Gas, solid and liquid. (Or some are more familiar with the Malay equivalent of gas, pepejal dan cecair) These are the three classical states of matter. Among the three, liquids and gases can be grouped together to be fluids. So some students get confused when they try to read out of the syllabus. Concentrating back on gases, students typically underestimate the difficulty of “gas” when they study chemistry. Admittedly, it is rather easy to understand gases. One word of caution though from Berry Berry Easy because it is rather hard to master gas, especially gas related calculations. (So pay more attention to it).
At the minimum, students must try to master the concept Ideal Gas Law as soon as possible before they proceed to the calculations. Ideal Gas Law is simple but yet for some unknown reasons, students find it hard to make sense out of it. You will use the Ideal Gas Law in most of the calculations even up to university level for most engineering courses. So for the sake of your STPM and future studies, try and master this simple yet surprisingly difficult (oxymoron) topic of Gas. We’ll kick off with some revision of the states of matter.
STPM Chemistry Form 6 – Terminology and Concepts: Gas
Kinetics Theory of Matter
- describe the behaviour of particles in solids, liquid and gas.
- particles are held rigidly in fixed positions by strong attractive forces in an orderly arrangement;
- particles cannot move freely;
- particles can only vibrate or rotate about their mean position;
- particles have less energy (compared to liquids and gases);
- solids cannot be compressed;
- solids have fixed shapes;
- solids have fixed volume
- particles are packed closely together in cluster;
- particles are not in an orderly arrangement;
- particles can vibrate, rotate and move freely;
- particles have more energy (compared to solids) but have less energy (compared to gases);
- liquids are not easily compressed;
- liquids have no fixed shape (take the shape of the container);
- liquids have fixed volume.
- particles are separated from each other by distance far greater than their own size;
- particles have no forces between the particles.
- particles are not in an orderly arrangement;
- particles can vibrate, rotate and move freely within the container;
- particles have more energy (compared to liquids and solids);
- particles are in constant random motion, moving in straight lines;
- particles collide (elastic) with the walls of the container, they exert a pressure on the container and there is no loss of kinetics energy during the collision;
- gases are easily compressed;
- gases have no fixed shape (take the shape of the container);
- gases have no fixed volume.
Kinetics Theory of Gases
- describe the behaviour of ideal gas.
- the average kinetics energy of gases particles is directly proportional to the absolute temperature of the gas (Kelvin).
- four assumptions associated with this theory:
i.) particles are small compared to the distances between particles that their volumes are negligible.
ii.) particles move in straight lines. The direction of a particle’s motion is changed only by its collision with either another molecule or the walls of the container. All the collisions are to be elastic (no loss of energy).
iii.) particles are in constant random motion. Gas pressure is only caused by collisions of the particles against the walls of the container.
iv.) Gas molecules exhibit no intermolecular forces. The particles neither attract nor repel one another.
- three common gas laws to know: Avogadro’s Law, Boyle’s Law and Charles’ Law – A, B and C laws of gases.
(If you find yourself about to get confused, here is a simple story about how the scientist, Avogadro might have made his discovery: Avogadro was into counting big numbers, so his law focuses on the number of molecules. Therefore, Avogadro’s law deals with the relationship between moles of gas and volume. Big Boy Boyle sat on his lunch and smashed it (decreased the volume of his sandwich), by increasing the pressure on it. Therefore, Boyle’s law deals with the relationship between pressure and volume. Good ol’ Chuck overheated his popcorn and it scattered all over (increased its volume). Therefore, Charles’s law deals with the relationship between temperature and volume.) taken from General Chemistry Part II Sections VI-X pg 13. (2001) Berkeley Review.
1. Avogadro’s Law
- Amedeo Avogadro (1811)
- equal volumes of all gases at the same temperature and pressure contain equal numbers of molecules.
V / n = k (a constant)
V1 / n1 = V2 / n2
Where n = number of moles of gas
* Molar volume of a gas (volume occupied by 1 mol of any gas) at standard temperature and pressure (s.t.p.) is 22.4 dm3 (Condition: 0˚C / 273 K and 101.3 kNm-1 / 1 atm.).
2. Boyle’s Law
- Robert Boyle (1662)
- the volume occupied by fixed mass of gas is inversely proportional to its pressure at constant temperature.
- applies under isothermal conditions in a closed container.
pV = k (a constant)
p1V1 = p2V2
* Real gases obey Boyle’s law only at low pressures and high temperatures (ideal gas or perfect gas).
* Real gases do not obey Boyle’s law at high pressures and low temperatures (non-ideal behaviour).
3. Charles’ Law
- Jacques Charles (1780)
- the volume occupied by fixed mass of gas is directly proportional to its absolute temperature at constant pressure.
V / T = k (a constant)
V1 / T1 = V2 / T2
* Temperature is the absolute temperature (-273˚C / 0 K)
* Absolute temperature scale (Kelvin scale) as the temperature -273˚C was adopted as the ‘zero’.
Ideal Gas Equation
Combining Avogrado’s law, Boyle’s law and Charles’ law
- Ideal gas equation:
pV = nRT
where R is a constant and its value of 8.31 J mol-1 K-1
pressure: Pa or Nm-2 (1 atm = 101 kPa)
volume: m3 (1 cm3 = 1 x 10-6 m3; 1 dm3 = 1 x 10-3 m3)
n = m / Mr
where m = mass of gas and Mr = relative molecular mass of gas
m / V = ρ
where ρ = density of a gas
4. Dalton’s Law
- the total pressure of a mixture of gases do not react is the sum of the partial pressures of the constituent gases on the mixture.
PT = PA + PB + PC + …
where PT = total pressure of the mixture and
PA, PB, PC = partial pressure of gases A, B and C.
Mole fraction of A (XA) in a mixture of A and B
= (number of moles of A) / (total number of moles of A + B)
= nA / (nA + nB)
PA = PT x XA
where PT = total pressure, PA = partial pressure of gas A, XA = mole fraction of gas A
5. Deviation from Ideal Behaviour
- molecular size
- intermolecular forces
Positive deviation (volume of gas molecules):
- low pressures (molecules are very far apart – volume of the gas molecules by comparison is extremely small and can be ignored)
- high pressures (molecules are closer together – volume of the gas molecules cannot be ignored)
Negative deviation (intermolecular forces of attraction):
- low temperature (intermolecular forces of attraction between the molecules will reduce the force exerted by the impact of the molecules collide the wall of container. Pressure exerted by the gas is reduced).
- high temperature (kinetics energy of the molecules is so high that the intermolecular forces between gas molecules can be ignored).
Negative deviation (polar bonds)
- Least deviation – hydrogen gas (small molecular size and non-polar. It possesses very weak intermolecular forces of attraction).
- Marked deviation from ideal behaviour – carbon monoxide gas (polar bonds. It possesses stronger intermolecular forces)