UNIVERSITY OF SARGODHA, SARGODHA
.
.
ASSIGNMENT TOPIC:
Non-Aqueous Solvent
PRESENTED TO:
Dr. Ashraf Shaheen
PRESENTED BY:
Sidra Hameed: CHEM62F20R018
Khadija Masood: CHEM62F20R019
Muhammad Siddique: CHEM62F20R020
Tanzeela Sattar: CHEM62F20R021
Non- Aqueous Solvent
Definition:
The Solvent which are used in non-aqueous titration is called non-aqueous solvent. It is
the solvent other than water that is not the organic compound.
Examples:
Following are the examples of non-aqueous solvent:




Liquid Ammonia (NH3)
Liquid Hydrogen Fluoride (HF)
Liquid Sulphur Dioxide (SO2)
Liquid Dinitrogen Tetraoxide (N2O4)
Classification of Non-Aqueous Solvent:
The solvent have been classified in the number of ways depending on the properties of
non-aqueous solvent. The most convenient classification based on the electrolytic
characteristics of non-aqueous solvents.
The non-aqueous solvents have been classified as follows:
 Non-Ionizing solvent
 Ionizing solvent
This classification is based on “like dissolve like” principle.
Non-Ionizing solvent:
The non-ionizing solvent are non-polar or non-ionic in nature. These solvents dissolve
only non-polar or neutral compounds and donot initiate ionic reactions. These solvents
have low dielectric constant and have little associating tendency between the solute and
solvent. They donot undergo self-ionization.
Examples:
Following are the examples of non-ionizing solvents:
 Benzene (C6H6)
 Carbon Tetrachloride (CCl4)
Ionizing Solvent:
The ionizing solvent are polar or ionic in nature. These solvents dissolve ionic
compounds and initiate ionic reactions. These solvents exist as ions in their pure state
and thus are weak conductors of electricity.
These solvents have high values of dielectric constant. Because of their polar nature,
they have strong tendency to form associated structures. These solvents also show selfionization.
The non-polar substances do not dissolve in in these solvents because of the squeezing
effect produced by their associated structures . However, ionic and covalent polar
substances dissolve in these solvents. The dissolution of the ionic substances is due to
their greater salvation energy than the lattice energy of the salts. Several polar covalent
substances dissolve forming hydrogen bonds.
Examples:
Following are the examples of ionizing solvents:




Ammonia (NH3)
Hydrogen Flouride (HF)
Water (H2O)
Sulphur Dioxide (SO2)
The self-ionization of these solvents are as follows:
H2O + H2O
SO2 + SO2
3HF
⟶
NH3 + NH3
⟶ H3O+ + OH⟶ SO+2 + SO3-2
H2F+ + HF2⟶
NH4+ + NH2-
The non-aqeous solvents have been classified as:
 Amphiprotic or Amphoteric solvents
 Aprotic or non-protic solvents
 Protic or Protonic Solvents
This classification is based on proton-donor or proton-acceptor property.
Amphiprotic or Amphoteric Solvents:
The solvents which have hydrogen in their molecule and can act both as acids and bases
consequently are amphoteric in nature are called amphoteric or amphiprotic solvents.
These solvents undergo auto-ionization in which a proton transfer between two similar
neutral molecules take place and a cation-anion pair of the solvent is obtained. These
solvents show dual character i.e. they can lose as well as accept protons depending on
the nature of reacting species.
Examples:
Following are the examples of amphiprotic or amphoteric solvents:
 Acetic acid (CH3COOH)
 Water (H2O)
Aprotic or Non-protic Solvents:
The Solvents which may or may not have hydrogen in their molecule and neither donate
nor accept protons is called aprotic or non- protic solvents.
Due to self-ionization, they also furnish cations and anions similar to protonic or protic
solvents.
Examples:
Following are the examples of aprotic or non-protic solvents:
 Benzene(C6H6)
 Carbon Tetrachloride (CCl4)
 Sulphur Dioxide (SO2)
Protic or Protonic solvents:
The solvents which have hydrogen as their constituents are called protic or protonic
solvents.
