Hey Tara! Welcome to your very first Chemistry topic. States of Matter might look simple at first glance, but it is the foundation for almost everything else in Chemistry - so let's build it properly. We are going to think about matter as tiny particles that are constantly moving, and use that one idea to explain everything from why chai steams to why you can smell samosas frying from across the house. Take it slowly, use the diagrams, and try every question. You've got this!
Properties of Solids, Liquids and Gases
Everything around you - your steel water bottle, the water inside it, and the air you breathe - is made of tiny particles (atoms, molecules or ions). What makes something a solid, a liquid or a gas is simply how those particles are arranged and how they move. Let's compare them properly.
| Property | Solid | Liquid | Gas |
|---|---|---|---|
| Shape | Fixed shape | Takes the shape of the container (flows to the bottom) | Fills the whole container completely |
| Volume | Fixed volume | Fixed volume | No fixed volume - expands to fill any space |
| Compressibility | Cannot be compressed (squashed) | Very difficult to compress | Easily compressed |
| Density | High density | Slightly lower than solid, still high | Very low density |
| Ability to flow | Does not flow | Flows easily | Flows easily and spreads out |
Remember the order of increasing particle energy and spacing with "SLC" - Solid, Liquid, Gas (say it as "Silly Lizards Gallop") - particle energy and the space between particles increase as you go from Solid to Liquid to Gas.
Particle Arrangement, Separation and Motion
Now let's look INSIDE each state. This is the part examiners love to ask about, so make sure you can describe all three properly: arrangement (how they're packed), separation (how far apart), and motion (how they move).
Solids
- Arrangement: Particles are packed in a regular, ordered pattern (like a crate of mangoes stacked neatly at a Bangalore fruit market).
- Separation: Particles are very close together, touching their neighbours.
- Motion: Particles do NOT move from place to place. They can only vibrate (jiggle) about a fixed position. This is why a solid keeps its shape.
- Forces of attraction: Very strong forces hold the particles in place.
Liquids
- Arrangement: Particles are randomly arranged - no neat pattern - but still close together.
- Separation: Particles are close together, touching, but with slightly more space than in a solid.
- Motion: Particles can move around and slide past each other. This is why liquids flow and take the shape of their container.
- Forces of attraction: Weaker than in a solid, but still fairly strong - strong enough to keep the particles close.
Gases
- Arrangement: Particles are randomly arranged with no pattern at all.
- Separation: Particles are very far apart compared to their own size - there is mostly empty space between them.
- Motion: Particles move rapidly and randomly in straight lines, in all directions, until they collide with something (another particle or the container wall).
- Forces of attraction: Extremely weak - almost negligible, which is why gas particles can fly off in every direction.
Never say particles "get bigger" or "expand" when heated. The particles themselves stay exactly the same size. What changes is the SPACE between them and how fast they move. This is one of the most common mark-losing mistakes in IGCSE Chemistry.
Changes of State
Matter can change from one state to another. Each change has a specific name, and you must know all of them - including which direction they go and what energy is doing.
| Change of State | From → To | What Happens to Energy |
|---|---|---|
| Melting | Solid → Liquid | Energy is absorbed (taken in) |
| Boiling | Liquid → Gas (throughout the liquid, at ONE fixed temperature) | Energy is absorbed |
| Evaporation | Liquid → Gas (only at the surface, at ANY temperature) | Energy is absorbed |
| Freezing | Liquid → Solid | Energy is released (given out) |
| Condensing | Gas → Liquid | Energy is released |
| Sublimation | Solid → Gas directly (no liquid stage), and the reverse: Gas → Solid | Energy is absorbed (solid to gas) |
Melting and Freezing
Melting is when a solid changes to a liquid. This happens at a specific temperature called the melting point. For pure ice, this is exactly 0 °C at normal atmospheric pressure. Freezing is the exact reverse - a liquid turning into a solid, at the same fixed temperature (the freezing point, which equals the melting point for a pure substance).
Indian example: When you make ice candy (kulfi or ice popsicles) at home in Bangalore, you pour flavoured liquid into moulds and place them in the freezer. The liquid freezes into a solid as heat is removed by the freezer.
Boiling and Evaporation - Know the Difference!
