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FSc Notes Chemistry Part 1 Chapter 4 Liquids and Solids Lecture 1

FSc Notes Chemistry Part 1 Chapter 4 Liquids and Solids Lecture 1


A substance is said to be in liquid stat; if it has a fixed volume but no fixed shape. liquids always acquire the shape of container. According to kinetic molecular theory, there are some sort of intermolecular attractive forces which do not let the molecules to have a random motion. thus the molecules of liquids can move inside the liquid with a rotator motion but can not go out of the liquid. that is why the volume of a liquid remains constant however its. shape is always like the shape of its container.

Intra-Molecules attractive forces:

the forces of attraction between the atoms of a molecule are known as intra-molecular attractive forces.
For example the covalent bonds between H and o atoms in H2O Molecule. The forces of attraction which keep the particles (atoms or molecules) of a substance together, are known as inter molecular attractive forces.
For example hydrogen bonding in water. intermolecular attractive forces are not true chemical, bonds they are just a temporary sort of forces to keep the particles of a substance together. thus we can say that intermolecular attractive forces keep a substance in a particular physical state. intermolecular attractive forces are found in all kinds of atoms and molecules when they are sufficiently close to each other. " such intermolecular attractive force which are found in all kinds of atoms and molecules when they are sufficiently close to each other are known as " Vander-wall's force"
from the above discussion. it is clear that inter molecular attractive forces bring the molecules of a substance together and particular physical properties to the substance in gaseous liquid and solid.

Types of Intra-Molecules attractive force:
following are the important types of intermolecular attractive forces.
  1. Dipole - dipole interactions
  2. Hydrogen bonding
  3. London dispersion forces
  4. ion-dipole interactions
  5. Dipole - induced dipole interactions.

(1) Dipole - Dipole Interactions:

Any polar molecule (i.e. molecule having positive and negative poles) is known as a dipole. "When ever the positive pole of one dipole attracts the negative pole of another dipole, this force of attraction is known as dipole - dipole interaction". e.g. the HCl molecules are polar in nature because of having a more electro negative " cl" and less electronegative "H". cl has a partial +ve (+s) change. the +ve pole of one HCl molecule attracts the -ve pole of another HCl molecule and in such a way all the HCl molecules get attracted with each other by means of dipole - dipole interactions. The dipole - dipole force are temporary forces of attraction i.e. they are not true chemical bonds. Dipole - Dipole forces are approximately one percent as effective as a covalent bond. Factors Affecting Dipole - Dipole interactions: 
Some important factors upon which the strength of dipole -dipole interactions depend are:

(a) Electron affinity difference b/w:

The bonded Atoms: Greater the electrongeativity difference b/w the covalently bond atoms of a polar molecular (dipole) greater will be the polarity and hence stronger will be the dipole - dipole interactions and vice versa.
e.g: the dipole - dipole interactions b/w HCl molecules are are more stronger than the dipole - dipole interactions b/w HBr molecules. it is because of the fact that cl of HCl is more electronegative then Br of HBr. thus HCl has more polarity than HBr.

(b) Distance Between the Molecules:

The distance between the molecules of a substance weaker will be the dipole -dipole interactions between them and vice versa.
e.g: in case of gases the distance between the molecules is very large thus gases have very weak Dipole - Dipole forces on the other hand the distance between the molecules of a liquid is smaller thus then are stronger dipole -Dipole interactions. some important substances having dipole -dipole interactions between their molecules are H2S, HI, HBr, PH3, etc. it is important to note that great are the strength of dipole- dipole interaction of a substance, greater will be the rules of its differ thermodynamic parameters like m.p; b.p; heat of vaporization head of sublimation etc.