Examples:
Following are the examples of protic or protonic solvents:
 Sulhpuric Acid (H2SO4)
 Acetic Acid (CH3COOH)




Hydrogen Flouride (HF)
Ammonia ( NH3)
Hydrazine (N2H4)
Pyridine
Types of protic or protonic solvents:
Following are the types of protic or protonic solvents:
 Proto-genic or Acidic Solvents
 Proto-philic or Basic Solvents
Proto-genic or Acidic Solvents:
The solvents which have strong tendency to donate protons is called proto-genic or
acidic solvents.
Examples:
Following are the examples of proto-genic or acidic solvents:
 Sulphonic Acid (H2SO4)
 Acetic Acid (CH3COOH)
 Hydrogen Fluoride (HF)
Photophilic or Basic Solvents:
The solvents which have strong tendency to accept protons is called Photophilic or Basic
Solvents.
Examples:
Following are the examples of proto-philic or basic solvents:



Ammonia ( NH3)
Hydrazine (N2H4)
Pyridine
Physical Properties of Solvents and their Role in Chemical Reactions:
Every liquid cannot be
used as a solvent in a chemical reaction. Here we therefore concentrate overselveson the
properties of solvents which make the solvent suitable for using it as a goodsolvent.
Here we shall also compare the properties of some non aqueous solvent with the
properties of H2O.
Some important properties of solvent which make the solvent a useful solvent in a
chemical reactions are as follows:
Melting point and boiling point:
Most of chemical reaction are carried out in liquid phase hence the melting point and
boiling point of a solvent indicate the range of temperature within which the solvent can
be used .Melting point and boiling point (°C) of some solvent are given below :
H2o=0.0,100 ; NH3= -77.7, -33.4 ; SO2= -75.5 , -10.2
HF= 89.3, 19.5 ; CH3COOH =16.6, 118.1
These value shows that H2Ohas a very convenient liquid temperature range between 0
and 100°C . CH3COOH can act as a solvent at ordinary temperature range (16.6°C to
118.1°C). NH3 and SO2 exist gases at ordinary temperature and pressure and hence act as
solvent only at low temperature.
Heat of fusion (ΔHf) and heat of vaporization (∆Hvap):
The heat absorbed by one mole of a substance to change from solid to liquid state called
its molar heat of fusion.
Solid (one mole) →Liquid,
∆H =Heat of fusion =+ ∆Hfus.
Similarly, the heat absorbed by one mole of substance to change form liquid to vapour
state is called is called molar heat of vaporization.
Liquid (one mole) →Vapour,
∆H =Heat of vaporization =+ ∆Hvap
The heat of fusion and vaporization of a solvent indicate the nature strength of forces
with which the molecules of the solvent are held together in the solid or in the liquid
state. A high value of ∆Hvap Of a liquid indicate that the intermolecular forces in the
liquid are strong. The ratio of ∆Hvap( joule) and boiling point, Tb ( in K) of a liquid is a
constant which is known as Troutant constant. i.e.
∆Hvap (JOULE)
= Troutant constant.
Tb (K)
For normal! Liquids the value of Troutant constant is about 90JK-1. Such liquid does not
have any bonds between their molecules. Thus, these liquids exist as independent
molecules.
Value of Troutant constant greater than90JK-1 for liquid (solvent) indicates that the
molecules of the liquid are associated. Since the common solvent like H2O, NH3 and HF
have Troutant constant value greater than 90JK-1 these solvents are associated solvents.
These liquids are polar. The values of ∆Hfus and ∆Hvap (in KJ mol-1) for H2O, NH3, SO2
and HF are great below..
H 2O = 6.02, 40.65 , NH3 = 5.65,23.34.