This is one of the most important comparisons in this whole topic, and IGCSE loves to test it directly.
| Feature | Boiling | Evaporation |
|---|---|---|
| Where it happens | Throughout the whole liquid (bubbles form inside the liquid, not just at the top) | Only at the surface of the liquid |
| Temperature | Happens at ONE fixed temperature (the boiling point) | Happens at ANY temperature, even well below the boiling point |
| Speed | Fast - large amounts of liquid become gas quickly | Slow - only the fastest-moving surface particles escape |
| Energy source | Needs continuous heating from an external source | Can happen using energy already present (e.g. from the surroundings), no external heat required |
Evaporation happens at ANY temperature; boiling happens at ONE specific temperature. This is exactly why a puddle of water on the road dries up in the sun (evaporation, at around 30 °C) without ever reaching 100 °C. Water only boils at 100 °C at normal atmospheric pressure.
Why does evaporation happen at any temperature? Even in a liquid at room temperature, not all particles move at the same speed - there is a range of speeds. A few particles at the surface happen to be moving fast enough (have enough energy) to break away from the attractive forces of their neighbours and escape as a gas, even though the average temperature of the liquid is well below the boiling point.
Indian example: Wet clothes hung out to dry on a terrace in Bangalore dry through evaporation, not boiling - the water definitely does not reach 100 °C, but the fastest-moving water particles at the surface still escape into the air over time, especially helped along by wind and sunlight.
Sublimation
Sublimation is the direct change from a solid to a gas (or gas to solid) without passing through the liquid state at all. Only a few substances do this - the classic examples are solid carbon dioxide (dry ice), iodine, and ammonium chloride.
Example: Dry ice (solid CO₂) used to keep ice cream cold during transport does not melt into a puddle - it sublimes straight into carbon dioxide gas, which is why you see that dramatic white "smoke" (actually condensed water vapour in the cold gas).
Do not confuse "evaporate" with "boil" in your exam answers - examiners specifically check for this. Also remember: melting, boiling, evaporating and subliming (solid to gas) all need energy IN; freezing, condensing and subliming (gas to solid) all give energy OUT.
Effect of Temperature and Pressure on Gas Volume
Effect of Temperature on Volume
If you keep the pressure on a gas constant and increase its temperature, the volume of the gas increases. Heating a gas makes it expand.
Indian example: A rubber balloon left out on a hot afternoon in Chennai will appear slightly larger than the same balloon kept in an air-conditioned room, because the warmer air inside has expanded.
Effect of Pressure on Volume
If you keep the temperature of a gas constant and increase the pressure on it (squeeze it into a smaller space), the volume of the gas decreases. Higher pressure means a smaller volume, and lower pressure means a larger volume - the two are inversely related.
Indian example: When you use a bicycle pump, you push the piston down, squeezing the same amount of air into a smaller and smaller volume inside the pump chamber, which is why the pump gets harder to push as you compress the air.
Keep these two relationships separate in your head: higher temperature → larger volume (at constant pressure), and higher pressure → smaller volume (at constant temperature). Examiners often ask you to describe one while keeping the other constant.
Kinetic Particle Theory and Heating/Cooling Curves
Now let's explain WHY changes of state happen, using kinetic particle theory - the idea that particles are always moving, and that temperature is a measure of how much kinetic (movement) energy they have on average.
What Happens When You Heat a Solid
When you supply heat energy to a solid, the particles absorb this energy and vibrate more vigorously about their fixed positions. Eventually, the particles vibrate so much that they gain enough energy to overcome (partly break) the strong forces of attraction holding them in their fixed positions, and the solid starts to melt.
The Heating Curve
If you heat a solid steadily and plot temperature against time, you get a graph with a very distinctive shape - flat sections mixed with rising sections. This is called a heating curve, and reading it correctly is a classic exam skill.
Why Are There Flat Sections on the Heating Curve?
This is the key idea to explain clearly in your answer. During melting and boiling, heat energy is STILL being supplied continuously - but the temperature does not rise. Why not?
The energy being supplied during a flat section is being used entirely to weaken and overcome the forces of attraction between particles (to break the particles free from their fixed/close arrangement), NOT to increase the average kinetic energy (speed) of the particles. Since temperature is a measure of average kinetic energy, and the kinetic energy is not increasing during this stage, the temperature stays constant even though heating continues.
"Flat means Fighting" - on a heating curve, a flat section means the energy is "fighting" to break the forces of attraction between particles, not raising the temperature.
Never write "no energy is being supplied" to explain a flat section on a heating curve - this is WRONG and loses marks. Energy IS being supplied continuously; it is being used to break/weaken forces of attraction between particles instead of raising temperature.
The Cooling Curve
A cooling curve is the mirror image. As a gas cools, particles slow down and lose kinetic energy - temperature falls. When the gas reaches its condensation point, temperature stays constant (flat section) while particles come close enough for attractive forces to pull them into the liquid state (energy is being released as bonds/attractions form). The same happens again at the freezing point as the liquid becomes a solid.
