Matter Revision Notes

Understanding Changes of State

Changes of state are physical processes, not chemical reactions. This means that the chemical identity of the substance does not change. For instance, water (H₂O) remains water whether it is ice (solid), liquid water, or steam (gas). The particles themselves are preserved throughout the change. These processes include melting (solid to liquid), freezing (liquid to solid), evaporation (liquid to gas), condensation (gas to liquid), and sublimation (solid to gas or gas to solid).

The Principle of Mass Conservation

The law of conservation of mass states that mass cannot be created or destroyed in a closed system. During changes of state, the total number and type of particles (atoms or molecules) of a substance remain unchanged. They are simply rearranged or spread out. Therefore, the total mass of the substance before and after a change of state will always be the same, provided no part of the substance is allowed to escape or be added to the system.

Practical Demonstration: Closed Systems

To observe mass conservation experimentally, a closed system is crucial. For example, if you melt ice in a sealed container, the total mass of the container plus its contents will not change. If water evaporates in a sealed flask, the mass of the flask and its contents (liquid water and water vapour) will remain constant. This is because all the particles, even those that have changed state, are still contained within the system and contribute to its overall mass.

Apparent Mass Changes in Open Systems

Sometimes, it might seem like mass is not conserved, especially in an open system. For example, if liquid water boils in an open beaker, the mass of the beaker and its remaining water will decrease because water vapour (gas) escapes into the atmosphere. However, the mass of the water itself is still conserved; the escaped water vapour simply moved out of the measured system. If you could capture and weigh all the escaped water vapour, you would find that the total initial mass is equal to the mass of the remaining liquid plus the mass of the collected vapour.

Understanding Density

Density is defined as mass per unit volume. It indicates how tightly packed the particles (atoms or molecules) are within a substance. Substances with particles more closely packed into a given volume will have a higher density. Understanding density helps explain many physical properties and phenomena, such as why some objects float and others sink.

Solids: High Density

In solids, particles are arranged in a fixed, regular pattern and are very closely packed together. They vibrate about fixed positions but do not move past each other. This tight, ordered arrangement means that there are many particles in a small volume, giving solids generally the highest densities among the three states of matter. For example, a block of iron is much denser than liquid iron or iron vapour.

Liquids: Intermediate Density

In liquids, particles are closely packed but are arranged randomly and can move and slide past one another. This allows liquids to flow and take the shape of their container. Although the particles are still close together, the random arrangement and ability to move mean they are typically slightly less densely packed than in solids, resulting in intermediate densities. An important exception is water, where ice (solid) is less dense than liquid water due to its crystal structure.

Gases: Low Density

In gases, particles are widely spread out, move randomly and rapidly, and have large distances between them. There are very few particles in a given volume compared to solids or liquids. This significant separation of particles leads to gases having very low densities. Gases can be easily compressed because there are large empty spaces between their particles.

Defining Specific Heat Capacity (c)

Specific heat capacity (c) is a measure of how much thermal energy is needed to change the temperature of a material. Specifically, it is the amount of energy required to raise the temperature of 1 kilogram (kg) of a substance by 1 degree Celsius (°C) or 1 Kelvin (K). Its units are joules per kilogram per degree Celsius (J/kg°C) or joules per kilogram per Kelvin (J/kgK). Materials with a high specific heat capacity require more energy to change their temperature, while materials with a low specific heat capacity require less energy for the same temperature change.

Defining Specific Latent Heat (L)

Specific latent heat (L) is the amount of thermal energy absorbed or released when 1 kilogram (kg) of a substance undergoes a change of state (e.g., melts, boils, freezes, condenses) without any change in its temperature. The energy supplied is used to overcome the intermolecular forces during the phase transition, not to increase the kinetic energy of the particles. Its units are joules per kilogram (J/kg). There are two main types: specific latent heat of fusion (for melting/freezing) and specific latent heat of vaporisation (for boiling/condensing).

Key Distinction Between c and L

The fundamental difference between specific heat capacity and specific latent heat lies in their effect on temperature. Specific heat capacity deals with the energy required to change the *temperature* of a substance while it remains in the *same state*. For example, heating liquid water from 20°C to 80°C. Specific latent heat, on the other hand, deals with the energy required to change the *state* of a substance (e.g., from solid to liquid or liquid to gas) *without any change in temperature*. For example, melting ice at 0°C to form water at 0°C, or boiling water at 100°C to form steam at 100°C.

Practical Examples

Water has a very high specific heat capacity. This is why it is used as a coolant in car engines and in central heating systems; it can absorb or release a large amount of energy with a relatively small temperature change. For specific latent heat, consider boiling water: you can keep adding heat to water at 100°C, but its temperature will not rise above 100°C; instead, it will turn into steam. This energy is the specific latent heat of vaporisation. Similarly, ice can absorb a significant amount of energy (specific latent heat of fusion) to melt at 0°C without its temperature increasing, making it an effective cooling agent.

Kinetic Model of Gases

Gases consist of a large number of molecules (or atoms, for noble gases) that are in continuous, random motion. These molecules move at high speeds, colliding with each other and with the walls of their container. The collisions with the container walls exert a force on the walls, and this force distributed over the area of the walls is what we define as the gas pressure.

Temperature and Molecular Motion

Temperature is a direct measure of the average kinetic energy of the gas molecules. When a gas is heated, its temperature increases, which means the average kinetic energy of its molecules increases. Consequently, the molecules move faster on average. If a gas is cooled, its temperature decreases, the average kinetic energy of its molecules decreases, and they move slower.

Temperature, Pressure and Volume Relationship

If the temperature of a gas increases (and its volume is kept constant), the faster-moving molecules collide with the container walls more frequently and with greater force. This leads to an increase in the pressure exerted by the gas. Conversely, a decrease in temperature, with constant volume, causes molecules to move slower, resulting in less frequent and less forceful collisions, and thus a decrease in pressure. If the volume of the container decreases, the molecules have a smaller space to move in, leading to more frequent collisions with the walls, which increases the pressure, even if the temperature stays the same.