Rate of diffusion \(\text { (r) } \alpha \frac{1}{\sqrt{M}}\)
\(\mathrm{M}_{\mathrm{NH}_{3}}=17 ; \mathrm{M}_{\mathrm{HCl}}=36.5 \)
\(\therefore \mathrm{r}_{\mathrm{NH}_{3}}>\mathrm{r}_{\mathrm{HCl}} \)
Hence white fumes are first formed near HCL.

On heating incineration might take place due to the increase of pressure

If there is a tiny sticky area on the walls of the container of a gas, then the collision of the molecules will not be elastic. So the gas will behave as a real gas. So the pressure will be less than the calculated value.

The following table gives the pressure of a gas with increasing temperature at constant volume for two cylinders.

Both the models explain Gay - Lussac's Law. P \(\propto\) T are constant V & n
For 50 L container:
The pressure increases with rise in temperature (at constant volume)
For 75 L container:
The pressure increases with rise in temperature (at constant volume)

The standard atmospheric pressure (1atm.) is the pressure that supports a column of mercury exactly 760 mm (or 76 cm) high at 0°C at sea level. The standard atmosphere is thus equal to a pressure of 760 mm Hg, in which mm Hg represents the pressure exerted by a column of mercury 1 mm high. Thus 1 atm. = 760 mm Hg

If the molecular collisions were inelastic, there would have been loss of energy, which ultimately slows down the molecules.
Finally the pressure would gradually reduce to zero.

Units for Vanderwaal's constant:
The dimensions of the vanderwaal's constant a and b depend upon the units of P and V respectively.
\(a=\frac { Pressure.({ Volume) }^{ 2 } }{ { mole }^{ 2 } } \)
a = atm . dm⁶ . mol^-2 (1 litre = 1 dm³)
If volume is expressed in dm³. Then b is expressed as,
\(\frac { Vol }{ n } =\frac { { dm }^{ 3 } }{ mol } \)
\(\frac { V }{ n } ={ dm }^{ 3 }{ mol }^{ -1 }\)
unit of b = dm³ mol^-1

Based on the values of the critical temperature, the gases are classified into permanent and temporary gases.
(i) Gases having very low critical temperature belong to permanent type.
Eg: H₂, N₂, He, etc.
(ii) Gases having critical temperature in the ordinary range of temperatures belong to the temporary type.
Eg: NH₂, CO₂, SO₂, HCi, etc.

It is defined as the force exerted by a gas on unit area of the wall
\(Pressure=\frac { Force }{ Area } \)
The SI unit of pressure is the Pascal (Pa), named after the French Mathematician and Physicist Blaise Pascal, and
1 Pa = 1 N m^-2
1 N = 1 Kg m s^-2

(i) Isotherm:
The pressure-volume curve at constant temperature are known as isotherm.
(ii) Isobar:
The volume - temperature curve at constant pressure are known as isobar.
(iii) Isochore:
The pressure-temperature curve at constant volume are known as isochore.

The pressure decreases with the increase in altitude because there are fewer molecules per unit volume of air. Above 9200 m (30000 ft.) the pressure is so low that one could pass out for lack of oxygen. For this reason most airplanes cabins are artificially pressurized.

For ideal gases, a = 0. The value of 'a' is a measure of the attractive forces between the molecules. There will not be any intermolecular forces of attraction. So it cannot be liquefied.

According to Avogadro's hypothesis: V \(\alpha\) n
\(\therefore \frac{V_{1}}{n_{1}}=\frac{V_{2}}{n_{2}}=\text { constant }\)
V₁ & n₂ are the volume and number of moles of a gas and V₂ & n₂ are a different set of values of volume and number of moles of the same gas at same temperature and pressure.

Astronauts must wear space suits filled with air, whenever they leave a space craft and are exposed to the environment of space.
Dangers experienced on moon's space :
1. In space there is no air to breathe and no air pressure.
2. Space is extremely cold and filled with dangerous radiations.
3. If they do not wear space suits, there may be bleeding due to high body pressure.
4. The space suits prevent astronauts from impacts of small bits of space dust.
5. There is no atmosphere on the moon and there are dangers from micro-meteorite impacts.
6. The astronauts are protected from these dangers on wearing space suits.