WBCHSE Class 11 Chemistry Notes For Intermolecular Forces – Definition and Types

States Of Matter

That matter exists in three physical states—solid, liquid and gaseous—is something you know from everyday experience. You must also be aware of the fact that a particular substance can exist as a solid, a liquid, or a gas, depending upon the temperature and pressure, example, water can exist in all three states.

In fact almost all substances can exist in the three states if the conditions of pressure and temperature are appropriate. Also, every pure substance can exist in the three states simultaneously (the three states are said to be in equilibrium) at a characteristic temperature and pressure called the triple point, about which you will learn later.

Generally speaking, we call a substance a solid if its melting point is above room temperature, a liquid if its melting point is below room temperature and boiling point is above room temperature, and a gas if its boiling point is below room temperature. This is under atmospheric pressure.

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• Certain other differences in the properties or behaviour of solids, liquids and gases are obvious from common observation. A solid, for example, has a definite shape and volume. It cannot be compressed easily and generally has a high density.
• A liquid has a definite volume but not a definite shape—it takes on the shape of the vessel it is poured into. It can be compressed more easily than a solid and generally has a lower density than a solid. A gas has neither a definite shape, nor a definite volume. It is highly compressible and has the least density among the three, i.e., solid, liquid and gas.
• Tire differences in these macroscopic or bulk properties of the three states of matter are due to the different arrangements of the constituent particles in the different states.
• In solids, the constituent particles are packed closely and are held together by a strong force of attraction between the particles. In liquids, the distance between the particles is greater and this force is weaker.
• In gases, the particles are far apart and this force is negligible. The particles in a solid are not really free to move. They possess only vibratory motion. The particles in a liquid enjoy greater freedom of movement. The particles constituting a gas have the greatest freedom of movement.

Intermolecular Forces

Intermolecular forces are responsible for the bulk properties of matter such as melting point and boiling point. The word bulk indicates a collection of atoms, molecules or ions. Intermolecular forces are those which exist between molecules as against intramolecular forces which hold the atoms together.

• In order to understand the difference between the two types of forces, let us take the example of water whose chemical formula is H₂O. There are two O—H bonds in the molecule and the energy required to break the two bonds is high; it is 930 kJ per mole of water.
• However, if we want to evaporate water, it requires only 41 kJ per mole of water. Generally, intermolecular forces are weaker than intramolecular forces.What are the forces that hold together the molecules in the different states of matter?
  What are the physical laws that govern the formation of a particular state of matter?
  Let us now study the types of forces existing between the particles of a matter that make it acquire a particular state.
  The particles may be molecules, atoms or ions.

van der Waals forces: Dipole-dipole forces, dipole-induced dipole and dispersion forces make up van der Waals forces, called so after the Dutch physicist Johannes van der Waals. These occur between molecules (intermolecular forces) or atoms (interatomic forces).

Dipole-dipole forces Neutral but polar molecules experience dipole forces as a result of electrical interactions among dipoles on neighbouring molecules. There is partial positive (δ+) or partial negative (δ-) charge associated with the two ends of the dipoles in a molecule.

• The molecules in a sample could be oriented in various ways and so the forces could be attractive (when unlike charges are close together) or repulsive (when like charges are close together) between the dipoles of neighbouring molecules.
• The overall effect is the average of all such forces. The potential energy is lower for the attractive orientation and so this predominates. The net potential energy will, therefore, be an average of the different orientations.
• It increases as the distance between the interacting dipoles decreases. The potential energy of interaction is proportional to 1/r⁶ where r is the distance between the dipoles (molecules). It also varies inversely as a function of temperature.

The strength of the given dipole-dipole interaction depends on the sizes of the dipole moments involved. The more polar the substance, the greater is the strength of its dipole-dipole interaction.

This is reflected in the boiling point values for different substances. Substances with higher dipole moments tend to have higher boiling points.

Dipole-induced dipole forces Let us first understand what an induced dipole is. A molecule having a permanent dipole moment (polar molecule) can induce a dipole in another molecule by deforming its electronic cloud. The latter molecule, thus, gets an induced dipole.

• This induced dipole interacts with the dipole of the other molecule and the forces involved are called dipole-induced dipole forces. The strength of the interaction depends on the dipole moment of the polar molecule and the polarisability (the ease with which the electron cloud can get deformed) of the second molecule.
• The potential energy of interaction is directly proportional to these two factors. It also varies as 1/r⁶ as in the case of dipole-dipole forces.

London dispersion forces Intermolecular forces not only exist between charged and polar particles but also among nonpolar molecules. These are called London dispersion forces after Fritz London, who first explained them. In fact, all atoms and molecules experience London dispersion forces which result from the motion of electrons.