Kinetic Theory Explanation of Temperature and Pressure Effects on Gas Volume
Now let's explain (not just describe) why temperature and pressure affect the volume of a gas, using kinetic particle theory.
Why does increasing temperature increase volume (at constant pressure)? Heating a gas gives its particles more kinetic energy, so they move faster. Faster-moving particles hit the walls of the container more often and with greater force, increasing the pressure they would exert. To keep the pressure constant (as stated), the container must expand - the particles need more space to spread into so that the frequency and force of collisions with the walls returns to the original (constant) pressure. So the volume increases.
Why does increasing pressure decrease volume (at constant temperature)? If you squeeze a gas into a smaller volume, the same number of particles are now confined to a smaller space. This means the particles collide with the walls of the container more frequently (they have less distance to travel between collisions), which increases the pressure. Conversely, if you want to increase the pressure applied to a gas while keeping temperature (and therefore average particle speed) constant, the particles must be squeezed into a smaller volume so that collisions with the walls happen more often, generating higher pressure.
When explaining gas pressure using kinetic theory, always mention collisions between particles and the container walls. Pressure is caused by billions of tiny particle collisions against the walls per second - more frequent or more forceful collisions mean higher pressure.
2. Evaporation: Even before boiling starts (and even at the liquid surface while it boils), fast-moving water particles at the surface escape into the air as steam - this is evaporation, happening continuously at the surface.
3. Condensation: The water vapour (steam) that hits the cool underside of the lid loses energy to the metal lid, slows down, and the particles come close enough for forces of attraction to pull them back into liquid droplets - this is condensation.
Because liquid water takes up slightly less volume than the same mass of ice, high pressure encourages the ice to change into the lower-volume state - liquid water - at a slightly lower temperature than normal. This creates a very thin film of liquid under the blade.
This thin liquid film reduces friction between the blade and the ice, letting the skater glide smoothly.
As the temperature rises from the morning to midday, the air particles inside the tyre gain kinetic energy and move faster.
Because the tyre's volume cannot expand much (it is a rigid, fixed shape), the faster-moving particles collide with the inside walls of the tyre more frequently and with greater force.
More frequent, more forceful collisions with the tyre walls means higher pressure - which is exactly why the tyre feels harder.
In the very low-pressure chamber, the particles at the surface of the ice absorb enough energy to overcome the forces of attraction completely and escape directly as gas particles (water vapour), without first forming a liquid layer.
The water vapour is then pumped away, leaving the solid coffee solids behind as dry, light granules.
In a sealed pressure cooker, the trapped steam builds up extra pressure above the liquid. This higher pressure pushes down harder on the liquid surface, making it more difficult for particles to escape into the gas state.
As a result, the water particles need MORE energy (a higher temperature, often around 120 °C) before they can overcome this increased pressure and boil.
Cooking at this higher temperature makes chemical and physical changes in the food (like softening the dal) happen faster, which is why pressure-cooked food cooks in a fraction of the time.
What Is Diffusion?
Diffusion is the net (overall) movement of particles from a region of higher concentration to a region of lower concentration, as a result of the random motion of particles. It happens in liquids and gases (not really in solids, because particles there cannot move from place to place).
Diffusion is net movement from HIGH to LOW concentration. Always use the word "net" in your definition - individual particles move randomly in all directions, but overall, more particles move from the crowded area to the less crowded area until concentration is even everywhere.
Explaining Diffusion Using Kinetic Particle Theory
Particles in gases and liquids are always moving randomly, in all directions, because they possess kinetic energy. If you release a gas or a coloured substance in one spot, its particles do not simply "spread out" on purpose - they move around randomly, colliding with other particles. Because there are initially more particles in the high-concentration region, more particles happen to move OUT of that crowded region than move back INTO it (simply because there are more of them there to begin with). Over time, this random motion causes a net movement from high to low concentration, until the particles are evenly spread throughout the available space.
The Bromine Jar Experiment
A classic demonstration of diffusion in gases: a gas jar of orange-brown bromine gas is placed at the bottom, with a gas jar of air placed upside-down on top of it (connected mouth-to-mouth), separated by a glass cover slide. When the slide is removed, the denser bromine gas particles gradually move upward and the air particles move downward. Over several minutes, the orange-brown colour spreads evenly throughout both jars, showing that diffusion has occurred - bromine particles have moved by diffusion from the high-concentration jar into the low-concentration jar (and vice versa for air particles), even against gravity, purely due to random particle motion.
