Introducing Properties of Gases (VCE SSCE Chemistry): Revision Notes

Introducing Properties of Gases

Understanding the properties of gases is essential for chemistry students, particularly when studying atmospheric composition, greenhouse effects, and gas behaviour. Gases behave very differently from liquids and solids, and these differences can be observed, measured, and predicted using mathematical relationships.

Comparing states of matter

Gases differ significantly from liquids and solids in several key ways. Understanding these differences helps explain why gases behave as they do.

Property	Gases	Liquids	Solids
Density	Low	High	High
Volume and shape	Fill all available space because particles move independently	Fixed volume, adopt shape of container because particles are affected by attractive forces	Fixed volume and shape because particles are affected by attractive forces
Compressibility	Compress easily	Almost incompressible	Almost incompressible
Ability to mix	Mix together rapidly	Mix together slowly unless stirred	Do not mix unless finely divided

These property differences are evident in everyday examples. Air in vehicle tyres compresses to absorb impacts. The aroma from fresh coffee rapidly fills a room because gas particles mix readily and occupy all available space. Weather balloons expand as they rise because of pressure changes in the atmosphere.

Physical properties of gases

Physical properties are those properties that can be observed and measured without changing the nature of the substance itself. For gases, physical properties include how they compress, expand, mix, and respond to changes in temperature and pressure.

Kinetic molecular theory

Scientists developed the kinetic molecular theory to explain gas behaviour based on particle movement. According to this theory:

This theoretical model helps explain why gases behave the way they do and allows us to predict their behaviour under different conditions.

Measurable quantities of gases

Four key quantities are commonly measured to describe a gas: amount, volume, pressure, and temperature. Changing any one of these affects one or more of the others.

Amount

Like solids and liquids, the amount of gas present in a sample is measured in moles (mol). The amount of gas can be calculated from its mass using the formula:

$n = \frac{m}{M}$

where $n$ is the number of moles, $m$ is the mass in grams, and $M$ is the molar mass in $\text{g mol}^{-1}$.

Volume

Volume is the space that a gas occupies. Common units for measuring gas volume include millilitres (mL) and litres (L). The relationship between these units is:

$1 \text{ L} = 1000 \text{ mL}$

To convert from millilitres to litres, divide by 1000. To convert from litres to millilitres, multiply by 1000.

Pressure

When gas particles collide with the walls of their container, they exert a force. Pressure is defined as the force applied per unit area:

$\text{pressure} = \frac{\text{force}}{\text{area}}$

The smaller the volume occupied by a gas, the more frequently particles collide with the container walls, increasing the pressure.

Several units are used to measure pressure:

Unit name	Symbol	Conversions
Pascal	Pa	1 Pa = 1 N m⁻²
Kilopascal	kPa	1 kPa = 1 × 10³ Pa
Atmosphere	atm	0.987 atm = 100 kPa

The pascal (Pa) is the SI unit for pressure and equals one newton per square metre. The standard atmosphere (atm) is the average atmospheric pressure at sea level, where 1 atm = 101.3 kPa.

Torricelli's barometer

Partial pressure

Air is a mixture of gases including nitrogen, oxygen, carbon dioxide, and argon. In a gas mixture, each gas exerts its own partial pressure. The total measured pressure equals the sum of all partial pressures.

This principle, known as Dalton's Law of Partial Pressures, is important when working with gas mixtures in the laboratory.

Temperature

Temperature is a crucial quantity for describing gases. According to kinetic molecular theory, increasing temperature increases the average kinetic energy of gas particles.

Experiments show that gas volume increases as temperature increases when pressure remains constant. When experimental data plotting volume against temperature in degrees Celsius is extrapolated to zero volume, it crosses the temperature axis at -273°C.

This observation led scientists to develop the kelvin scale (also called the absolute temperature scale). On this scale:

• Each degree increment equals one Celsius degree
• 0°C equals 273 K
• 0 K equals -273°C

The temperature 0 K is called absolute zero - the lowest temperature theoretically possible. At absolute zero, all molecules and atoms have minimum kinetic energy.

The relationship between Celsius and kelvin scales is:

$T \text{ (in K)} = T \text{ (in °C)} + 273$

Relationship between volume and temperature

For a fixed amount of gas at constant pressure, volume is directly proportional to temperature (in kelvin):

$V \propto T$

This can also be written as:

$\frac{V}{T} = k$

where $k$ is a constant for the fixed amount of gas at constant pressure.

Molar volume of a gas

The molar volume of a gas ($V_m$) is the volume occupied by one mole of gas at a particular temperature and pressure. Importantly, molar volume does not depend on the identity of the gas.

One mole of neon, one mole of nitrogen, and one mole of ozone all occupy the same volume at the same temperature and pressure, even though their masses differ. This is because all contain the same number of particles (Avogadro's number, $6.02 \times 10^{23}$).

Standard laboratory conditions

Standard laboratory conditions (SLC) refer to a temperature of 25°C (298 K) and a pressure of 100 kPa. These represent typical laboratory conditions.

An ideal gas is a theoretical gas where particles do not interact except during elastic collisions. Real gases behave very similarly to ideal gases at SLC.

Gas	Formula	Molar volume at SLC (L mol⁻¹)
Ideal gas	-	24.79
Helium	He	24.83

The amount of gas can be calculated from its volume at SLC using:

$n = \frac{V}{V_m}$

The ideal gas equation

Mathematical relationships linking volume, pressure, temperature, and amount were developed by different scientists through experiments. Robert Boyle determined the pressure-volume relationship, while Jacques Charles identified the volume-temperature relationship.

Volume and pressure

For a fixed amount of gas at constant temperature, volume is inversely proportional to pressure:

$P \propto \frac{1}{V}$

This can also be written as:

$PV = k$

where $k$ is a constant at constant temperature.

When volume decreases, particles are more closely spaced and collide more frequently with container walls, increasing pressure. When volume increases, particles are more widely spaced and pressure decreases.

Volume and amount of gas

At constant temperature and pressure, volume is directly proportional to the amount of gas:

$V \propto n$

This can also be written as:

$\frac{V}{n} = k$

where $k$ is a constant at constant temperature and pressure.

Doubling the amount of gas doubles the volume if pressure and temperature remain constant.

Combining the relationships

Three mathematical relationships describe gas behaviour:

• $V \propto T$ (for constant $P$ and $n$)
• $V \propto \frac{1}{P}$ (for constant $T$ and $n$)
• $V \propto n$ (for constant $P$ and $T$)

Calculating molar mass or molar volume

Experimental data from laboratory investigations can be used to determine the molar volume or molar mass of a gas.

Molar volume determination

If you know the mass and volume of a gas sample, you can calculate its molar volume. First, calculate the amount in moles using:

$n = \frac{m}{M}$

Then calculate molar volume using:

$V_m = \frac{V}{n}$

Molar mass determination

Molar mass can be determined from experimental measurements of mass, volume, pressure, and temperature. Use the ideal gas equation to find the amount in moles:

$n = \frac{PV}{RT}$

Then calculate molar mass using:

$M = \frac{m}{n}$