, Cu+
4, 6
Cu2+, Zn2+, Pd2+, Pt2+
6, 4
Pt4+, Cr3+, Co3+, Fe3+
8
Sn4+
Aqueous solutions that contain [Ni(H2O)6]2+, [Ni(NH3)6]2+ and [Ni(en)3]2+
(from left to right). The two solutions on the right were prepared by
adding ammonia and ethylenediamine, respectively, to aqueous
nickel(II) nitrate.
Naming Coordination Compounds
Names of Some Common Metallate
Anions
Names of Some Common Ligands
Examples of Complexes with
Various Coordination Numbers
 Ligands have at least one lone pair of electrons
that can be used to form a coordinate covalent
bond to a metal ion.
 They can be classified as monodentate or
polydentate, depending on the number of ligand
donor atoms that bond to the metal.
Ligands such as H2O, NH3 or Cl- that bond using the
electron pair of a single donor atom are called
monodentate ligands (literally, “onetoothed”
ligands).
 Those that bond through electron pairs on more than
one donor atom are termed polydentate ligands
(“many-toothed” ligands).
For example, ethylenediamine (NH2CH2CH2NH2
abbreviated en) is a bidentate ligand because it
bonds to a metal using an electron pair on each of its
two nitrogen atoms.
The hexadentate ligand ethylenediaminetetraacetate
ion (EDTA4-) bonds to a metal ion through electron
pairs on six donor atoms (two N atoms and four O
atoms).
Structures of some common ligands
Types of complex:
1. Ionic associates (ionic pairs) in solutions are
formed as a result only electrostatic interaction
between opposite charged ions, for example
Kt+ + An-[Kt+, An-]
+
(CH3)2N
N(CH3)2
C
-
[SbCl6] +
Malachite green
+
(CH3)2N
N(CH3)2
C
[SbCl6]-
 2. Complexes without the coordination
centre
Hydroquinone
Quinhydrone
Quinone
 3. Coordination complex compounds
Coordination complex
compounds:
1. One-nuclear complexes

One-ligandly: metallamine [Cu(NH3)4]SO4
aquacomlexes [Co(H2O)6]Cl2
acidocomplexes K2[PtCl4]; H2[SiF6];
 Combination-ligandly: [Pt(NH3)Cl2];
[Pt(NH3)Cl3].
2. Poly-nuclear complexes
 bridging complex [Cr(NH3)5-OH-(NH3)5Cr]Cl5
 cluster complex
Br
Br
Br
Re
Br
2-
Br
Re
Br
Br
Br
 isopoly acids
Н4Р2О7, Н2В4О7
 heteropoly acids
H3PO4·12MoО3·nН2O
H3PO4·12WО3·nН2O
H4SiО4·12MoО3·nН2O
H4SiО4·12WО3·nН2O
 A complex such as [Co(en)3]3+ or Co(EDTA)]- that
contains one or more chelate rings is known as a metal
chelate.
The resulting five-membered ring consisting of the Co(III)
ion, two N atoms, and two C atoms of the ligand is called
a chelate ring.
[Co(en)3]3+
Co(EDTA)]-
Scheme of copper chelation [Cu(NH3)4]2+
Octahedral structure of the
[Co(NH3)6]3+
Idiosyncrasy of chelate – it is presence
of cycles.
Diethylenediaminocopper (ІІ)
Diglycinatocopper (ІІ)
active site of chlorophyll
active site of hemoglobin
hemoglobin
Structure of molecule of cyancobalamin
(vitamin В12)
Mechanism of action Tetacinum-calcium
Ions Hg2+ and Cd2+ displace ions Ca2+ from
Tetacinum
Color changes produced by adding various reagents to an
equilibrium mixture of Fe3+ (pale yellow), SCN- (colorless),
and FeNCS2+ (red): (a) The original solution. (b) After
adding to FeCl3 the original solution, the red color is darker
because of an increase in [FeNCS2+]. (c) After adding
KSCN to the original solution, the red color again deepens.
(d) After adding H2C2O4 to the original solution, the red
color disappears because of a decrease in [FeNCS2+] the
yellow color is due to Fe(C2O4)33-. (e) After adding HgCl2
to the original solution, the red color again vanishes.
Necessary parts of ligands for chelate
formation
1. Functional-analytical groups (FAG) - are
specific groups which provide occurrence of
donor-acceptor bond.
-ОН, -SH, =NH, -COOH, -SO3H, -AsО3H2,
C=Ö: і т.д.
2. Analytical-active groups (ААG) – are the groups
of atoms which change analytical properties of
reaction products (solubility, intensity of
colouring).
Auxochrome - this is a group of atoms attached to a
chromophore which modifies the ability of that
chromophore to absorb light.
An auxochrome is a functional group of atoms with
nonbonded electrons which, when attached to a
chromophore, alters both the wavelength and
intensity of absorption.
If these groups are in direct conjugation with the
pi-system of the chromophore, they may increase
the wavelength at which the light is absorbed and
as a result intensify the absorption (-Cl, -Br, -J, C6H5).
A feature of these auxochromes is the presence of at
least one lone pair of electrons which can be
viewed as extending the conjugated system by
resonance. Also that groups which improve
solubility of complexes (-SO3H,-COOH).
Process of complexing
stepwise fashion
cumulative (common)
Me + L ↔ MeL
Me + L ↔ MeL
MeL + L ↔ MeL2
Me + 2L ↔ MeL2
MeL2 + L ↔ MeL3
Me + 3L ↔ MeL3
··································
··································
MeLn-1+ L ↔ MeLn
Me + n L ↔MeLn
The formation of a metal–ligand complex is
described by a formation constant, Kf.
Process of complex dissociate
stepwise fashion
MeLn  MeLn-1+ L
MeLn-1 MeLn-2+ L
…………………….
MeL2  MeL + L
MeL Me + L
cumulative (common)
MeLn  Me + nL
МeLn-1  Me + (n-1)L
……………………..
MeL2  Me + 2L
MeL  Me +L
The reverse of reaction complexing is called a
dissociation reaction and is characterized by a
dissociation constant, Kd
 Stepwise formation constants
The formation constant for a metal–ligand
complex in which only one ligand is added
to the metal ion or to a metal–ligand
complex (Ki)
 Cumulative formation constant
The formation constant for a metal–ligand
complex in which two or more ligands are
simultaneously added to a metal ion or to a
metal–ligand complex (βi).
For example, the reaction between Cd2+ and
NH3 involves four successive reactions
So
Relationship between Kf() and Kd
Me + nL ↔MeLn
[MeL ]
 
