Diffusion
1.2 Diffusion
1. Describe and explain diffusion in terms of kinetic particle theory
1.1 Definition of Diffusion
- Diffusion is the spreading out and mixing of particles of one substance with the particles of another due to their random motion.
- It happens without the need for stirring and without requiring external forces.
- It occurs in gases and liquids because their particles are free to move; in solids, diffusion is extremely slow because particles are fixed in a lattice.
1.2 Everyday Examples of Diffusion
- Gases:
- Smell of perfume spreading through a room.
- Cooking smells moving from the kitchen into other rooms.
- Spread of cigarette smoke in air.
- Liquids:
- Colour from potassium permanganate (KMnO₄) crystals slowly spreading in water without stirring.
- Salt diffusing in water.
1.3 Kinetic Particle Theory (KPT) Explanation
- In gases:
- Gas particles are far apart, move rapidly in random directions.
- Random motion means particles spread into available space.
- Collisions with other particles and container walls cause change of direction but do not stop motion.
- Over time, the concentration becomes uniform throughout the space.
- In liquids:
- Particles are close but can slide over one another.
- Movement is slower than in gases due to stronger forces between particles and higher density.
- In solids:
- Particles vibrate in fixed positions.
- Diffusion occurs only very slowly if atoms can swap places in the lattice (e.g., diffusion of one metal into another at high temperature).
1.4 Factors Affecting Rate of Diffusion
- Temperature:
- Higher temperature → higher kinetic energy → faster particle motion → faster diffusion.
- Example: Sugar dissolves and diffuses faster in hot water than cold.
- State of matter:
- Fastest in gases → slower in liquids → very slow in solids.
- Concentration gradient:
- Steeper gradient → faster diffusion rate until equilibrium.
- Presence of barriers:
- Porous membranes or small openings slow down diffusion.
1.5 Diffusion in Air Demonstrations
- Ammonia and Hydrogen Chloride Experiment:
- A long glass tube with cotton wool soaked in ammonia solution (NH₃) at one end and hydrochloric acid (HCl) at the other.
- Gases diffuse towards each other; where they meet, they form a white ring of ammonium chloride (NH₄Cl).
- The ring forms closer to the HCl end because NH₃ has a lower relative molecular mass and diffuses faster.
- Bromine and Air Experiment:
- Two gas jars: one with bromine vapour, the other with air.
- When lids are removed, bromine diffuses into the air-filled jar and air diffuses into the bromine jar until both jars have uniform colour.
1.6 Diffusion in Liquids Demonstrations
- Potassium Permanganate in Water:
- A few crystals placed in water at rest will slowly colour the entire container.
- Faster in warm water due to greater particle speed.
- Copper(II) Sulfate Solution Layering:
- A dense CuSO₄ solution placed beneath clear water.
- Over time, blue Cu²⁺ ions diffuse upward, mixing with water.
2. Describe and explain the effect of relative molecular mass on the rate of diffusion of gases
2.1 Definition of Relative Molecular Mass (Mr)
- Relative Molecular Mass (Mr) is the sum of the relative atomic masses (Ar) of all atoms in a molecule.
- Example: NH₃ → (14 + 3×1) = 17; HCl → (1 + 35.5) = 36.5.
2.2 Graham’s Law of Diffusion
- Statement: The rate of diffusion of a gas is inversely proportional to the square root of its relative molecular mass.
- Mathematical form:
Rate ∝ 1/√Mr
OR
rate₁ / rate₂ = √(Mr₂ / Mr₁)
2.3 Implications
- Gases with lower relative molecular mass diffuse faster than those with higher Mr.
- Example:
- Ammonia (NH₃): Mr = 17
- Hydrogen Chloride (HCl): Mr = 36.5
- Ratio of diffusion rates:
rate(NH₃)/rate(HCl) = √(36.5 / 17) ≈ 1.46
→ NH₃ diffuses about 1.46× faster than HCl.
2.4 Experimental Demonstration – NH₃ and HCl
- Setup:
- Long horizontal glass tube.
- Cotton wool soaked in concentrated NH₃ at one end; cotton wool soaked in concentrated HCl at other end.
- Observation:
- White ring of NH₄Cl forms closer to HCl end.
- Explanation:
- NH₃ has lower Mr → faster diffusion → meets slower-moving HCl closer to HCl source.
2.5 Experimental Demonstration – Hydrogen vs Oxygen
- Setup:
- Two gases released into opposite ends of a chamber.
- Data:
- Mr(H₂) = 2, Mr(O₂) = 32
- Ratio of diffusion rates:
rate(H₂)/rate(O₂) = √(32 / 2) = √16 = 4 - Hydrogen diffuses 4× faster than oxygen.
2.6 Real-life Applications
- Smell detection:
- Lighter gases (e.g., natural gas – mainly methane, Mr = 16) spread faster, which is why gas leaks are detected quickly.
- Respiration:
- Oxygen (O₂, Mr = 32) diffuses from alveoli into blood; carbon dioxide (CO₂, Mr = 44) diffuses out; diffusion rate is influenced by molecular mass and solubility.
- Industrial processes:
- Diffusion rate differences used in gas separation, e.g., uranium isotope enrichment.
3. Summary Table – Diffusion and Molecular Mass
| Gas | Mr | Relative Speed (to each other) | Notes |
|---|---|---|---|
| Hydrogen | 2 | Fastest | Very low molecular mass |
| Helium | 4 | Slightly slower than hydrogen | Used in balloons |
| Ammonia | 17 | Faster than most air gases | Strong smell |
| Oxygen | 32 | Slower | Supports combustion |
| Chlorine | 71 | Much slower | Heavy gas, green-yellow colour |
4. Kinetic Particle Theory Recap Applied to Diffusion
- Particle motion: Random and continuous in gases and liquids.
- Temperature effect: Higher temperature → greater kinetic energy → faster motion → faster diffusion.
- Mass effect: Lighter particles have higher average speeds at the same temperature (from kinetic energy formula KE = ½mv²).
5. Key Points for Exam Answers
- Always mention random motion of particles in KPT explanations.
- State that diffusion occurs from a region of high concentration to low concentration.
- For molecular mass effect: Quote Graham’s Law or describe using comparative particle speeds.
- Use experimental examples (NH₃ & HCl) to illustrate molecular mass effect