| Ideal solutions | Non-ideal solutions | ||
| Positive deviation from Raoult's law | Negative deviation from Raoult's law | ||
| 1. Obey Raoult's law at every range of concentration. | 1. Do not obey Raoult's law. | 1. Do not obey Raoult's law. | |
| 2. \[\Delta {{H}_{\text{mix}}}=0;\]neither heat is evolved nor absorbed during dissolution. | 2. \[\Delta {{H}_{\text{mix}}}>0.\]Endothermic dissolution; heat is absorbed. | 2. \[\Delta {{H}_{\text{mix}}}<0.\] Exothermic dissolution; heat is evolved. | |
| 3. \[\Delta {{V}_{\text{mix}}}=0;\]total volume of solution is equal to sum of volumes of the components. | 3. \[\Delta {{V}_{\text{mix}}}>0.\]Volume is increased after dissolution. | 3. \[\Delta {{V}_{\text{mix}}}<0.\] Volume is decreased during dissolution. | |
| 4. \[P={{p}_{A}}+{{p}_{B}}=p_{A}^{0}{{X}_{A}}+p_{B}^{0}{{X}_{B}}\]i.e., \[{{p}_{A}}=p_{A}^{0}{{X}_{A}}:{{p}_{B}}=p_{B}^{0}{{X}_{B}}\] | 4. \[{{p}_{A}}>p_{A}^{0}{{X}_{A}};\] \[{{p}_{B}}>p_{B}^{0}{{X}_{B}}\] \[\therefore \] \[{{p}_{A}}+{{p}_{B}}>p_{A}^{0}{{X}_{A}}+p_{B}^{0}{{X}_{B}}\] | 4. \[{{p}_{A}}<p_{A}^{0}{{X}_{A}};\]\[{{p}_{B}}<p_{B}^{0}{{X}_{B}}\] \[\therefore \] \[{{p}_{A}}+{{p}_{B}}<p_{A}^{0}{{X}_{A}}+p_{B}^{0}{{X}_{B}}\] | |
| 5. \[A-A,\ A-B,\ B-B\] interactions should be same, i.e., 'A' and 'B' are identical in shape, size and character. | 5. \[A-B\] attractive force should be weaker than \[A-A\] and \[B-B\] attractive forces. 'A' and 'B' have different shape, size and character. | 5. \[A-B\]attractive force should be greater than \[A-A\] and \[B-B\] attractive forces. 'A' and 'B' have different shape, size and character. | |
| 6. Escaping tendency of 'A' and 'B' should be same in pure liquids and in the solution. | 6. 'A' and 'B' escape easily showing higher vapour pressure than the expected value. | 6. Escaping tendency of both components 'A' and 'B' is lowered showing lower vapour pressure than expected ideally. | |
| Examples: Dilute solutions; benzene + toluene: n-hexane + n-heptane; chlorobenzene + bromobenzene; ethyl bromide + ethyl iodide; n-butyl chloride + n-butyl bromide | Examples: Acetone +ethanol acetone +\[C{{S}_{2}}\]: water + methanol; water + ethanol; \[CC{{l}_{4}}+\]toluene; \[CC{{l}_{4}}+CHC{{l}_{3}}\]; acetone + benzene; \[CC{{l}_{4}}+C{{H}_{3}}OH\]; cyclohexane + ethanol | Examples: Acetone + aniline; acetone + chloroform; \[C{{H}_{3}}OH\ +\ C{{H}_{3}}COOH\]; \[{{H}_{2}}O\ +\ HN{{O}_{3}}\] chloroform + diethyl ether; water + HCl; acetic acid + pyridine; chloroform + benzene | |
| Solvent | Solute | Example |
| Gas | Gas | Mixture of gases, air. |
| Gas | Liquid | Water vapours in air, mist. |
| Gas | Solid | Sublimation of a solid into a gas, smoke. |
| Liquid | Gas | CO2 gas dissolved in water (aerated drinks). |
| Liquid | Liquid | Mixture of miscible liquids, e.g., alcohol in water. |
| Liquid | more...