SO2 = 7.40, 24.93 , HF =4.58, 30.28
Above values show that the value of ∆Hfus for H2O and NH3 are almost the same. This
means that the forces which hold the molecule of H2O and NH3 are of the same
magnitude. Thus these solvents will serve as similar solvent it may also be noted that
since the value of ∆Hfus of SO2 is comparatively.
Dielectric constant (ε):
Dielectric constant (ε) of a solvent determines the ability of the solvent to dissolve polar
and non-polar substances in it. We know that the coulombic force (F) between a cation
and an anion of an ionic compound is given by the expression:
In this expression q1 and q2, are the charges on cation and anion respectively r1 and r2
are the radii of the two ions and e is the dielectric constant of solvent. The value of e
depends on the nature of the solvent in which t ionic compound is dissolved. The value
of dielectric constant of some solvent are given below.
H2O=78.5(25°C)
,
NH3= 22.0 (-33.5°C),
SO2= 17.3(-16°C)
,
HF=83.6(0°C)
N2O4= 2.42(0°C)
,
CH3COOH= 9.7(18°C)
It may be clear from the above expression that if ε of a solvent is large, F would be small
i.e. if εis large, small a mount of energy would be required to separate the ions and
hence it would be easy for a solvent having a high value of εto dissolve an ionic
compound in it. For example, since anhydrous HF and H20 have high value of ε ,these
are the best solvents for ionic compounds. On the other hand, since liq. NH3 and liq.
SO2, have low values of ε, these solvents show amaller ability to dissolve ionic
compounds especially those containing multi-charged ions. Thus carbonates, sulphates
and phosphales which contain multi-charged ions are insoluble in liq. NH3, and liq.SO2,.
Dipole moment:
Greater is the polarity of the bond in a solvent molecule, greater is the charge
separation and higher will be the value for dipole moment, Substances having high
dipole moment values are good solvents for polar solutes.
This is because of the fact that greater is the polartiy of a solvent molecule, greater is the
solvation energy released on dissolution of a solute. Dipole moment value of a solvent
also gives an idea about the extent of association of the molecules of a liquid and hence
its liquid range. Dipole moment values (in D) of H2O, NH3, and SO2, are given below:
H2O = 1.85
, NH3 =1.47
,
SO2 =1.61
Viscosity:
Viscosity gives a measure of the fluidity of the solvent. Solvents like water, carbon
tetrachloride have low viscoscities and flow rapidly under ordinary temperature. In
solvents of low viscosities, the operations such as precipitation, crystallization,
filteration, ete. can be easily carried out without any difficulty. With increasing viscositiy
of a liquid, the difficulty of such operations increases. Solvente like anhydrous sulphuric
acid have higher viscosities and this reduces their usefulness as solvent, Viscosity of
H2O, NH3 and SO2 is 1.00,0. 241 and 0.009 respectivly.
Proton affinity:
It is applicable for protonic solvents only. It greatly affects the behaviour of a solute in a
given solvent. NH 3, has greater proton affinity than H2O. Hence acetamide (CH3 CO
NH2,) which behaves as a very weak base in aqueous solution shows acidic properties in
liq. NH3.
CH3CONH2 +H2O
⇌ CH3 CONH3+ +OH
CH3 + CONH2 +NH3⇌ CH3 COHN +NH4+
Proton affinity (in kJ mol-1) for H2O and NH3 are 760 and 865 respectively. The
properties of H2O, NH3 and SO2 other than those mentioned above are given below:
Solvents
Properties
H2O
NH3
Equivalence
conductance (ohm-1)
6 x 10-8
5 x 10-9
1 x 10-7
374
132.4
157.5
217.7
112.0
77.8
0.96
0.68
1.46
Critical temperature
Critical pressure
Density (g/cc)
SO2
Types of chemical Reactions taking place in non aqueous solvents:
Following are the types of reactions in non-aqueous solvents:
Metathetical (Precipitation) Reactions:
The reactions in which precipitate is formed by mixing two solutions of two compounds
are called metathetical or precipitation reactions. Thus, precipitation reactions are
normally double decomposition. The formation of a precipitate in different solvents
depend on the solubilities of the products in those solvents. For example, the precipitate
of AgCl is obtained by mixing BaCl2, and AGNO3, in aqueous medium.
