The behaviour of matter in different states is governed by various physical laws. According to you, what are the factors that determine the state of matter?
Temperature, pressure, mass and volume are the factors which determine the different states of matter. i.e., solid, liquid and gas.
Use the information and data given below to answer the questions (a) to (c),
Stronger intermolecular forces result in higher boiling point.
Strength of London forces increases with the number of electrons in the molecule.
Boiling point of $\mathrm{HF}, \mathrm{HCl}, \mathrm{HBr}$ and HI are $293 \mathrm{~K}, 189 \mathrm{~K}, 206 \mathrm{~K}$ and 238 K respectively.
(a) Which type of intermolecular forces are present in the molecules HF, $\mathrm{HCl}, \mathrm{HBr}$ and HI ?
(b) Looking at the trend of boiling points of $\mathrm{HCl}, \mathrm{HBr}$ and HI , explain out of dipole-dipole interaction and London interaction, which one is predominant here.
(c) Why is boiling point of hydrogen fluoride highest while that of hydrogen chloride lowest?
From the information and data given in the question, we concluded that
(a) In $\mathrm{HCl}, \mathrm{HBr}$ and HI , dipole-dipole and London forces are present because molecules possess permanent dipole. In HF dipole-dipole, London forces and hydrogen bonding are present.
(b) Electronegativity of chlorine, bromine and iodine decreases in the order are present
$$\mathrm{Cl}>\mathrm{Br}>\mathrm{I}$$
Therefore, dipole moment should decrease from HCl to HI Thus, dipole-dipole interaction should decrease from HCI to HI But boiling point increases on moving from HCl to HI This means that London forces are predominant.
This is so because London forces increases as the number of electrons in a molecule increases and in this case number of electrons is increasing from HCl towards HI
(c) Hydrogen fluoride has highest dipole moment attributes due to highest electronegativity of fluorine as well as presence of hydrogen bonding in HF. Therefore, HF has highest boiling point.
What will be the molar volume of nitrogen and argon at 273.15 K and 1 atm?
When temperature and pressure of a gas is 273.15 K (or $0{ }^{\circ} \mathrm{C}$ ) and 1 atm (or 1 bar or $10^5$ pascal), such conditions are called standard temperature and pressure conditions (STP). Under these conditions, the volume occupied by 1 mole of each and every gas is 22.4 L . Hence, the moler volume of $\mathrm{N}_2$ and Ar at 273.15 K and 1 atm is 22.4 L .
A gas that follows Boyle's law, Charle's law and Avogadro's law is called an ideal gas. Under what conditions a real gas would behave ideally?
At low pressure and high temperature, a real gas behaves as an ideal gas. Almost all gases are real gas.
Two different gases ' $A$ ' and ' $B$ ' are filled in separate containers of equal capacity under the same conditions of temperature and pressure. On increasing the pressure slightly the gas ' $A$ ' liquefies but gas $B$ does not liquify even on applying high pressure until it is cooled. Explain this phenomenon.
The temperature above which a gas cannot be liquefied howsoever high pressure may be applied on the gas is called critical temperature. Since, gas 'A' liquifies easily, this suggests gas ' $A$ ' is below its critical temperature. On the other hand, gas 'B' does not liquefy easily even on applying high pressure. This suggests that gas ' $B$ ' is above its chitical temperature.