"Solubility of a substance may be defined as the amount of solute dissolved in 100 gms of a solvent to form a saturated solution at a given temperature". A saturated solution is a solution which contains at a given temperature as much solute as it can hold in presence of dissolveding solvent. Any solution may contain less solute than would be necessary to saturate it. Such a solution is known as unsaturated solution. When the solution contains more solute than would be necessary to saturate it then it is termed as supersaturated solution.
In 1920, Latimer and Rodebush introduced the idea of "hydrogen bond".
For the formation of H-bonding the molecule should contain an atom of high electronegativity such as F, O or N bonded to hydrogen atom and the size of the electronegative atom should be quite small.
Types of hydrogen bonding
(1) Intermolecular hydrogen bond : Intermolecular hydrogen bond is formed between two different molecules of the same or different substances.
(i) Hydrogen bond between the molecules of hydrogen fluoride.
(ii) Hydrogen bond in alcohol and water molecules
(2) Intramolecular hydrogen bond (Chelation)
Intramolecular hydrogen bond is formed between the hydrogen atom and the highly electronegative atom (F, O or N) present in the same molecule. Intramolecular hydrogen bond results in the cyclisation of the molecules and prevents their association. Consequently, the effect of intramolecular hydrogen bond on the physical properties is negligible.
For example : Intramolecular hydrogen bonds are present in molecules such as o-nitrophenol, o-nitrobenzoic acid, etc.
The extent of both intramolecular and intermolecular hydrogen bonding depends on temperature.
Effects of hydrogen bonding
Hydrogen bond helps in explaining the abnormal physical properties in several cases. Some of the properties affected by H-bond are given below,
(1) Dissociation : In aqueous solution, hydrogen fluoride dissociates and gives the difluoride ion \[(HF_{2}^{-})\] instead of fluoride ion \[({{F}^{-}})\]. This is due to H-bonding in HF. This explains the existence of\[KH{{F}_{2}}\]. H-bond formed is usually longer than the covalent bond present in the molecule (e.g. in \[{{H}_{2}}O,\,\,O-H\] bond = 0.99 Å but H-bond = 1.77 Å).
(2) Association : The molecules of carboxylic acids exist as dimers because of the hydrogen bonding. The molecular masses of such compounds are found to be double than those calculated from their simple formulae. For example, molecular mass of acetic acid is found to be 120.
(3) High melting and boiling point : The compounds having hydrogen bonding show abnormally high melting and boiling points.
The high melting points and boiling points of the compounds \[({{H}_{2}}O,\ HF\] and \[N{{H}_{3}})\] containing hydrogen bonds is due to the fact that some extra energy is needed to break these bonds.
(4) Solubility : The compound which can form hydrogen bonds with the covalent molecules are soluble in such solvents. For example, lower alcohols are soluble in water because of the hydrogen bonding which can take place between water and alcohol molecules as shown below,
\[\begin{align} & \overset{\delta +}{\mathop{H}}\,-\overset{\delta -}{\mathop{O}}\,................\overset{\delta +}{\mathop{H}}\,-\overset{\delta -}{\mathop{O}}\,...............\overset{\delta +}{\mathop{H}}\,-\overset{\delta -}{\mathop{O}}\, \\ & \,\,\,\,\,\,\overset{\,}{\mathop{\,}}\,\,{{C}_{2}}{{H}_{5}}\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,H\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,\,{{C}_{2}}{{H}_{5}} \\ \end{align}\]
The intermolecular hydrogen bonding increases solubility of the compound in water while, the intramolecular hydrogen bonding decreases.
(5) As the compounds involving hydrogen bonding between different molecules (intermolecular hydrogen bonding) have higher boiling points, so they are less volatile. more...
Molecular orbital theory was given by Hund and Mulliken in 1932.
The main ideas of this theory are,
(1) When two atomic orbitals combine or overlap, they lose their identity and form new orbitals. The new orbitals thus formed are called molecular orbitals.
(2) Molecular orbitals are the energy states of a molecule in which the electrons of the molecule are filled just as atomic orbitals are the energy states of an atom in which the electrons of the atom are filled.
