Particle model Revision Notes

In solids, particles (atoms, molecules, or ions) are tightly packed in a regular, fixed arrangement. This close packing means particles are very close together with almost no space between them.

The particles are held in place by strong forces of attraction, which keep them fixed in position. Because of this, solids have a definite shape and volume.

Particles in solids do not move from their fixed positions but vibrate about these positions. The vibration increases with temperature but does not allow the particles to move freely.

• Particles are closely packed in a fixed, regular pattern.
• Very little space between particles.
• Strong forces hold particles in fixed positions.
• Particles vibrate about fixed points but do not move around.
• Solids have definite shape and volume.

In liquids, particles are still close together but are not arranged in a fixed pattern. The forces of attraction between particles are weaker than in solids, allowing particles to move around each other.

There is some space between particles, more than in solids but less than in gases. This allows liquids to flow and take the shape of their container while maintaining a definite volume.

Particles in liquids move randomly and slide past each other, which explains why liquids can flow but are not easily compressed.

• Particles are close but arranged irregularly.
• Some space between particles, more than solids.
• Forces of attraction are weaker than in solids.
• Particles move randomly and slide past each other.
• Liquids have definite volume but no fixed shape.

In gases, particles are far apart with large spaces between them compared to solids and liquids. The forces of attraction between particles are very weak or negligible.

Particles move very rapidly and randomly in all directions. Because of this, gases have neither a definite shape nor a definite volume and will expand to fill any container.

The large spaces between particles mean gases are easily compressed compared to solids and liquids.

• Particles are widely spaced with large gaps.
• Forces of attraction are negligible.
• Particles move rapidly and randomly in all directions.
• Gases have no definite shape or volume.
• Gases expand to fill their container and are compressible.

The movement of particles in each state is related to their kinetic energy. In solids, particles have the least kinetic energy and only vibrate in place.

In liquids, particles have more kinetic energy, allowing them to move around each other but still remain close.

In gases, particles have the highest kinetic energy, moving freely and quickly in all directions.

• Particle movement increases from solid to liquid to gas.
• Kinetic energy of particles is lowest in solids and highest in gases.
• Particle movement explains the physical properties of each state.

To illustrate the states of matter, draw simple particle diagrams following these instructions:

For solids: draw particles as small circles arranged closely in a regular grid pattern with no gaps. Indicate vibration by small double-headed arrows around each particle.

For liquids: draw particles as small circles close together but arranged irregularly with small gaps. Show movement by arrows between particles indicating sliding past each other randomly in different directions but close proximity remains constant due to attraction forces. Avoid a fixed pattern or grid arrangement in the drawing to reflect fluidity and disorder compared to solids. The particles should be touching or nearly touching but not in straight lines or rows as in solids. For gases: draw particles as small circles spaced far apart randomly scattered across the space. Use arrows of varying lengths and directions to show rapid, random movement in all directions. The large gaps between particles should be clearly visible to indicate low density and high compressibility.

• Solid diagram: closely packed, regular pattern, vibration arrows.
• Liquid diagram: close but irregular arrangement, sliding movement arrows.
• Gas diagram: widely spaced, random positions, long movement arrows.
• Use consistent particle size and clear arrows to show movement.
• Label diagrams with state names and key features.

Temperature is a measure of the average kinetic energy of the particles in a substance. Kinetic energy is the energy that particles have due to their motion. As temperature increases, the particles move faster, meaning their kinetic energy increases.

Particles in solids, liquids, and gases are always moving, but the speed and type of motion depend on the temperature and state of matter. When a substance is heated, its particles gain kinetic energy and move more rapidly.

The relationship between temperature and average kinetic energy can be expressed as: \n\n\tAverage kinetic energy ∝ Temperature (in kelvin)\n\nThis means that if the temperature doubles (in kelvin), the average kinetic energy of the particles also doubles.

• Temperature is directly proportional to the average kinetic energy of particles.
• Higher temperature means faster particle motion.
• Particle motion includes vibrations, rotations, and translations depending on the state of matter.

Absolute zero is the lowest possible temperature, defined as 0 kelvin (K), which is equivalent to -273 °C. At this temperature, the particles have minimal kinetic energy and theoretically stop moving completely.

