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UNIT 9: GASES
GAS PRESSURE
Composition of dry air: N (78%), O (21%), others (1%)
Pressure is defined as the force exerted on a given area
- In the context of gases, it arises due to the collision of gas molecules with the walls of their container.
Mathematically, pressure (P) is expressed as: P = F/A
Where:
● F = Force exerted (in newtons, N) ● A = Area (in square meters, m^2)
PRESSURE CONVERSIONS:
1.013 bar = 1 atm = 760 mmHg = 760 torr = 14.7 psi = 101.3 kPa = 101300 Pa
Tools for Measuring Gas Pressure
● Barometer : Measures atmospheric pressure. It uses a column of mercury or liquid, where the height of the liquid is proportional to atmospheric pressure. ● Manometer : Measures gas pressure relative to atmospheric pressure. It typically consists of a U-shaped tube filled with a liquid (like mercury or water). ○ Open-end manometers compare gas pressure with atmospheric pressure. ○ Closed-end manometers measure absolute gas pressure. ● Pressure Gauges : Devices like Bourdon gauges or digital pressure sensors measure pressure in closed systems.
At constant temperature and pressure, a gas expands uniformly to fill the container in which it is contained.
All gases have the same dependence on those four properties: volume, amount, temperature, pressure. All are combined to create…
THE IDEAL GAS LAW: PV = nRT
Standard temperature and pressure (STP)
- 0 oC = 273 K
- 1 atm
- 1 mol @ 22.4L
DENSITY OF GAS: d = M (P/RT) where M is molar mass, P is pressure, T is temp.
MOLAR MASS OF GAS: PV = m/M (RT) where M = grams/mole = m/n is changed to n = m/M
Dalton’s law of partial pressure: Total pressure of a mixture of ideal gases is the sum of the partial pressures of individual gases.
PT = PA + PB + PC …
● Partial Pressure: The pressure that each gas would exert if it occupied the container alone.
The Kinetic-Molecular Theory of Gases:
- Gas particles are in constant, random motion (move in straight lines until they collide with each other or the walls of their container)
- Gas particles occupy negligible volume ( mostly empty space )
- Collisions between gas particles and the container walls are elastic (no energy is lost during collisions; the total kinetic energy of the system remains constant).
- No intermolecular forces between gas particles (gas particles do not attract or repel each other, behaving independently)
- The average kinetic energy of gas particles is proportional to the temperature (in Kelvin). ○ Higher temperatures result in faster-moving particles, increasing their kinetic energy.
Non-Ideal Gas Behavior:
Intermolecular Forces:
● At low temperatures or high pressures, gas particles are close together. ● Attractive forces (e.g., van der Waals forces) reduce the force of particle collisions with the container, causing the pressure to be lower than predicted by the Ideal Gas Law.
Finite Volume of Gas Particles:
● Gas molecules occupy a non-negligible volume. ● At high pressures (small volumes), the actual free space available for particle motion is less than the container volume, leading to deviations.
Van der Waals equation modifies the Ideal Gas Law to account for intermolecular forces and the finite volume of particles:
