States of Matter - CXC Chemistry

Note to Students: This resource covers the States of Matter topics required for the CXC Chemistry syllabus (2024-2025). Use it for comprehensive review and understanding of key concepts.

Introduction to States of Matter

Matter is anything that has mass and takes up space. All matter around us exists in three main physical states: solid, liquid, and gas. Some substances can also exist in a fourth state known as plasma, though this is less common in everyday experiences and not a focus of the CXC syllabus.

SOLID LIQUID GAS

Figure 1: Particle arrangement in the three states of matter

Characteristics of the States of Matter

Solids

Liquids

Gases

Property Solid Liquid Gas
Shape Fixed Variable (container shape) Variable (container shape)
Volume Fixed Fixed Variable
Particle arrangement Regular, close-packed Irregular, close Random, far apart
Particle movement Vibration only Can slide past each other Random, high-speed motion
Intermolecular forces Strong Moderate Weak
Density High Moderate Low
Compressibility Very low Low High
Diffusion Extremely slow Slow Rapid

Kinetic Molecular Theory

The Kinetic Molecular Theory explains the behavior of matter in different states by describing the motion and energy of particles.

Key Principles:

Example: Why does a gas expand to fill its container while a solid maintains its shape?

According to Kinetic Molecular Theory, gas particles have high kinetic energy and weak intermolecular forces, allowing them to move rapidly and independently throughout the available space. In contrast, solid particles have lower kinetic energy and stronger intermolecular forces that hold them in fixed positions relative to each other, maintaining the solid's shape.

Changes of State

Matter can transition between states when energy is added or removed. These transitions are called changes of state or phase changes.

SOLID LIQUID GAS Melting Freezing Vaporization Condensation Sublimation Deposition

Figure 2: Changes of state between solids, liquids, and gases

Types of Phase Changes:

Important for CXC: Remember that phase changes that involve breaking intermolecular forces (melting, vaporization, sublimation) require energy input and are endothermic processes. Phase changes that involve forming intermolecular forces (freezing, condensation, deposition) release energy and are exothermic processes.

Energy and Phase Changes

When matter changes state, energy is either absorbed or released while the temperature remains constant. This energy is called latent heat.

Types of Latent Heat:

Heat Added Temperature (°C) 0°C 100°C Solid Melting Liquid Boiling Gas Latent Heat of Fusion Latent Heat of Vaporization

Figure 3: Heating curve showing temperature changes and phase transitions for water

A heating curve shows how the temperature of a substance changes as heat energy is added. The flat portions of the curve represent phase changes where energy is being used to break bonds rather than increase temperature.

Example calculation: How much energy is required to convert 2 kg of ice at 0°C to water at 0°C?

Given: Mass of ice = 2 kg, Latent heat of fusion of water = 334 kJ/kg

Energy required = Mass × Latent heat of fusion

Energy required = 2 kg × 334 kJ/kg = 668 kJ

Properties of Gases

Gases have unique properties that are explained by the Kinetic Molecular Theory and quantified by gas laws.

The Gas Laws:

Boyle's Law P₁V₁ = P₂V₂ Charles' Law V₁/T₁ = V₂/T₂ Gay-Lussac's Law P₁/T₁ = P₂/T₂ Ideal Gas Law PV = nRT

Figure 4: Visualization of gas laws with particles

Example: A gas occupies a volume of 250 cm³ at 27°C and 100 kPa. What volume will it occupy at 127°C and 150 kPa?

Using the Combined Gas Law: P₁V₁/T₁ = P₂V₂/T₂

We must convert temperatures to Kelvin: T₁ = 27°C + 273 = 300 K, T₂ = 127°C + 273 = 400 K

Rearranging for V₂: V₂ = (P₁V₁T₂)/(P₂T₁)

V₂ = (100 kPa × 250 cm³ × 400 K)/(150 kPa × 300 K) = 222.2 cm³

Diffusion and Brownian Motion

The movement of particles in matter leads to two important phenomena:

Diffusion

Brownian Motion

Diffusion Brownian Motion

Figure 5: Diffusion and Brownian motion

Intermolecular Forces

The forces of attraction between molecules determine many physical properties of substances, including their state at room temperature.

Types of Intermolecular Forces:

Hydrogen Bonding in Water O H H O H H O H H O H H O H H

Figure 6: Hydrogen bonding network in water molecules

Example: Why does water have a higher boiling point than expected for its molecular mass?

Water has a relatively high boiling point (100°C) compared to other molecules of similar mass due to hydrogen bonding. Each water molecule can form up to four hydrogen bonds with neighboring water molecules. These hydrogen bonds are strong intermolecular forces that require additional energy to break during the liquid-to-gas phase transition, resulting in a higher boiling point than would be expected based on molecular mass alone.

Real-Life Applications

Understanding states of matter helps explain many phenomena in everyday life and is crucial for technological applications.