[Me]  [L]
MeLn↔Me + nL
'
d
n
n
[ Me ]  [ L ]
K 
[ MeLn ]
n
n
1
n  '
Kd
 β (Kf) - formation constant (or stability constant)
! So, Kd, which is the reciprocal of Kf.
2. Stability of complexes and influence
of different factors on it.
Kinetic stability:
 Labile complexes
 Inert complexes
Thermodynamic stability:
 formation constant (dissociation constant)
Factors which influence stability of complex
connections:
 The ion nature of metal and ligand;
 The charge of an metal ion;
 Ionic radius of the metal-complexing agent;
 The nature of medium.
Influence of different factors on
complexing in solution.
1. Ionic strength of solution
2. рН
3. concentration of ligand
4. temperature
5. stranger ions, which form slightly soluble
compound with metal-complexing agent or
ligand.
3. Influence of complexing on precipitate
solubility and oxidation-reduction
potential of system.
 the solubility of precipitate increases
 oxidizing and reducing properties of redoxpair can increase or decrease (depending
on the nature of comlexes, which will form
with oxidizing and reduction redox-pair
forms)










4. Usage of complexing in analytical
chemistry.
masking of іоns
determination of cations and anions
separation
concentrating and determination of ions
precipitation of cations and anions from the solutions
dissolution of precipitate
definition identity of drugs on functional groups
change red-ox potential
determination of ions by fluorescence analysis
for fixing of equivalence point in titrimetric analysis
The qualitative analysis
Silver chloride is insoluble in water (left) but
dissolves on addition of an excess of
aqueous ammonia (right).
Chelatometry
Complexon І: nitrilotriacetic acid
(tetradentate)
CH2-COOH
HOOC-CH2-N
CH2-COOH
Complexon ІІ: (EDТА)
ethylenediaminetetraacetic acid
HOOC-CH2
..
..
N-CH2-CH2-N
HOOC-CH2
CH2-COOH
CH2-COOH
Complexon ІІІ: sodium
ethylenediaminetetraacetate (Na-EDТА, trylon
B, chelaton) - Na2H2Y
NaOOC-CH2
HOOC-CH2
..
..
N-CH2-CH2-N
CH2-COOH
CH2-COONa
Complexon
acid
ІV:
cyclohexyldiaminetetraacetic
CH2-COOH
N
CH2-COOH
CH2-COOH
N
CH2-COOH
 All metal-EDTA complexes have a 1:1
stoichiometry.
 These complexes are dissolved in water.
 Metal-EDTA complexes are named – metal
complexonate.
Thanks for your attention!