• Let us consider the example of a nonpolar molecule Br₂. It has a zero dipole moment averaged over time but at any point of time, the electrons may be more towards one end than the other resulting in a short-lived dipole. This dipole may induce temporary dipoles in neighbouring molecules resulting in attractive forces between the molecule with the permanent dipole and the neighbouring molecule.
• These forces are weak and the energy of interaction is small, in the range of 1-10 kJ mol^-1. The exact magnitude of the energy depends on the polarisability of the molecule. Smaller atoms or molecules are less polarisable as compared to larger ones and the dispersion forces are also loss. This can be easily understood by taking halogens as examples. F₂ and Cl₂ are gases at room temperature, Br₂ is a liquid and I₂ is a solid.

The intermolecular force of attraction increases in the order: F₂ <Cl₂; <Br₂, <I₂.

The increase in force of attraction with the increase in the size of atoms involved can be explained by London’s theory, In a larger atom, the valence electrons are, on an average, farther from the nuclei than those in a smaller atom.

• Also, they are loss tightly held and can more easily form the temporary dipoles that cause the attraction between molecules. Dispersion forces usually increase with molar mass because such molecules have more electrons and dispersion forces increase in strength with the number of electrons. Another factor which is important in determining the magnitude of the dispersion forces is its shape.
• Shapes which have larger molecular surface area allow greater contact between molecules and give rise to higher dispersion forces than do compact molecules. For instance, the boiling point of u-pentane is 309.4 K while that of neopentane or 2,2-dimethyl propane is 282.7 K though both of them have the same molecular formula.
• Tire reason behind this behaviour is that in n-pentane there is a possibility of greater contact between molecules because of its spread-out structure. Hence, this leads to stronger intermolecular dispersion forces and, therefore, a higher boiling point.
• The potential energy associated with London dispersion forces is again inversely proportional to the sixth power of the distance between tire molecules.

Hydrogen bonding: As already explained in Chapter 4, a hydrogen bond is formed whenever a hydrogen atom is bonded to one or more electronegative atoms (fluorine, oxygen or nitrogen).

The large electronegativity difference between the atoms leads to a highly polar covalent bond with a partial positive charge on the hydrogen atom and a partial negative charge obviously on the electronegative atom.

• Tire electrostatic attraction between the partially positive H atom in one molecule and the partially negative atom in another molecule gives rise to a strong dipole-dipole attraction—the hydrogen bond. Hydrogen bonds are weaker than ordinary covalent bonds but stronger than other dipole-dipole attractions and London dispersion forces.
• Hydrogen bonds have a pronounced effect on the properties of condensed phases—solid and liquids.

Generally, the energy required to overcome the intramolecular forces decreases with a decrease in molar mass. The amount of heat required to evaporate 1 mole of the hydrides in a particular group generally decreases with decrease In molar mass but H₂O, HF and NH₃ exhibit exceptionally high values.

This is because more energy is required to overcome H-bonding in these three cases.

Thermal Energy: Thermal energy (i.e., heat energy) is the kinetic energy associated with the random motion of atoms and molecules. The greater the amount of thermal energy in a sample, the more vigorous is the motion of atoms and molecules. This is called thermal motion. Thermal energy is directly proportional to temperature.

• If we increase the amount of thermal energy in a sample of matter, then either the temperature of the matter increases or the substance melts or evaporates with no change in temperature. If we decrease the thermal energy, then either the temperature decreases or the substance condenses or freezes with no change in temperature. The change in thermal energy of a substance can be observed by observing the change in temperature or state of that substance.
• Thus, the state of a matter depends on its intermolecular forces and thermal energy. An increase in the thermal energy of a solid or liquid helps to overcome the attractive forces between the particles and the substance melts (or evaporates). Similarly, if the thermal energy is decreased the intermolecular forces play a greater role than thermal motion and a liquid (or gas) is transformed to a solid (or liquid). Now we will learn about the laws which govern the behaviour of matter in different states and thus know more about the states of matter.

Measurement Of Properties Of Gases

Let us begin our study of the states of matter with gases and liquids. The behaviour of gases is studied in terms of their mass, volume, temperature and pressure and the relationship between these properties. So, we must know how these properties are measured before we examine the behaviour of gases.

Mass: The mass of a gas can be determined by weighing a container filled with the gas and then subtracting the weight of the empty container from this weight. The number of moles or the amount of the gas can be determined from the following equation.

n = m/M,

where n = number of moles,

m = mass of gas, and

M = molar mass of gas.