Natural Phenomena:

Industrial Applications:

Glossary of Key Terms

Boiling point
The temperature at which a liquid changes to a gas at a specific pressure (usually atmospheric pressure).
Brownian motion
The random movement of particles in a fluid due to collisions with molecules of the fluid.
Condensation
The process by which a gas changes to a liquid.
Deposition
The process by which a gas changes directly to a solid without passing through the liquid state.
Diffusion
The movement of particles from a region of higher concentration to a region of lower concentration.
Evaporation
The process by which molecules at the surface of a liquid escape into the gas phase at temperatures below the boiling point.
Freezing
The process by which a liquid changes to a solid.
Hydrogen bond
A strong type of dipole-dipole attraction between a hydrogen atom bonded to a highly electronegative atom (N, O, or F) and another electronegative atom.
Intermolecular forces
The forces of attraction between molecules.
Kinetic Molecular Theory
A theory that explains the behavior of gases by considering the motion of gas particles.
Latent heat
The energy absorbed or released by a substance during a phase change without a change in temperature.
Melting point
The temperature at which a solid changes to a liquid at a specific pressure (usually atmospheric pressure).
Phase
A distinct form of matter with uniform physical properties.
Sublimation
The process by which a solid changes directly to a gas without passing through the liquid state.
Van der Waals forces
Weak forces of attraction between molecules, including dispersion forces and dipole-dipole interactions.
Vapor pressure
The pressure exerted by a vapor in equilibrium with its liquid or solid phase at a given temperature.

Self-Assessment Questions

Multiple Choice Questions:

  1. Which of the following is NOT a characteristic of gases?
    1. They are easily compressed
    2. They have a fixed shape
    3. They fill the entire container they are in
    4. They diffuse rapidly

  2. During which phase change is energy absorbed?
    1. Condensation
    2. Freezing
    3. Melting
    4. Deposition

  3. According to Boyle's Law, if the pressure on a gas is doubled while temperature remains constant, what happens to the volume?
    1. It doubles
    2. It halves
    3. It remains the same
    4. It quadruples

  4. Which intermolecular force is primarily responsible for water's unusually high boiling point?
    1. Dispersion forces
    2. Dipole-dipole forces
    3. Hydrogen bonding
    4. Ionic bonding

  5. During which phase change does a substance absorb the most energy per gram?
    1. Melting
    2. Freezing
    3. Vaporization
    4. Condensation

Short Answer Questions:

  1. Explain why diffusion occurs faster in gases than in liquids.

  2. A sample of gas occupies 250 cm³ at 25°C and 1 atmosphere pressure. Calculate its volume at 50°C and 2 atmospheres pressure.

  3. Why does the temperature remain constant during a phase change even though energy is being added or removed?

  4. Describe the particle arrangement and movement in each state of matter.

  5. Explain why water expands when it freezes, unlike most substances.

Answers:

Multiple Choice:

  1. b - Gases have variable shape, not fixed shape
  2. c - Melting is an endothermic process (absorbs energy)
  3. b - Volume halves (inverse relationship)
  4. c - Hydrogen bonding
  5. c - Vaporization (breaking all intermolecular bonds requires more energy than melting)

Short Answer:

  1. Diffusion occurs faster in gases than in liquids because:

    • Gas particles are further apart than liquid particles, allowing for freer movement
    • Gas particles move at higher speeds than liquid particles due to weaker intermolecular forces
    • There are fewer collisions between gas particles that would slow down diffusion
    • Liquid particles are held more closely together by stronger intermolecular forces, restricting their movement

  2. Using the Combined Gas Law: P₁V₁/T₁ = P₂V₂/T₂

    First, convert temperatures to Kelvin:

    T₁ = 25°C + 273 = 298 K

    T₂ = 50°C + 273 = 323 K

    Rearranging to solve for V₂:

    V₂ = (P₁V₁T₂)/(P₂T₁)

    V₂ = (1 atm × 250 cm³ × 323 K)/(2 atm × 298 K)

    V₂ = 250 × 323 / (2 × 298) = 135.6 cm³

  3. The temperature remains constant during a phase change because the energy being added or removed is used to change the arrangement of the particles rather than increase their kinetic energy. During a phase change, the energy goes into breaking or forming intermolecular bonds between particles. For example, when ice melts, the energy absorbed goes into breaking some of the hydrogen bonds that hold the water molecules in a rigid lattice, allowing them to move more freely in the liquid state. The average kinetic energy of the particles (which determines temperature) remains the same until the phase change is complete.

  4. Solid: Particles are arranged in a regular, fixed pattern and are closely packed together. They can only vibrate about their fixed positions and cannot move past each other. The strong intermolecular forces hold the particles rigidly in place.

    Liquid: Particles are still close together but not in a regular pattern. They can move past each other and slide around, giving liquids their ability to flow. Intermolecular forces are strong enough to keep particles close but not fixed in position.

    Gas: Particles are far apart and arranged randomly. They move freely at high speeds in all directions, colliding with each other and the walls of their container. Intermolecular forces are very weak, allowing particles to move independently.

  5. Water expands when it freezes because of its unique hydrogen bonding structure. In liquid water, molecules are relatively close together with hydrogen bonds constantly breaking and reforming. As water cools and approaches freezing, the decreased kinetic energy allows hydrogen bonds to form a more stable, open hexagonal crystal lattice structure in ice. This hexagonal arrangement actually takes up more space than the more random arrangement in liquid water, causing ice to be less dense than liquid water. The hydrogen bonds force water molecules to maintain certain distances and angles from each other, creating open spaces within the ice crystal. This is why ice floats on water, unlike most solids which sink in their liquid form.

Conclusion

Understanding the states of matter and their transitions is fundamental to chemistry and explains countless phenomena in our daily lives and in industrial processes. The behavior of matter in its different states is determined by the arrangement and energy of particles, as well as the intermolecular forces between them. For CXC Chemistry examinations, ensure you are comfortable with all the concepts covered in this resource, particularly:

Remember to practice applying these concepts to real-world situations and problems, as this is often a focus in CXC examinations.