BaCl2 + AGNO3
AgCl ↓ +Be (NO3)
⟶
AgCl + Be (NO3)
BaCl2↓ + AGNO3
⟶
Thus, we see that reaction (i) is reversed on changing the solvent.
Acid-Base Reactions:
We have already said that the ionic solvents are polar compounds and undergo selfionization. Self-ionization of some important solvents is given below:
CH3COOH + NH3
Acid
base
⇌
NH4+ CH3COO-
acid
base
The accepting ability of H2O less than that of liq. NH 3, CH3COOH behaves as weak acid
in aqueous solution and acts as a strong acid in liq. NH3.
Comparison of acidic strength of HCIO4, HCI, HBr. HI and HNO3, in
aqueous solution:
When the given compounds (represented as H X) react with H,2O. proton is donated by
HX to H2O to produce H3 O because HX molecule has stronger proton donating
property than H2O. Due to the production of H30 ion. HCIO4 HCI, HBr, HI and HNO3,
behave as acids in aqueous medium.
HX + H 2O
→
H 3O+ + X
Stronger
proton donor
In aqueous medium all the given acids dissociate (or ionize) complete and hence behave
as equally strong acids. Thus, H2O acts as a levelling solvent for the given acids. Here it
should be noted that HX and H2O both are protonic substances.
Solvation Reaction:
(Formation of solvates)
The solvation reaction is the general reaction in which a solute (cation, anion or neutral
molecule)reacts with one or more molecules of a solvent(for example, liquid ammonia,
water, liquid Sulphur dioxide) to form a product in which the solute and solvent species
are attached to each other by a hydrogen bond or by a coordinate bond. The product
formed is called solvate. Solvate is an addition compound and hence is called an addict.
The addition compounds contain solvent of crystallization.
In the formation of a solvate, the solvent act as a Lewis base while the solute behaves as
Lewis’s acid.
When the solvent used in water, the solvent reaction is called hydration and the
addition compound formed is called hydrate. Hydrate contains one or more molecules
of water as water of crystallization.
Similarly, when the solvent is liquid ammonia, the reaction is called ammonization and
the addition compound formed is called ammoniate. Ammoniate contains one or more
molecules of ammonia as ammonia of crystallization.
BF3 + NH3 ⟶
BF3.NH3
Ammoniate
KI
+
Lewis base
(solute)
4SO2
⟶
Lewis acid
K+[(SO2 )4I]Solvate
(Solvent)
Solvolytic Reaction:
(Solvolysis)
The solvolysis reaction is the reaction in which the solvent molecules react with the
solute molecules(salt or ion) in the way which consist of the following steps:
The solvent molecules undergo auto-ionization or self-ionization to give solvent cations
and solvent anions.
The solute salt splits in to the solute cations or solute anions interact with the solvent
cations od solvent anions. Due to this interaction, the concentration of the solvent
anions or solvent cations are increased.
When water is used as solvent in the solvolysis reaction, this reaction is called
hydrolysis.
When ammonia is used as solvent in the solvolysis reaction, this reaction is called
ammonolysis or ammonolysis reaction.
In the solvolysis reaction in liquid ammonia, the concentration of either NH4+ or NH2- is
increased.
NH3 + NH3⟶
NH4+ + NH2-
SiCl4 + 8NH3 ⟶ Si(NH2)4 + 4NH4+ + 4ClH- + NH3
⟶ NH-2 + H2
In the solvolysis reaction in liquid Sulphur dioxide, the concentration of either SO+2 or
SO3-2is increased.
SO2 + SO2 ⟶
SO+2 + SO3-2
Zn(C2H5)2 + 2SO2 ⟶
Zn+2 + SO3-2 + (C2H5)2SO
Liquid Ammonia:
Ammonia has a reasonable liquid range (-77 to –33 °C), and as such it can be readily
liquefied with dry ice (solid CO2, Tsub = -78.5 °C), and handled in a thermos flask.
