Diffusion Revision Notes

Diffusion is the process by which particles spread out from an area where they are in higher concentration to an area where they are in lower concentration. This movement occurs naturally and results in the gradual mixing of substances until the concentration is uniform throughout the space.

For example, when a drop of coloured ink is placed in water, the ink particles move from the concentrated drop into the surrounding water until the colour is evenly spread.

• Diffusion occurs down a concentration gradient (high to low concentration).
• It is a passive process requiring no external energy input.

The kinetic particle theory states that all matter is made up of tiny particles (atoms, ions, or molecules) that are in constant random motion. The particles collide with each other and with the walls of their container, moving in all directions.

In gases and liquids, particles have enough energy to move freely and randomly, which allows them to spread out and mix with other particles.

The speed and energy of particle motion increase as temperature rises, causing particles to move faster.

• Particles are in constant, random motion.
• Particles collide and change direction frequently.
• Temperature affects particle speed and energy.

Diffusion occurs because particles move randomly and collide with each other. In a region of high concentration, there are more particles, so collisions are more frequent, causing particles to move away into regions of lower concentration.

Over time, this random movement and collision cause particles to spread out evenly, resulting in an equal concentration throughout the available space.

No external force is needed; diffusion is driven by the kinetic energy of particles and their random motion.

• Random motion causes particles to move from crowded to less crowded areas.
• Collisions between particles cause spreading.
• Diffusion continues until equilibrium (uniform concentration) is reached.

Increasing temperature increases the kinetic energy of particles, making them move faster. Faster-moving particles collide more often and spread out more quickly, increasing the rate of diffusion.

Conversely, at lower temperatures, particles move more slowly, and diffusion occurs more slowly.

This explains why diffusion is faster in hot air or liquids than in cold ones.

• Higher temperature → faster particle motion → faster diffusion.
• Lower temperature → slower particle motion → slower diffusion.

In gases, diffusion can be observed when a scented gas spreads through a room, gradually filling the space with its smell.

In liquids, diffusion is seen when a coloured substance, such as potassium permanganate crystals, dissolves and spreads out evenly in water without stirring.

These observations support the kinetic particle theory explanation that particles move randomly and spread from high to low concentration areas.

• Scent spreading in air is diffusion of gas particles.
• Colour spreading in water is diffusion of solute particles.
• No stirring is needed; diffusion is spontaneous.

Diffusion is the spontaneous movement of gas molecules from a region of higher concentration to a region of lower concentration, resulting from their random motion. It occurs without any external force and continues until the gases are evenly mixed.

• Diffusion depends on the random motion of gas particles.
• It results in mixing of gases over time.

The rate of diffusion of a gas is influenced by the molecular mass of its particles. At a given temperature, all gases have the same average kinetic energy, but their average speeds differ depending on their molecular masses. Lighter gas molecules move faster than heavier ones, causing them to diffuse more rapidly.

This relationship is quantitatively described by Graham's law of diffusion, which states that the rate of diffusion (r) of a gas is inversely proportional to the square root of its molecular mass (M):

\[ r \propto \frac{1}{\sqrt{M}} \] This means that if one gas has a molecular mass four times that of another, its diffusion rate will be half as fast.

• Lighter gases have higher average speeds at the same temperature.
• Rate of diffusion ∝ 1/√(molecular mass).
• Heavier gases diffuse more slowly than lighter gases.

Calculate the ratio of diffusion rates of hydrogen gas (H₂, molecular mass = 2 g/mol) and oxygen gas (O₂, molecular mass = 32 g/mol) at the same temperature.

Using Graham's law: \[ \frac{r_{\text{H}_2}}{r_{\text{O}_2}} = \sqrt{\frac{M_{\text{O}_2}}{M_{\text{H}_2}}} = \sqrt{\frac{32}{2}} = \sqrt{16} = 4 \]

This means hydrogen gas diffuses four times faster than oxygen gas under identical conditions.

• Use molecular masses in grams per mole for calculation.
• Apply the square root ratio to find relative diffusion rates.
• Hydrogen diffuses significantly faster than oxygen.

While molecular mass is a primary factor, other conditions also influence diffusion rates:

Temperature: Increasing temperature increases the average kinetic energy and speed of gas molecules, thus increasing diffusion rates.

Pressure: Higher pressure increases gas density, which can reduce diffusion rates due to more frequent collisions between molecules, although this effect is less direct than molecular mass or temperature changes.

• Higher temperature increases diffusion rate.
• Higher pressure can decrease diffusion rate due to increased collisions.
• Molecular mass remains the key factor for comparing different gases at the same temperature and pressure.

In summary, the molecular mass of a gas determines the average speed of its molecules at a given temperature, which directly affects the rate of diffusion. Graham's law provides a quantitative method to compare diffusion rates between gases. Understanding this principle is essential for explaining phenomena such as the mixing of gases, the behaviour of gases in different environments, and practical applications like gas separation techniques.

• Molecular mass inversely affects diffusion rate.
• Graham's law allows calculation of relative diffusion rates.
• Temperature and pressure modify diffusion but do not override molecular mass effects.