In reality, particles cannot be completely at rest due to quantum mechanical effects, but absolute zero represents the point where particle motion is as low as physically possible.

Absolute zero is important because it sets the zero point on the Kelvin temperature scale, which is used in scientific measurements of temperature.

• Absolute zero = 0 K = -273 °C.
• At absolute zero, particle kinetic energy is minimal.
• Particles theoretically have no motion at absolute zero.

Temperature is measured in degrees Celsius (°C) or kelvin (K). The kelvin scale is an absolute scale starting at absolute zero.

The relationship between Celsius and kelvin is: \n\n\tT(K) = T(°C) + 273\n\nThis means that 0 °C corresponds to 273 K, and absolute zero is 0 K or -273 °C.

Using kelvin is important when relating temperature to kinetic energy because it starts at the point where particle motion is minimal.

• Kelvin scale starts at absolute zero (0 K).
• T(K) = T(°C) + 273.
• Kelvin is used in physics to relate temperature to kinetic energy.

Thermal energy is the total kinetic energy of all particles in a substance. When a substance is heated, thermal energy increases because particles move faster.

The average kinetic energy of particles is related to temperature, but thermal energy depends on both kinetic energy and the number of particles.

This explains why larger amounts of a substance contain more thermal energy at the same temperature.

• Thermal energy = total kinetic energy of particles.
• Temperature measures average kinetic energy, not total energy.
• More particles at the same temperature means more thermal energy.

To draw a particle diagram showing the effect of temperature on particle motion:

1. Draw three boxes labelled 'Low Temperature', 'Room Temperature', and 'High Temperature'.

2. Inside each box, draw several small circles representing particles. At low temperature, draw particles close together with small arrows indicating slow motion (short arrows). At room temperature, draw particles slightly further apart with medium-length arrows. At high temperature, draw particles more spread out with long arrows showing fast motion.

• Particle speed is shown by arrow length.
• Particle spacing may increase with temperature in gases.
• Diagram illustrates increasing kinetic energy with temperature.

Gas pressure is defined as the force exerted per unit area on the walls of a container by gas particles. It is measured in pascals (Pa), where 1 Pa = 1 newton per square metre (N/m²).

Gas particles move randomly and constantly in all directions. When these particles collide with the container walls, they exert a force on the surface. The cumulative effect of many such collisions per unit area results in gas pressure.

• Gas pressure = force exerted by particle collisions per unit area (Pa = N/m²).
• Particles move randomly and collide with container walls, exerting force.

Gas particles are in constant, random motion with a range of speeds. Each collision with the container wall transfers momentum, exerting a force. The total pressure depends on the number of collisions per second and the average force per collision.

The kinetic theory of gases explains that the pressure is caused by these collisions. Faster particles collide more frequently and with greater force, increasing pressure.

• Pressure arises from particle collisions transferring momentum to container walls.
• More frequent or more forceful collisions increase pressure.

Temperature is a measure of the average kinetic energy of gas particles. Increasing temperature increases particle speed.

Faster particles collide more often and with greater force against the container walls, increasing the pressure if volume is constant.

This relationship is described by Gay-Lussac's law: at constant volume, pressure is directly proportional to temperature (in kelvin).

• Temperature increase → higher particle speed → higher pressure (if volume constant).
• Pressure ∝ Temperature (K) at constant volume (Gay-Lussac's law).

Reducing the volume of a gas (while keeping temperature constant) means particles have less space to move.

Particles collide more frequently with the container walls, increasing the pressure.

This relationship is described by Boyle's law: at constant temperature, pressure is inversely proportional to volume.

• Decreasing volume → more frequent collisions → higher pressure (if temperature constant).
• Pressure ∝ 1/Volume at constant temperature (Boyle's law).

Gas pressure depends on particle motion and collisions with surfaces.

Pressure increases with increasing temperature (at constant volume) due to faster particle speeds.

Pressure increases with decreasing volume (at constant temperature) due to more frequent collisions with container walls.

• Pressure (P) measured in pascals (Pa).
• Temperature (T) measured in kelvin (K).
• Volume (V) measured in cubic metres (m³).

