Non-Aqueous Solvent
Submitted to: Dr. Ashraf Shaheen
Submitted by:Group Number 5
Group Members:
1)
2)
3)
Sidra Hameed: CHEM62F20R018
Khadija Masood: CHEM62F20R019
Muhammad Siddique: CHEM62F20R020
4)
Tanzeela Sattar: CHEM62F20R021
Non- Aqueous Solvent
Definition:
The Solvent which are used in non-aqueous titration is called nonaqueous 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 Tetroxide (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.
1. 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-he

Benzene (C6H6)

Carbon Tetrachloride (CCl4)
2. 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 self-ionization.
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
⟶
H3O+ + OHSO+2 + SO3-2
⟶
3HF
⟶
NH3 + NH3
⟶
H2F+ + HF2NH4+ + 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.
1. 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:

Acetic acid (CH3COOH)

Water (H2O)
2. 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)
3. 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 overselves on the properties of solvents which
make the solvent suitable for using it as a good solvent. 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 discussed below .
1. 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 .
2. 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 Trouton constant. i.e
∆Hvap (JOULE)÷ Tb (K) = Troutant constant.
For normal! Liquids the value of Troutant constant is about 90 JK-1. Such
liquid does not have any bonds between their molecules. Thus, these
liquids exist as independent molecules.
Value of Troutant constant greater than 90 JK-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 90 JK-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..
3. 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), CH3COOH= 9.7(18°C)
HF= 83.6(0°C),
N2O4 = 2.42(0°C)
It may be clear from the above expression that if ε of a solvent is large, F
would be small i.e. if ε is 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 ε to 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,.
4. 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
5. 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 viscotiy 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.
6. 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 + CONH2 +NH3
⇌ CH3 CONH3+ +OH
⇌
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
Equivalence
H2 O
NH3
SO2
6 x 10-8
5 x 10-9
374
132.4
157.5
217.7
112.0
77.8
0.96
0.68
1.46
1 x 10-7
conductance (ohm-1)
Critical
temperature
Critical pressure
Density (g/cc)
Types of chemical Reactions taking place
in non aqueous solvents
1. 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
AgCl ↓
⟶
+ Be (NO3) ⟶
+ Be (NO3)
BaCl2 ↓ + AGNO3
Thus, we see that reaction (i) is reversed on changing the solvent.
2. Acid-Base Reactions
We have already said that the ionic solvents are polar compounds and
undergo self-ionization. Self-ionization of some important solvents is
given below:
CH3 COOH + NH3
Acid
base
⇌
NH4
acid
+ CHCOO_
base
Now since the proton accepting ability of H2O E tess 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 H 2O to produce H3 O because HX molecule has
stronger proton donating property than H2O. Due to the production of H
3
0 ion. HCIO 4 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
3. 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.
4. 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
SiCl4 + 8NH3 ⟶
H- + NH3
⟶
NH4+ + NH2-
Si (NH2)4 + 4NH4+ + 4Cl⟶
NH-2 + H2
In the solvolysis reaction in liquid Sulphur dioxide, the
concentration of either SO+2 or SO3-2 is increased.
SO2 + SO2 ⟶
Zn (C2H5)2 + 2SO2
⟶
SO+2 + SO3-2
Zn+2 + SO3-2 + (C2H5)2SO
Reaction of 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.
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
Generally insoluble in liquid NH3
NH3
F-, Cl-, Br-, CO32-, SO42-, O2-, OH-, S2SCN-, I-, NH4+, NO3-, NO2-, ClO4General 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
Equation
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
Reaction
Ammonolysis
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.
Reaction
The reaction of esters, Equation, and aryl halides, Equation, are also examples of
ammonolysis reactions
Homoleptic amides
A homoleptic compound is a compound with all the ligands being identical, e.g.,
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. The solubility of the Group 1 amides is given in
Table. Since all amides are insoluble (except those of the Group 1 metals) are
insoluble in liquid ammonia, the resulting amide may be readily isolated.
Rection
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.
Equation
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].
Equation
The solvated electrons provide a suitable and powerful reducing agent for a range
of reactions that are not ordinarily accomplished.
Equation