Sciences - Co-ordinated (Double) (0654) | States of matter

Welcome to States of Matter! (Chemistry Chapter C1)

Hello future scientists! This chapter is all about understanding the building blocks of chemistry: the physical states that everything around us exists in—solids, liquids, and gases.
This isn't just common sense; we will explore why these states behave differently using a super important idea called the Kinetic Particle Theory. Mastering this foundation is key to understanding everything else in Chemistry!

C1.1 Solids, Liquids and Gases

The Kinetic Particle Theory (KPT) - The Big Idea

The Kinetic Particle Theory states that all matter is made up of tiny particles (atoms, molecules, or ions) that are always moving. The state of matter depends on how much energy these particles have, and the strength of the forces holding them together.

Understanding Structure through KPT (C1.1.2)

The differences in properties come down to three things: separation, arrangement, and motion of the particles.

1. Solids

• Separation: Particles are very close together.
• Arrangement: Arranged in a fixed, regular pattern (a lattice).
  Analogy: Imagine soldiers standing rigidly in perfect rows.
• Motion: Particles only vibrate about fixed positions. They do not move past each other.
• Forces: Very strong forces of attraction between particles.

2. Liquids

• Separation: Particles are still close together, almost the same separation as in a solid.
• Arrangement: Arranged randomly; there is no fixed pattern.
  Analogy: Imagine a crowded party where people are close but moving around freely.
• Motion: Particles move randomly and slide past each other, allowing the liquid to flow.
• Forces: Forces are weaker than in solids, allowing particles to move.

3. Gases

• Separation: Particles are very far apart (the distance between them is about 10 times their size).
• Arrangement: Arranged randomly.
• Motion: Particles move rapidly and randomly in all directions.
• Forces: Extremely weak forces of attraction (almost negligible).

Key Takeaway: The state of matter depends entirely on how closely packed the particles are and how strongly the forces hold them together, controlling their ability to move.

C1.1 Changes of State

Describing Phase Changes (C1.1.3 Core)

We need to know the specific names for the six changes between states:

• Melting: Solid to Liquid (e.g., ice becoming water).
• Boiling: Liquid to Gas (happens throughout the liquid, e.g., water boiling vigorously).
• Evaporating: Liquid to Gas (happens only at the surface, below the boiling point, e.g., a puddle drying up).
• Freezing (Solidification): Liquid to Solid (e.g., water becoming ice).
• Condensing: Gas to Liquid (e.g., steam hitting a cold mirror).

Explaining Changes of State using KPT (C1.1.5 Supplement)

What happens at the particle level during heating and cooling?

1. Melting (Solid $\rightarrow$ Liquid) and Boiling (Liquid $\rightarrow$ Gas)

When you heat a substance, the particles gain kinetic energy and vibrate faster. When enough energy is gained, the particles can break free from their fixed positions (melting) or overcome the remaining forces completely (boiling).

• Melting Point: The specific temperature where the strong forces holding the solid lattice break down, allowing particles to slide past one another.
• Boiling Point: The specific temperature where particles gain enough energy to completely overcome the forces of attraction and escape as a gas.

2. Evaporation vs. Boiling

Evaporation occurs because particles at the liquid surface have a range of kinetic energies. The most energetic particles have enough energy to overcome the surface forces and escape into the air. Because the most energetic particles leave, the average kinetic energy of the remaining liquid particles drops, causing cooling (e.g., sweat cooling your body).

Interpreting Heating and Cooling Curves (C1.1.5 Supplement)

When you graph temperature against time during heating, you get flat sections (plateaus). These show periods where the substance is absorbing energy, but its temperature is not increasing.

Step-by-Step Explanation:

1. Sloping Section (Temperature Rises): Energy is being absorbed and converted into kinetic energy, making the particles move faster (increasing temperature).
2. Plateau (Phase Change Occurs, e.g., Melting): Energy is still being absorbed, but it is now being used as potential energy to break the forces of attraction between particles. The temperature stays constant until all the solid has turned into a liquid.
3. Next Sloping Section (Temperature Rises Again): All the substance is now liquid. Energy again becomes kinetic energy, and the temperature rises until the boiling point is reached.

Cooling curves are simply the reverse! Energy is released (exothermic) during condensation and freezing, but the temperature stays constant during these phase changes.

Key Takeaway: During a phase change, absorbed energy is used to break bonds (potential energy), not to make particles move faster (kinetic energy).

C1.1 The Behavior of Gases

Effects of Temperature and Pressure on Gas Volume (C1.1.4 Core, C1.1.6 Supplement)

Gases are highly compressible because the particles are very far apart. Their volume is highly sensitive to changes in temperature and pressure. We can explain these effects using KPT.

1. Effect of Temperature on Volume (at constant pressure)

Observation (Core): If you increase the temperature of a gas, its volume increases (it expands).

Explanation (Supplement via KPT):

• Increasing temperature gives gas particles more kinetic energy.
• They move faster and hit the walls of the container harder and more frequently.
• To keep the pressure constant (as required by the observation), the volume must increase, giving the particles more room to move.

2. Effect of Pressure on Volume (at constant temperature)

Observation (Core): If you increase the external pressure on a gas, its volume decreases (it is compressed).

Explanation (Supplement via KPT):

• The pressure exerted by a gas is caused by particles colliding with the container walls.
• When you decrease the volume (by applying external pressure), the particles have less space.
• This results in the particles hitting the container walls more frequently, increasing the internal pressure until it balances the external pressure.

Did you know? Air bags in cars rely on gases behaving exactly like this. A tiny chemical reaction produces a large volume of gas very quickly to inflate the bag!

Key Takeaway: Gas particles move randomly and rapidly. Changes in T or P directly affect how often and how hard they collide with the walls, dictating the volume.

C1.2 Diffusion

What is Diffusion? (C1.2.1 Core)

Diffusion is the net movement of particles from a region where they are in higher concentration to a region where they are in lower concentration, as a result of their random motion.

• It happens spontaneously in gases and liquids.
• Example: When you spray perfume, the scent molecules slowly spread out across the room.

Explanation using KPT:

1. Particles in the high concentration area are moving randomly (high kinetic energy).
2. They collide with other particles and eventually bounce/spread into the area of low concentration.
3. This continues until the particles are evenly distributed, achieving a uniform concentration.

The Effect of Relative Molecular Mass ($M_r$) on Diffusion Rate (C1.2.2 Supplement)

The rate at which a gas or liquid diffuses is not always the same. It is affected by the mass of the particles.

Rule: Substances with a lower relative molecular mass (\(M_r\)) diffuse faster.

Explanation (via KPT):

• At the same temperature, all particles have the same average kinetic energy.
• Kinetic energy (KE) is calculated using the formula: \(KE = \frac{1}{2} m v^2\), where $m$ is mass and $v$ is velocity (speed).
• If KE is constant, then a particle with a smaller mass (\(m\)) must have a higher velocity (\(v\)).

Analogy: Imagine a marathon where everyone has the same total energy budget. The lighter runners (smaller mass) can use that energy to run much faster than the heavier runners.

Common Mistake to Avoid: Diffusion rate depends on particle mass, not particle size. While related, focus on $M_r$.

Key Takeaway: Lighter gases (smaller $M_r$) move at higher speeds than heavy gases at the same temperature, so they spread out more quickly.