A suspension is a mixture in which tiny solid particles are dispersed throughout a liquid or gas but are not dissolved. These particles are large enough to be seen under a microscope but small enough to remain suspended for some time before settling.

The particles in a suspension move in a random, erratic manner. This movement is not caused by any external force but results from collisions with the much smaller, fast-moving molecules of the fluid in which they are suspended.

This random movement of particles in a suspension is observable under a microscope and is known as Brownian motion. It provides direct visual evidence that particles in matter are constantly moving.

• Suspension: solid particles dispersed in a fluid, not dissolved.
• Particles in suspension exhibit random, erratic movement.
• Brownian motion is the observed random movement of suspended particles.

The kinetic particle theory states that all matter is made up of tiny particles (atoms or molecules) that are in constant motion. The theory explains the properties of solids, liquids, and gases based on particle arrangement and movement.

In liquids and gases, particles move randomly with varying speeds and directions. This motion is continuous and results from the particles’ kinetic energy.

The random movement of particles in a suspension supports this theory by showing that particles are not stationary but constantly moving due to collisions with fluid molecules.

• Matter consists of tiny particles in constant motion.
• Particle motion explains states of matter and their properties.
• Random movement in suspensions confirms particles are always moving.

Brownian motion is the continuous, random movement of microscopic particles suspended in a fluid, caused by collisions with the fluid’s molecules.

This phenomenon was first observed by Robert Brown in 1827 when he looked at pollen grains in water under a microscope.

The irregular, jittery motion of suspended particles cannot be explained by external currents or convection but is due to the kinetic energy of the fluid’s molecules constantly bombarding the particles from all sides.

• Brownian motion is caused by collisions with fast-moving fluid molecules.
• It provides visual proof of the kinetic particle theory.
• Observed under a microscope as random, erratic particle movement.

To observe Brownian motion, a simple experiment can be performed using a microscope and a suspension such as fine smoke particles in air or pollen grains in water.

Apparatus: microscope, slide, cover slip, suspension sample (e.g., pollen in water).

Method: Place a drop of suspension on the slide, cover it with a cover slip, and observe under the microscope at high magnification. The tiny particles will be seen moving randomly in all directions without any apparent pattern or external cause.

• Use microscope to observe suspended particles.
• Particles move randomly due to collisions with fluid molecules.
• Observation supports kinetic particle theory.

The random movement of particles in suspensions confirms that particles have kinetic energy and are in constant motion, a fundamental assumption of the kinetic particle theory.

This understanding helps explain diffusion, gas pressure, and temperature effects in matter.

It also supports the idea that temperature is related to the average kinetic energy of particles, as faster-moving molecules cause more vigorous collisions and more noticeable Brownian motion.

• Random particle motion confirms constant kinetic energy of particles.
• Supports explanations of diffusion and gas pressure.
• Links particle motion to temperature and energy.

Brownian motion is the continuous, irregular, and random movement of microscopic particles suspended in a fluid (liquid or gas). These particles are typically visible under a microscope and include pollen grains, dust, or smoke particles.

When observed under a microscope, these particles do not remain still but move in a zigzag or jittery path. This movement is not caused by currents or convection but by collisions with the molecules of the surrounding fluid, which are themselves in constant motion due to thermal energy.

• Microscopic particles suspended in fluids exhibit random motion called Brownian motion.
• Movement is visible under a microscope as a jittery, zigzag path.
• First observed by Robert Brown in 1827 while looking at pollen grains in water.

Fluids consist of a vast number of molecules moving randomly at high speeds. These molecules collide with suspended microscopic particles from all directions.

Because the collisions are random and uneven in frequency and force, the suspended particles receive unequal pushes, causing them to move unpredictably.

This continuous bombardment results in the erratic motion characteristic of Brownian motion, demonstrating that fluid molecules are in constant motion even though they are too small to be seen.

• Fluid molecules move randomly and collide with suspended particles.
• Unequal collisions cause random changes in particle velocity and direction.
• Brownian motion is evidence of molecular motion in fluids.

The kinetic theory states that matter is made of tiny particles in constant motion. Brownian motion provides direct visual evidence supporting this theory.