Ammonia’s high boiling point relative to its heavier congeners is indicative of the
formation of strong hydrogen bonding, which also results in a high heat of vaporization
(23.35 kJ/mol). As a consequence ammonia can be conveniently used as a liquid at
room temperature despite its low boiling point.
Physical properties:
Selected physical properties of NH3 are compared with those of water; it has a liquid
range of 44.3 K . The lower boiling point than that of water suggests that hydrogen
bonding in liquid NH3 is less extensive than in liquid H2O and this is further illustrated
by the values of vap Ho(23.3 and 40.7 kJ mol-1 for NH3 and H2O respectively). This is
consistent with the presence of one lone pair on the nitrogen atom in NH3 compared
with two on the oxygen atom in H2O.
The relative permittivity of NH3is considerably less than that of H2O and, as a
consequence, the ability of liquid NH3to dissolve ionic compounds is generally
significantly less than that of water. Exceptions include ½NH4 salts, iodides and
nitrates which are usually readily soluble. For example, AgI which is sparingly soluble in
water dissolves easily in liquid. NH3 (solubility ¼ 206.8 g per 100 g of NH3), a fact that
indicates that both interact strongly with the solvent; Ag forms an ammine complex .
Changes in solubility patterns in going from water to liquid NH3lead to some interesting
precipitation reactions in NH3. Whereas in aqueous solution, BaCl2 reacts with AgNO3
to precipitate AgCl, in liquid NH3, AgCl and BaNO3Þ2react to precipitate BaCl2. Most
chlorides (and almost all fluorides) are practically insoluble in liquid NH3.Molecular
organic compounds are generally more soluble in NH3than in H2O.
Self-ionization:
Liquid NH3 undergoes self-ionization, and the small value of K itself indicates that the
equilibrium lies far over to the left-hand side. The ½NH4 and½NH2 ions have ionic
mobilities approximately equal to those of alkali metal and halide ions. This contrasts
with the situation in water, in which ½H3O and ½O are much more mobile than other
singly charged ions.
Self-ionization of ammonia is much "weaker" than water.
2NH3
⟶ NH+4 + NH-2
2NH3
⟶ NH+4 + NH-2
with K≈10-30K 223K. Since ammonia is better proton acceptor than water, the ionization
of acids is relatively enhanced in liquid ammonia. For example, acetic acid is a strong
acid in liquid ammonia. Liquid ammonia will therefore tolerate very strong bases such
as C5H-5C5H-5 that would otherwise be hydrolyzed in water.
Ammonia is kinetically stabilized to reduction (but easily oxidized) by many reagents
.For example, the reaction
Na+NH3⟶NaNH2+H2
Na+NH3⟶NaNH2+H(g)
and is very favorable but slow in the absence of a catalyst such as Fe3+.
Solubility of various substances in liquid ammonia:
Solubility of ioinic compounds or organic salts:
Liquid ammonia is a good solvent for organic molecules (e.g., esters, amines, benzene,
and alcohols). It is a better solvent for organic compounds than water, but a worse
solvent for inorganic compounds. The solubility of inorganic salts is highly dependant
on the identity of the counter ion.
Soluble in liquid NH3
SCN-, I-, NH4+, NO3-, NO2-, ClO4-
Insoluble in liquid NH3
F-, Cl-, Br-, CO3-2, SO4-2, O-2, OH-, S-2
General solubility of inorganic salts in liquid ammonia as a function of the counter
ion.The difference in solubility of inorganic salts in ammonia as compared to water, as
well as the lower temperature of liquid ammonia, can be used to good advantage in the
isolation of unstable compounds. For example, the attempted synthesis of ammonium
nitrate by the reaction of sodium nitrate and ammonium chloride in water results in the
formation of nitrogen and water due to the decomposition of the nitrate.
By contrast, if the reaction is carried out in liquid ammonia, the sodium chloride side
product is insoluble and the ammonium nitrate may be isolated as a white solid after
filtration and evaporation below its decomposition temperature of 0 °C.