(3) In terms of probability distribution, a molecular orbital gives the electron probability distribution around a group of nuclei just as an atomic orbital gives the electron probability distribution around the single nucleus.
(4) Only those atomic orbitals can combine to form molecular orbitals which have comparable energies and proper orientation.
(5) The number of molecular orbitals formed is equal to the number of combining atomic orbitals.
(6) When two atomic orbitals combine, they form two new orbitals called bonding molecular orbital and antibonding molecular orbital.
(7) The bonding molecular orbital has lower energy and hence greater stability than the corresponding antibonding molecular orbital.
(8) The bonding molecular orbitals are represented by
 etc, whereas the corresponding antibonding molecular orbitals are represented by  etc.
(9) The shapes of the molecular orbitals formed depend upon the type of combining atomic orbitals.
(10) The filling of molecular orbitals in a molecule takes place in accordance with Aufbau principle, Pauli's exclusion principle and Hund's rule. The general order of increasing energy among the molecular orbitals formed by the elements of second period and hydrogen and their general electronic configurations are given below.
(11) Electrons are filled in the increasing energy of the MO which is in order
(a) \[\sigma 1s<{{\sigma }^{*}}1s<\sigma 2s<{{\sigma }^{*}}2s<\sigma 2{{p}_{z}}<\pi 2{{p}_{y}}\]
\[=\pi 2{{p}_{x}}<{{\pi }^{*}}2{{p}_{x}}={{\pi }^{*}}2{{p}_{y}}={{\pi }^{*}}2{{p}_{z}}\]
(b) \[=\frac{Increasing\text{ }energy\text{ }\left( for\text{ }electrons\text{ }>\text{ }14 \right)}{\sigma 1s<{{\sigma }^{*}}1s<\sigma 2s<{{\sigma }^{*}}2s<\pi 2{{p}_{x}}=\pi 2{{p}_{y}}<\sigma 2{{p}_{z}}<{{\pi }^{*}}2{{p}_{x}}}\]
\[={{\pi }^{*}}2{{p}_{y}}<{{\sigma }^{*}}2{{p}_{z}}\]
(12) \[\frac{Increasing\text{ }energy\text{ }(for\text{ }electrons\le 14)}{Number\text{ }of\text{ }bonds\text{ }between\text{ }two\text{ }atoms\text{ }is\text{ }called}\]
bond order and is given by
where  number of electrons in bonding MO.
 number of electrons in antibonding MO.
For a stable molecule/ion, 
(13) Bond order µ Stability of molecule µ Dissociation energy µ  .
(14) If all the electrons in a molecule are paired then the substance is a diamagnetic on the other hand if there are unpaired electrons in the molecule, then the substance is paramagnetic. More the number of unpaired electron in the molecule greater is the paramagnetism of the substance.
 more...
The basic concept of the theory was suggested by Sidgwick and Powell (1940). It provides useful idea for predicting shapes and geometries of molecules. The concept tells that, the arrangement of bonds around the central atom depends upon the repulsions operating between electron pairs(bonded or non bonded) around the central atom. Gillespie and Nyholm developed this concept as VSEPR theory.
The main postulates of VSEPR theory are
(1) For polyatomic molecules containing 3 or more atoms, one of the atoms is called the central atom to which other atoms are linked.
(2) The geometry of a molecule depends upon the total number of valence shell electron pairs (bonded or not bonded) present around the central atom and their repulsion due to relative sizes and shapes.
(3) If the central atom is surrounded by bond pairs only. It gives the symmetrical shape to the molecule.
(4) If the central atom is surrounded by lone pairs (lp) as well as bond pairs (bp) of  then the molecule has a distorted geometry.
(5) The relative order of repulsion between electron pairs is as follows : lp - lp > lp - bp > bp - bp.
A lone pair is concentrated around the central atom while a bond pair is pulled out between two bonded atoms. As such repulsion becomes greater when a lone pair is involved.
Geometry of Molecules/Ions having bond pair as well as lone pair of electrons
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