The random movement of suspended particles shows that fluid molecules are moving and transferring momentum through collisions.

This supports the idea that temperature relates to the average kinetic energy of particles: higher temperature means faster molecular motion, increasing the intensity of Brownian motion.

• Brownian motion confirms that molecules in fluids are constantly moving.
• It supports the kinetic theory of matter and the concept of thermal energy.
• The intensity of Brownian motion increases with temperature due to faster molecular speeds.

The path of a particle undergoing Brownian motion is unpredictable and changes direction frequently.

The motion is three-dimensional and results from countless collisions occurring at different angles and speeds.

Because the collisions are random, the particle’s displacement over time follows a statistical pattern rather than a straight line or predictable curve.

• Brownian motion is random and erratic in direction and speed.
• Particles move in a zigzag, three-dimensional path.
• Displacement over time is statistical, not deterministic.

Brownian motion can be observed by placing a small amount of smoke or pollen in water and viewing it under a microscope.

This phenomenon helps scientists understand molecular motion and diffusion processes in fluids.

It also provides evidence for the existence of atoms and molecules, which were once theoretical concepts.

• Observed using microscopes with suspensions of pollen or smoke particles.
• Demonstrates molecular motion and supports atomic theory.
• Important in understanding diffusion and thermal processes in fluids.

Particles in matter include atoms, molecules, ions, and electrons. These particles exert forces on each other, which can be attractive or repulsive depending on their separation distance.

In solids and liquids, strong intermolecular forces hold particles close together. These forces include ionic bonds, covalent bonds, metallic bonds, and van der Waals forces.

In gases, particles are much further apart, so the forces between them are negligible except during collisions. This weak interaction allows gas particles to move freely and independently.

• Interparticle forces determine how closely particles are held together.
• Strong forces in solids and liquids maintain particle proximity.
• Negligible forces in gases allow free particle movement.

In solids, particles are tightly packed in a fixed, regular arrangement, often forming a lattice structure.

Particles vibrate about fixed positions but do not move freely. This vibration increases with temperature but does not change the overall shape or volume.

The strong forces between particles prevent them from moving apart or sliding past each other, giving solids a definite shape and volume.

• Particles closely packed in fixed positions.
• Vibrate about fixed points; no translational movement.
• Strong forces maintain fixed shape and volume.

Particles in liquids are close together but not in a fixed arrangement, allowing them to slide past each other.

This mobility enables liquids to flow and take the shape of their container while maintaining a fixed volume.

Interparticle forces are weaker than in solids but still significant enough to keep particles close, preventing large expansion or compression under normal conditions.

• Particles close but disordered arrangement.
• Particles move past each other, allowing flow.
• Fixed volume but shape changes with container.

Gas particles are widely spaced with negligible forces between them except during collisions.

They move rapidly and randomly in all directions, filling any container uniformly.

Because particles are free to move apart or come together, gases have neither fixed shape nor fixed volume and are easily compressible and expandable.

• Particles far apart with negligible forces.
• Random, rapid motion in all directions.
• No fixed shape or volume; gases fill containers.

The differences in forces and particle movement explain the macroscopic properties of matter states: solids have fixed shape and volume; liquids have fixed volume but variable shape; gases have neither fixed shape nor volume.

Temperature affects particle vibration and movement but does not change the fundamental forces unless a phase change occurs.

Understanding these microscopic behaviours is essential for explaining phenomena such as diffusion, pressure, and phase transitions.

• Strong forces + fixed positions = solid properties.
• Moderate forces + particle mobility = liquid properties.
• Weak forces + free movement = gas properties.
• Particle behaviour underpins macroscopic matter properties.

Gas pressure is the force exerted by gas particles when they collide with the walls of their container, distributed over the area of the surface. It is quantitatively defined as the force applied per unit area, with the SI unit of pressure being the pascal (Pa), where 1 Pa = 1 N/m².

Mathematically, pressure (P) is given by: P = F / A, where F is the total force exerted perpendicular to the surface and A is the area of that surface.