Solubility of non-ionic compounds or organic salts:
Halogen compounds, alcohols, ketones, esters, simple ethers, amines, phenols and its
derivates etc. are soluble. Alkanes are insoluble and alkenes and alkynes are slightly
soluble. In this sense liquid ammonia is a better solvent for non-ioinc and non- polar
compounds or organic salts.
Solubility of non-metals:
The non-metals like S, P , Se, I2 etc are soluble and they react with the solvents.
Self-ionization of ammonia is much "weaker" than water.
2NH3 ⟶NH+4+NH−2
2NH3 ⟶NH4++NH2−
with K≈10−30K 223K. Since ammonia is better proton acceptor than water, the
ionization of acids is relatively enhanced in liquid ammonia. For example, acetic acid is
a strong acid in liquid ammonia. Liquid ammonia will therefore tolerate very strong
bases such as C5H−5C5H5− that would otherwise be hydrolyzed in water.
Ammonia is kinetically stabilized to reduction (but easily oxidized) by many reagents,
e.g., the reaction
Na+NH3⟶NaNH2+H2
Na+NH3⟶NaNH2+H2(g
and is very favorable but slow in the absence of a catalyst such as Fe3+.
Chemical reactions taking place in liquid ammonia:
Following are the chemical reactions taking place in liquid ammonia:
Ammonation:
Ammonation is defined as a reaction in which ammonia is added to
other molecules or ions by covalent bond formation utilizing the unshared pair of
electrons on the nitrogen atom, or through ion-dipole electrostatic interactions. In
simple terms the resulting ammine complex is formed when the ammonia is acting as a
Lewis base to a Lewis acid,or as a ligand to a cation, e.g., [Pt(NH3)4]+2, [Ni(NH3)6]+2,
[Cr(NH3)6]+3, and [Co(NH3)6]+3
SiF4 + 2NH3 ⟶
BF3 +NH3 ⟶
SiF4(NH3)2
BF3(NH3)
Ammonolysis or Ammonolytic reactions:
Ammonolysis with ammonia is an analogous reaction to hydrolysis with water i.e. a
dissociation reaction of the ammonia molecule producing H+ and an NH2- species.
Ammonolysis reactions occur with inorganic halides,and organometallic compounds,
Equation. In both case the NH2- moiety forms a substituent or ligand.
The reaction of esters, Equation, and aryl halides, Equation, are also examples of
ammonolysis reactions
The reaction of ester and aryl halides are also examples of ammonolysis reactions.
Homoleptic amide:
A homoleptic compound is a compound with all the ligands being identical,. For
example, M(NH2)n. A general route to homoleptic amide compounds is accomplished
by the reaction of a salt of the desired metal that is soluble in liquid ammonia (Table)
with a soluble Group 1 amide. Since all amides are insoluble (except those of the Group 1
metals) are insoluble in liquid ammonia, the resulting amide may be readily isolated.
Amide
LiNH2
NaNH2
KNH2
RbNH2
CsNH2
Solubility in liquid ammonia
Sparingly soluble
Sparingly soluble
Soluble
Soluble
Soluble
Redox reactions:
Ammonia is poor as an oxidant since it is relatively easily oxidized. Thus, if it is
necessary to perform an oxidation reaction ammonia is not a suitable solvent; however,
it is a good solvent for reduction reactions.
Liquid ammonia will dissolve Group 1 (alkali) metals and other electropositive metals
such as calcium, strontium, barium, magnesium, aluminum, europium, and ytterbium.
At low concentrations (Ca: 0.06 mol/L), deep blue solutions are formed. these contain
metal cations and solvated electrons. The solvated electrons are stable in liquid
ammonia and form a complex.
[e-(NH3)6].
The solvated electrons provide a suitable and powerful reducing agent for a range of
reactions that are not ordinarily accomplished.