The number of molecules present in a particular mass of a gas can be determined by multiplying the number of moles (in that mass of the gas) by the Avogadro constant, 6.022 x 10²³ mol^-1.

Volume: The volume of a substance is the space occupied by it. A gas occupies the entire space available to it, so the volume of a gas is the same as the volume of the container it occupies. The SI unit of volume is m³, but usually volume is expressed in litres (L) or millilitres (mL).

1 m³ = 10³ dm³ = 10⁶ cm³

1 mL = 1 cm³

1 L = 1000 cm³ = 1 dm³

Pressure: Pressure is force per unit area. A gas confined in a container exerts uniform pressure on the walls of the container in all directions. The pressure of the gas is the outward force exerted by the gas per unit area on the walls of the container. The pressure exerted by the gases of the atmosphere on the surface of the earth is known as atmospheric pressure.

Barometer A (mercury) barometer is a simple device used to measure atmospheric pressure. It consists of a long glass tube (longer than 76 cm) filled with mercury, which is closed at one end and inverted (the open end) into an open vessel containing mercury.

The level of mercury in the tube drops until the downward pressure exerted by the column of mercury (in the tube) on the mercury in the vessel is exactly balanced by the atmospheric pressure, which acts on the open surface of the mercury in the vessel. Normally, the height of the mercury in the tube is 76 cm. It increases or decreases as the atmospheric pressure increases or decreases.

Atmospheric pressure The atmospheric pressure can be calculated from the pressure exerted by the column of mercury. Let the height of the column of mercury be h cm and its area of cross section be A cm². Then the pressure exerted by the column of mercury can be calculated as follows.

Pressure (p) = \(\frac{\text { force }}{\text { area }}=\frac{\text { mass } \times \text { acceleration }}{\text { cross-sectional area }}\)

= \(\frac{m \times g}{A}\)

where m = mass of mercury column and

g = acceleration due to gravity.

Let V be the volume of mercury in the tube and d its density.

Then p = \(\frac{V \times d \times g}{A}=\frac{A \times h \times d \times g}{A}\) = hdg (V =A x h)

Knowing the density of mercury (1.35951 x 10⁴ kg m^-3 at 0°C) and the acceleration due to gravity (9.8065 m s^-2), it is possible to calculate the pressure exerted by the column of mercury 76 cm in height.

p = \((0.760 \mathrm{~m})\left(1.35951 \times 10^4 \frac{\mathrm{kg}}{\mathrm{m}^3}\right)\left(9.8065 \frac{\mathrm{m}}{\mathrm{s}^2}\right)=101,325 \mathrm{~kg} \mathrm{~m}^{-1} \mathrm{~s}^{-2}\) = 101, 325 N m^-2

Thus the measurement of the atmospheric pressure does not depend upon the area of cross section of the glass tube of the barometer.

The SI unit of pressure is pascal (Pa), which is defined as the pressure exerted when a force of 1 N acts on an area of 1 m².

1 Pa = 1 N m².

As calculated above, 1 atm = 101,325 N m^-2 = 101,325 Pa = 101.325 kPa.

The unit pascal is too small for gases and generally we express the pressure of the gas in terms of bar. The standard atmospheric pressure is 1 bar.

1 bar = 10⁵ Pa = 100 kPa.

Pascal and bar are related to the older unit of measurement, atm, as

1 atm = 1.01325 bar = 1.01325 x 10² kPa

or 1 bar = 0.987 atm.

One atmosphere may approximately be taken to be equal to 10² kPa or 10⁵ Pa.

1 atm is also referred to as 760 torr. The ‘torr’ is still used as a unit of pressure in some laboratories.

Manometer A manometer is used to measure the pressure of a gas. It consists of a U-tube partially filled with mercury. One of the limbs of the tube is longer than the other. The longer limb is open, while the shorter limb with its horizontal arm is connected to a closed vessel containing the gas whose pressure is to be measured.

The mercury in the open limb is exposed to atmospheric pressure, while the mercury in the shorter tube is subjected to the pressure of the gas. The difference in the level of the mercury in the two limbs can be used to determine the pressure of the gas.

1. If the level of mercury is equal in both limbs, the pressure of the gas is equal to the atmospheric pressure.
2. If the level is higher in the open tube, the pressure of the gas is higher than the atmospheric pressure (that is why the mercury gets pushed into the longer limb). Then the pressure of the gas can be determined as follows. Gas pressure (p_g) = atmospheric pressure (p_atm) + difference in mercury levels in two columns (h)
3. If the level of mercury is higher in the shorter tube, it means the pressure exerted by the gas is less than the atmospheric pressure. In this case, the pressure of the gas is determined as follows.

