If the pressure of a given mass of gas is reduced to half and temperature is doubled simultaneously, the volume will be

The critical volume of a gas is $0.072 \mathrm{lit}. {\mathrm{mol}}^{–1}$. The radius of the molecule will be , in cm

32 gm of oxygen and 3 gm of hydrogen are mixed and kept in a vessel to 760 mm pressure and 0ºC. The total volume occupied by the mixture will be nearly

A closed vessel contains equal number of nitrogen and oxygen molecules at pressure of P mm. If nitrogen is removed from the system, then the pressure will be

Two vessels of capacities 3 litres and 4 litres are separately filled with a gas. The pressures are respectively 202 kPa and 101 kPa. The two vessels are connected. The gas pressure will be now, at constant temperature.

Two flasks A and B of equal volume containing ${\mathrm{NH}}_3$ and HCl gases, are connected by a narrow tube of negligible volume. The two gases were prevented from mixing by stopper fitted in connecting tube. For further detail of experiment refer to the given figure. What will be final pressure in each flask when passage connecting two tubes are opened. Assume ideal gas behaviour of ${\mathrm{NH}}_3$ and $\mathrm{HCl}$ gas and the reaction.

The ratio of average molecular kinetic energy of ${\mathrm{UF}}_6$ to that of ${\mathrm{H}}_2$, both at 300 K is$-$

A mono atomic gas diatomic gas and triatomic gas are mixed, taking one mole of each ${\mathrm{C}}_{\mathrm{p}}/{\mathrm{C}}_{\mathrm{v}}$ for the mixture is$-$

According to kinetic theory of gases, for a diatomic molecule

The values of vander waals constant ‘a’ for the gases ${\mathrm{O}}_2$, ${\mathrm{N}}_2$, ${\mathrm{NH}}_3$ and ${\mathrm{CH}}_4$ are 1.36, 1.39, 4.17 and 2.253 ${\mathrm{lit}}^2$ atom mol^-2 respectively. The gas which can most easily be liquefied is$-$

At low pressure, the vander waals equation is written as :

X ml of ${\mathrm{H}}_2$ gas effuses through a hole in a container in 5 seconds. The time taken for the effusion of the same volume of the gas specified below under ideal condition is

The valves X and Y are opened simultaneously. The white fumes of ${\mathrm{NH}}_4\mathrm{Cl}$ will first form at :

A gas mixture consists of 2 moles of oxygen and 4 moles of a argon at temperature T. Neglecting all vibrational modes, the total internal energy of the system is –

Given reaction : $\mathrm{C}\left(\mathrm{s}\right) + {\mathrm{H}}_2\mathrm{O}\left(\mathrm{g}\right) \to \mathrm{CO}\left(\mathrm{g}\right) + {\mathrm{H}}_2\left(\mathrm{g}\right)$. Calculate the volume at STP from 48 gm of carbon and excess ${\mathrm{H}}_2\mathrm{O}-$

Two gases occupy two containers A and B the gas in A, of volume 0.10 ${\mathrm{m}}^3$, exerts a pressure of 1.40 MPa and that in B of volume 0.15 ${\mathrm{m}}^3$ exerts a pressure 0.7 MPa. The two containers are united by a tube of negligible volume and the gases are allowed to intermingle. Then if the temperature remains constant, the final pressure in the container will be (in MPa)

The average molecular weight of air is $28.8 \mathrm{g} {\mathrm{mol}}^{–1}$. At 20ºC, the pressure of air at a height of 6 km is half of that at the sea level. Assuming that air contains minute quantities of hydrogen, at what height the partial pressure of hydrogen would be one fourth of the partial pressure at the sea level ? The temperature may be assumed to be the same.

2.3 grams of a mixture of ${\mathrm{NO}}_2$ and ${\mathrm{N}}_2{\mathrm{O}}_4$ have a pressure of 0.82 atm at temperature, T K and volume, V litres. If $\frac{\mathrm{V}}{\mathrm{T}}=\frac{1}{300}$, calculate ${{\mathrm{P}}_{\mathrm{NO}}}_2.$ Assume that all the ${\mathrm{NO}}_2$ was converted into ${\mathrm{N}}_2{\mathrm{O}}_4$.

40 milligram diatomic volatile substance $\left({\mathrm{X}}_2\right)$ is converted to vapour that displaced 4.92 mL of air at 1 atm and 300 K. Atomic weight of element X is nearly:

An evacuated glass vessel weighs $50.0 \mathrm{g}$ when empty, $148.0 \mathrm{g}$ when filled with a liquid of density $0.98 \mathrm{g} {\mathrm{mL}}^{–1}$ and 50.5 g when filled with an ideal gas at $760 \mathrm{mm} \mathrm{Hg}$ at $300 \mathrm{K}$. Determine the molar mass of the gas.

The average speed at ${\mathrm{T}}_1$ (in kelvin) and the most probable speed at ${\mathrm{T}}_2$ (in kelvin) of ${\mathrm{CO}}_2$ gas is $9.0 \times 104 \mathrm{cm} {\mathrm{s}}^{–1}$. Calculate the values of ${\mathrm{T}}_1$ and ${\mathrm{T}}_2$.

Two flasks of equal volume are connected by a narrow tube (of negligible volume) at 27ºC and contain 0.70 mole of ${\mathrm{H}}_2$ at 0.5 atm. One of the flasks is then immersed in a hot bath, kept at 127 ºC, while the other remains at 27 ºC. The final pressure is -

A gas bulb of 1 litre capacity contains $2.0 \times {10}^{21}$ molecules of nitrogen exerting a pressure of $7.57 \times {10}^3 {\mathrm{Nm}}^{–2}$. Calculate the root mean square (rms) speed and the temperature of the gas molecules. If the ratio of most probable speed to the root mean square speed is 0.84, calculate the most probable speed for these molecules at this temperature.

A mixture of 10 ml ${\mathrm{CH}}_4, {\mathrm{C}}_2{\mathrm{H}}_4 \mathrm{and} {\mathrm{C}}_2{\mathrm{H}}_2$ has a vapour density of 11.3. Mixture contains x ml of ${\mathrm{CH}}_4$ , y ml of ${\mathrm{C}}_2{\mathrm{H}}_4$ and z ml of ${\mathrm{C}}_2{\mathrm{H}}_2$. When 30 ml of oxygen are sparked together over aqueous KOH, the volume contracts to 5.5 ml and then disappears when pyrogallol is introduced. If volumes are measured in the same conditions of pressure, temperature and humidity, value of x, y and z is–

The pressure in bulb dropped from 2000 to 1500 mm Hg in 47 mins. when the ${\mathrm{O}}_2$ present in the bulb leaked through a small hole. The bulb was then completely evacuated. A mixture of ${\mathrm{O}}_2$ and another gas of molecular weight of 79 in the molar ratio 1 : 1 at a total pressure of 4000 mm Hg was introduced. Find the mole ratio of two gases remaining in the bulb after a period of 74 mins.