(2) Hydrogen Bonding:

it is a special type of intermolecular attractive forces. Hydrogen bonding is actually Dipole - Dipole interactions but due to the presence of most electronegative elements (i.e. F, O & N) and the least electronegative element "H", it has got extra strength, therefore has been given a special name "Hydrogen bonding"
Hydrogen bonding can be defined as " The force of attraction between the partial positive Hydrogen of one dipole and partial negative N, O or F of another dipole." All these substances which have O-H bond, N-H bond or F-H bond in their formula have hydrogen bonding among their molecules.
For example :
  1. The chemical formula of water is H2O which contains O-H bond. Thus there is hydrogen bonding among water molecules. it is clear as follow,
  2. Chemical formula of ammonia is NH3 which contains N-H bonds. Thus there is hydrogen bonding among NH3 molecules
  3. The chemical formula of hydrofluoric acid is HF which contains H-F bond and thus there is hydrogen bonding among its molecules. There are some substances like proteins, Ammines etc which have hydrogen bonding.

Evidence For Hydrogen bonding:

As it is clear that hydrogen bonding is a special type of intermolecular attractive forces and is more stronger than dipole -dipole interactions. To prove this fact let us consider the hydrides of element of Group VA, VIA, & VIIA. Hydrides are the binary compounds (in which there are only two different elements) of hydrogen. but before explaining this fact let us understand the trend of boiling points of hydrides of Group IVA elements i.e. CH4, SiH4, GeH4, SnH4 and pbH4. As we go down the group in GP IV A the b.p of their hydrides of elements of this group increases.
It is because of the fact that down the group, the size of molecules (hydrides) increases due to which their valence electrons get away from their nuclei and thus Vander waal's forces of attraction can easily an more strongly can operate. thus due to presence of stronger Vander waal's forces, more heat is required to separate the molecules and that is why the b.p; of the hydrides of Group IVA elements increases.

Now it is clear that the hydrides of elements of Group VA, VIA and VIIA should also show, a similar behavior (trend) of their boiling points. when it is practically noted, it was seen that of Group VA, VIA & VIIA showed the same behavior of their b.p; which proves that due to increase in their size down the group; Vander waal's forces become more and more stronger and thus down the group the b.p of hydrides of Group VA, VIA and VIIA increases.
But it was surprising to see that the hydrides of first elements of Group VA, VIA & VII A show very different behavior.
According to the routine their boiling points: should have been lower than all the other members according to the above explanation but their point are surprisingly very much higher than all other members. this proves that there is some special force of attraction between these hydrides, which named as hydrogen bonding.

The Origin of Hydrogen Bonding:

Hydrogen bonding is found between all those molecules in which hydrogen atom is directly bonded by means of a covalent bond to the most electromotive element of the periodic table (i.e mF, O, Or N).
Due to the presence of any of the above mentioned elements i.e. F, O & N bonded with hydrogen atom, the bond is much more polar between the shared pair is attracted more by the most electronegative element (F,O or N) and thus hydrogen atom gets a significant positive charge. The best examples of the molecules which have hydrogen bonding are H20, HF and NH3. in all the above mentioned molecules, there is a most polar H-O, H-F or H-N bond which is responsible for hydrogen bonding.
It is interesting to see that the most E.N element of H2o, HF and NH3 i.e. O,F & N has at least one " active" lone pair of electrons. The word active line pair means that this lone pair also takes part in hydrogen bonding.
The generally the no of active lone pairs of electrons in a molecule, greater will be the no of hydrogen bonds established with that molecule and hence the b.p of the substance will thus be higher, and vice versa.
now for example, in H2O molecule, oxygen atom has two active lone pairs of electrons both of which take part in hydrogen bonding. There are also two positively charged H atom in H2O molecule which also form hydrogen bonds with other water molecules. thus each water molecule has four sites of attachment to form 4 hydrogen bonds with 4 other H2O molecules.

Hydrogen Bonding in Alcohols:

Any organic molecule in which a –OH group (hydroxyl group) is directly attached with a carbon, atom by means of a covalent bond is known as an alcohol.
As there is O-H bond in every alcohol molecule, Thus there will be hydrogen bonding in alcohols. It can be proved by the following example.
Let us consider ethanol (an alcohol) and dim ethyl ether or methy oxy methane (an ether) i.e. Both these molecules have same chemical formula i.e. C2H6O. Thus both have same size and same no of electrons. So we can clearly say that both of them have same vander wall: attractions (i.e. dispersion force & dipole –dipole interactions. But their b.p shows that the molecules of ethanol are sticked together much firmly than the molecules of methoxy methane. If proves the presence of hydrogen bonding in alcohols (ethanol).