Gas pressure (p_g) = atmospheric pressure (p_atm) – difference in mercury levels in two columns (h)

The kind of manometer we have just discussed is called an open-end manometer. In some manometers, the longer limb is closed and the space above mercury on the closed end-side is completely evacuated. They are called closed-end manometers and are used only for gases at pressures less than the atmospheric pressure.

Temperature: The temperature of a gas, like that of other substances, is measured with the help of a thermometer. A thermometer, as you must already know, uses the fact that most substances expand on heating. The substance most commonly used in thermometers is mercury. Three scales are used for the measurement of temperature.

Celsius scale The Celsius (centigrade) scale, named after the Swedish astronomer, Anders Celsius, takes the freezing point of water as 0°C and the boiling point (both points at normal atmospheric pressure) of water as 100°C. The scale is divided into 100 equal parts. Meteorologists use this scale.

Fahrenheit scale The Fahrenheit scale, named after the German physicist, Daniel Fahrenheit, takes the freezing point of water as 32°F and the boiling point of water as 212°F. The scale is divided into 180 equal parts. Clinical thermometers use this scale.

Kelvin scale The Kelvin scale, named after the British physicist, Lord Kelvin, was the result of the study of gases. You will read more about this scale later in this chapter, kelvin is the SI unit of temperature and degrees are not used while expressing temperatures on this scale. For instance, 0°C – 273.15 K, not 273.15°K. The size of the degree on a Kelvin scale is the same as on the Celsius scale. Thus, on the Kelvin scale, water freezes at 273.15 K and boils at 373.15 K at 1 bar pressure.

We can measure the amount, pressure, volume and temperature of a gas in a sealed container. The observed relationships between these four interdependent variables are expressed in the form of gas laws. These laws are the quantitative relationships between any two of these variables when the other two are kept constant.

States Of Matter Multiple Choice Questions

Question 1. At the same temperature and pressure, which of the following gases will have the highest kinetic energy per mole?

1. Hydrogen
2. Oxygen
3. Methane
4. Same in all cases

Answer: 4. Same in all cases

Question 2. The mass of 5.6 L of a gas at stp is 11 g. The gas may be

1. PH₃
2. CoCl₂
3. NO
4. N₂O

Answer: 4. N₂O

Question 3. Equal weights of ethane and hydrogen are mixed in an empty container at 25°C. The fraction of the total pressure exerted by hydrogen is

1. 1/4
2. 1/2
3. 15/16
4. 1/16

Answer: 4. 1/16

Question 4. If you take 5.0 g of each of the following gases at 87°C and 750 mm Hg pressure, which will have the least volume?

1. HF
2. HCl
3. HBr
4. HI

Answer: 2. HCl

Question 5. The density of neon will be the highest at

1. stp
2. 0°C, 2 atm
3. 273°C, 1 atm
4. 273°C, 2 atm

Answer: 2. 0°C, 2 atm

Question 6. The temperature below which a gas does not obey the ideal gas law is called the

1. Critical temperature
2. Inversion temperature
3. Boyle temperature
4. Reduced temperature

Answer: 3. Boyle temperature

Question 7. 1.0 g of H₂ is heated from 300 K to 600 K in a dosed flask of volume 5 L. Which of the following statements is not correct?

1. The tire pressure of the gas increases.
2. The rate of collisions increases.
3. The number of moles increases.
4. The energy of the gaseous molecules increases.

Answer: 3. The number of moles increases.

Question 8. At constant temperature, for a given mass of an ideal gas,

1. The ratio of pressure and volume remains constant
2. The pressure remains constant
3. The volume remains constant
4. The product of pressure and volume remains constant

Answer: 4. The product of pressure and volume remains constant

Question 9. The compressibility factor for an ideal gas is

1. 1.5
2. 1.0
3. 2.0
4. ∞

Answer: 3. 2.0

Question 10. The molecular weights of O₂ and SO₂ are 32 and 64 respectively. If 1 L of O₂ at 15°C and 750 mm Hg contains N molecules, the number of molecules in 2 L of SO₂ under the same conditions of temperature and pressure will be

1. 2N
2. N
3. 2N
4. 4N

Answer: 3. 2N

Question 11. Two separate bulbs contain gas A and gas B. The density of gas A is twice that of B. The molecular mass of A is 12 that of B. If the temperature is constant, the ratio of the pressures of A and B is

1. 1:1
2. 1: 2
3. 4 :1
4. 2:1

Answer: 3. 4 :1