The Ammonia and Hydrogen Chloride Experiment
This is one of the most important diffusion experiments in the whole IGCSE syllabus, so learn it in detail. A long glass tube is set up horizontally. A piece of cotton wool soaked in concentrated ammonia solution (NH₃) is placed at one end, and a piece of cotton wool soaked in concentrated hydrochloric acid (HCl) is placed at the other end, at the same time. Both are sealed with bungs.
Ammonia gas (NH₃) and hydrogen chloride gas (HCl) both diffuse from their respective ends toward the middle of the tube. When they meet, they react to form a white solid ring of ammonium chloride (NH₃Cl):
The key observation: the white ring does NOT form exactly in the middle of the tube. It forms closer to the hydrochloric acid end.
If asked to explain why the ring forms closer to the HCl end, you must mention BOTH ideas: (1) ammonia has a smaller relative molecular mass than hydrogen chloride, and (2) lighter (less massive) particles move faster and diffuse faster, so ammonia travels further along the tube in the same time.
Effect of Relative Molecular Mass on Rate of Diffusion
Not all gases diffuse at the same speed. The rate of diffusion depends on the relative molecular mass (Mₜ) of the gas - lighter (less massive) particles diffuse faster than heavier particles, at the same temperature.
Why? At a given temperature, ALL gas particles have (on average) the same amount of kinetic energy, regardless of their mass. Since kinetic energy depends on both mass and speed (KE = ½mv²), a particle with a SMALLER mass must be moving at a HIGHER speed to have the same kinetic energy as a heavier particle. Faster-moving particles diffuse faster.
"Small is Speedy" - the smaller (lighter) the relative molecular mass of a gas, the faster it diffuses.
Applying This to the Ammonia-HCl Experiment
Let's calculate the relative molecular masses to see exactly why ammonia diffuses faster:
- Ammonia, NH₃: Mₜ = 14 (N) + 3 × 1 (H) = 14 + 3 = 17
- Hydrogen chloride, HCl: Mₜ = 1 (H) + 35.5 (Cl) = 36.5
Ammonia has a much smaller relative molecular mass (17) than hydrogen chloride (36.5). This means ammonia particles move faster and diffuse faster, travelling further along the tube before meeting the HCl particles. This is exactly why the white ring of ammonium chloride forms closer to the HCl end - the ammonia has "won the race" and travelled further.
You do NOT need to memorise a formula linking rate of diffusion and relative molecular mass for IGCSE (that level of detail, Graham's Law as a formula, is beyond this syllabus). You DO need to know the qualitative rule: lower Mₜ = faster diffusion, and be able to compare two given relative molecular masses to predict which gas diffuses faster.
These gas particles move randomly and rapidly in all directions due to their kinetic energy, constantly colliding with air particles (nitrogen, oxygen etc).
Because there are initially far more scent particles near the incense than near the gate 15 m away, there is a net movement of scent particles from high concentration (near the incense) to low concentration (near the gate) - this is diffusion.
Over a couple of minutes, enough scent particles diffuse through the air to reach Tara's nose in a high enough concentration for her to smell it.
Each collision changes the perfume particle's direction, so its overall path is a slow, zig-zag "random walk" rather than a straight line, even though individual particles move very fast between collisions.
This is why it takes tens of seconds (not milliseconds) for a noticeable concentration of perfume particles to diffuse several metres through the air, even though the particles themselves are moving at hundreds of metres per second.
However, your nose is extremely sensitive and can detect even a very low concentration of mercaptan particles.
As soon as gas starts leaking, mercaptan particles begin diffusing outward from the leak (high concentration) into the surrounding air (low concentration). Because your nose can detect even tiny concentrations, you smell the leak from a distance long before enough gas has diffused/accumulated to be dangerous or to be seen.
Faster-moving particles collide and mix more quickly, spreading the coloured tea particles from the high-concentration area (around the tea bag) to the low-concentration area (the rest of the glass) more rapidly.
In the cold water, all particles have less kinetic energy and move more slowly, so the same net movement from high to low concentration takes much longer to complete.
Ammonia (NH₃): Mₜ = 14 + (3 × 1) = 17
Hydrogen chloride (HCl): Mₜ = 1 + 35.5 = 36.5
Step 2: Apply the diffusion rule.
Ammonia has the smaller Mₜ (17), so ammonia particles move faster and diffuse faster than hydrogen chloride particles (Mₜ = 36.5).
Step 3: Predict the position of the ring.
Since ammonia travels further along the tube in the same amount of time, the two gases meet closer to the HCl end (right end) rather than in the exact middle of the tube.