Hydrogen Bonding Exists in Addition to Vander Wall's

It is important to note that hydrogen bonding exists in addition to Vander wall’s attractions. (i.e. dispersion forces or dipole – dipole interactions). Hydrogen bonding is established only between the H-O, H-P or H-N bonds of the molecules. Other parts of the molecules are attached together by means of Vander wall’s force of attractions (i.e. dispersion forces or dipole –dipole interactions). The molecules which are slim or narrow and longer have stronger dispersion forces than wider and shorter molecules.

Now to prove that hydrogen bonding exists in addition to Vander wall’s forces, consider the following examples.
(Pentane) CH3-CH2-CH2-CH2-CH3 b.p 36.3 degree Celsius
(1-butanol) CH3-CH2-CH2-CH2-OH b.p 117 degree Celsius
(2- Methyl – 1- Propanol) CH3-CH-CH2-OH-CH3 b.p 108 degree Celsius

All the above mentioned molecules have same number of electrons. The first two molecules i.e. pentane and 1-butanol have same length but due to hydrogen bonding in 1-butanol, its b.p. is much more than pentane.
On the other hand the two last molecules i.e. 1-butanol and 2-methyle – 1 Propanol have same chemical formula (i.e. C4H9OH) and both have O0H bond so both have hydrogen bonding of same strength. But still the b.p. Of 1-butanol is higher (117oc) than 3 – methyl -1- Propanol (108oc). it is because of the fact that the molecule of 1-butanol is thinner and longer than 2- methyl -1-propanol. So there are stronger dispersion forces in 1-butanol than in 2-methyle -1- Propanol.
Thus it is proved that hydrogen bonding exists in addition to Vander Waal s forces of attractions.
If there were no Vander Waal s forces dispersion forces) then there would have no difference in their b.p both have same formula (C4H9OH) and same strength of hydrogen bonding both have an O-H bond.

Application of Hydrogen Bonding:

Hydrogen bonding has so many applications. Some important are as following.
  1. As ‘F’ is more electronegative then “O” therefore the H_F bond of HF is more polar than H-O bond of H2O and thus the hydrogen bonding among HF molecules is expected to be more stronger than among (H2O) molecules and thus b.p. of HF should be higher than H2O. But actually the b.p of H2O is much more than the b.p. Of HF. It is because of the fact that in water, each H2O molecule is attached with four other bonding on the other hand each HF molecule is attached with only two other HF molecules by means of Hydrogen bonding. That is why b.p of water is higher than b.p. of HF (hydro floric acid).
  2. Among the hydrides of GP VII A elements (i.e. HF, HCL, HBr), HF has low acidic strength (i.e. weak acid as compared to HCl, HBr, and HI. It is because of the fact that there is hydrogen bonding among the molecules of hydrofluoric acid (HF) and thus these molecules cannot be separated easily while in case of HCl, HBr, and HI, there are dipole-dipole interactions among their molecules and thus there molecules can easily be separated. Thus HF is weak acid as compared to HCl, HBr, & HI.
  3. Any substance which has hydrogen bonding among its molecules is soluble in water of the formation of hydrogen bonding among the molecules of water and molecules of the dissolving substance. For example as ethanol (C2H5OH) has O-H bond, thus there is hydrogen bonding among ethanol molecules. Thus ethanol is water soluble b/c bonding among the H2O and C2H5-OH. i.e. Similarly carboxylic acids (they are organic molecules containing COOH group) are also soluble in water b/c of hydrogen bonding.
  4. Soaps and detergents perform the cleansing action because of the hydrogen bonding among the molecules of soaps or detergents and H2O molecules. Actually soap or detergents molecules have two parts. One is polar and the other is non-polar. The polar part gets attached with water molecules by means of hydrogen bonding while the non-polar part get attached with the dirt and thus the cleansing action is performed.
  5. Many organic molecules like proteins, DNA, fibers of hair, silk and muscles etc and paints dyes, glue and food materials like honey etc have hydrogen bonding because they either have NH or O-H bond, which are responsible for hydrogen bonding in them.

Written by: Asad Hussain