A hydrated ion is one kind of a complex ion, a species formed between a central metal
ion and one or more surrounding ligands , molecules or ions that contain at least one
lone pair of electrons, such as the [Al(H2O)6]3+ ion.
A complex ion forms from a metal ion and a ligand because of a Lewis acid–base
interaction. The positively charged metal ion acts as a Lewis acid, and the ligand, with
one or more lone pairs of electrons, acts as a Lewis base. Small, highly charged metal
ions, such as Cu2+ or Ru3+, have the greatest tendency to act as Lewis acids, and
consequently, they have the greatest tendency to form complex ions.
As an example of the formation of complex ions, consider the addition of ammonia to an
aqueous solution of the hydrated Cu2+ ion {[Cu(H2O)6]2+}. Because it is a stronger base
than H2O, ammonia replaces the water molecules in the hydrated ion to form the
[Cu(NH3)4(H2O)2]2+ ion. Formation of the [Cu(NH3)4(H2O)2]2+ complex is accompanied
by a dramatic color change. The solution changes from the light blue of [Cu(H2O)6]2+ to
the blue-violet characteristic of the [Cu(NH3)4(H2O)2]2+ ion.
An aqueous solution of CuSO4 consists of hydrated Cu2+ ions in the form of pale blue
[Cu(H2O)6]2+ (left). The addition of aqueous ammonia to the solution results in the
formation of the intensely blue-violet [Cu(NH3)4(H2O)2]2+ ions, usually written as
[Cu(NH3)4]2+ ion (right) because ammonia, a stronger base than H2O, replaces water
molecules from the hydrated Cu2+ ion.
The Formation Constant:
The replacement of water molecules from [Cu(H2O)6]2+ by ammonia occurs in
sequential steps. Omitting the water molecules bound to Cu2+ for simplicity, we can
write the equilibrium reactions as follows:
Cu2+(aq)+NH3(aq)[Cu(NH3)]2+(aq)+NH3(aq)[Cu(NH3)2]2+(aq)+NH3(aq)[Cu(NH3)
3]2+(aq)+NH3(aq)⇌[Cu(NH3)]2+(aq)
K1⇌[Cu(NH3)2]2+(aq)
K2⇌[Cu(NH3)3]2+(aq)
K3⇌[Cu(NH3)4]2+(aq)
K4⇌Cu2+(aq)+NH3(aq)⇌[Cu(NH3)](aq)2
K1[Cu(NH3)](aq)+NH3(aq)⇌[Cu(NH3)2](aq)2+K2([Cu(NH3)2](aq)2+NH3(aq)⇌[Cu(
NH3)3](aq)2+K3[Cu(NH3)3](aq)2+NH3(aq⇌[Cu(NH3)4](aq)2+K4
The sum of the stepwise reactions is the overall equation for the formation of the
complex ion: The hydrated Cu2+ ion contains six H2O ligands, but the complex ion that
is produced contains only four NH3NH3 ligands, not six.
Cu2+(aq)+4NH3(aq)⇌[Cu(NH3)4]2+(aq)
Cu(aq)2++4NH3(aq)⇌[Cu(NH3)4](aq)2+
The equilibrium constant for the formation of the complex ion from the hydrated ion is
called the formation constant (Kf). The equilibrium constant expression for Kf has the
same general form as any other equilibrium constant expression. In this case, the
expression is as follows:
Kf=[[Cu(NH3)4]2+][Cu2+][NH3]4=2.1×1013=K1K2K3K4
Kf=[[Cu(NH3)4]2+][Cu2+][NH3]4=2.1×1013=K1K2K3K4
The formation constant (Kf) has the same general form as any other equilibrium
constant expression.
Water, a pure liquid, does not appear explicitly in the equilibrium constant expression,
and the hydrated Cu2+(aq) ion is represented as Cu2+ for simplicity. As for any
equilibrium, the larger the value of the equilibrium constant (in this case, Kf), the more
stable the product. With Kf = 2.1 × 1013, the [Cu(NH3)4(H2O)2]2+ complex ion is very
stable.