• Pressure (P) = Force (F) ÷ Area (A)
• Unit of pressure: pascal (Pa) = newton per square metre (N/m²)

Gas consists of a large number of tiny particles (atoms or molecules) moving randomly at high speeds. These particles constantly collide with each other and with the walls of their container.

Each collision with a container wall exerts a tiny force on the surface. The cumulative effect of countless collisions per second over the entire surface area produces the measurable gas pressure.

The force from each collision arises because particles change momentum when they bounce off the surface, exerting an impulse force during the collision.

• Gas particles move randomly and collide with container walls.
• Each collision exerts a force on the surface.
• Pressure results from many such collisions per unit area per second.

The pressure of a gas depends on how often particles collide with the container walls and how forceful these collisions are.

Increasing the temperature of the gas increases the average speed of particles, causing more frequent and more forceful collisions, thus increasing pressure if volume is constant.

Decreasing the volume of the container reduces the space available, so particles collide more often with the walls, increasing pressure if temperature is constant.

• Higher temperature → faster particles → higher pressure (at constant volume)
• Smaller volume → more frequent collisions → higher pressure (at constant temperature)

To calculate gas pressure, measure the total force exerted by the gas particles on a surface and divide by the area of that surface.

Example: If a force of 10 N is exerted uniformly on a surface area of 2 m², the pressure is P = 10 N ÷ 2 m² = 5 Pa.

Common errors include forgetting to convert units or mixing up force and pressure. Always check units carefully.

• P = F / A, where P in pascals, F in newtons, A in square metres
• Example calculation: F = 10 N, A = 2 m² → P = 5 Pa

A simple diagram to illustrate gas pressure shows a container with gas particles inside moving randomly and colliding with the walls.

Draw arrows representing particle velocities pointing in various directions.

At the walls, show arrows indicating forces exerted by particles on the surface during collisions, perpendicular to the surface area element.

• Draw container walls as straight lines.
• Draw multiple small circles inside representing gas particles.
• Use arrows to show random particle motion and collision forces on walls.
• Label force (F), area (A), and pressure (P = F/A).

Microscopic particles are tiny particles that are too small to be seen with the naked eye. They include atoms, molecules, and other small particles such as tiny solid particles suspended in a fluid.

An atom is the smallest unit of a chemical element that retains its chemical properties. Atoms consist of a nucleus surrounded by electrons.

A molecule is formed when two or more atoms chemically bond together. Molecules can be made of the same type of atoms (e.g., O₂) or different types (e.g., H₂O).

• Microscopic particles include atoms, molecules, and other small particles.
• Atoms are the basic units of elements.
• Molecules are groups of atoms bonded together.

Microscopic particles are in constant motion due to their kinetic energy. In gases and liquids, molecules move rapidly and randomly.

When fast-moving molecules collide with microscopic particles suspended in a fluid, they transfer momentum to these particles. This causes the microscopic particles to move in a random, jittery manner.

These collisions are elastic, meaning kinetic energy is conserved during the interaction, although it is transferred between particles.

• Molecules move rapidly and randomly in fluids.
• Collisions between molecules and microscopic particles transfer momentum.
• These collisions cause random motion of microscopic particles.

The collisions of fast-moving molecules with microscopic particles cause the particles to move unpredictably. This phenomenon is observed as Brownian motion.

Brownian motion provides evidence for the existence of atoms and molecules, as it demonstrates that microscopic particles are constantly bombarded by smaller, fast-moving molecules.

The size and mass of microscopic particles affect how much they are influenced by these collisions; smaller particles move more noticeably.

• Collisions cause Brownian motion of microscopic particles.
• Brownian motion supports the kinetic theory of matter.
• Smaller particles are more affected by molecular collisions.

Brownian motion can be observed under a microscope by viewing tiny particles suspended in a liquid or gas.

This random motion is caused by countless collisions with molecules moving at high speeds due to thermal energy.

Understanding these interactions helps explain diffusion, gas pressure, and other physical phenomena at the microscopic level.

• Brownian motion is observable under a microscope.
• It results from molecular collisions due to thermal energy.
• These interactions explain diffusion and gas pressure.

Microscopic particles include atoms, molecules, and other tiny particles.

Atoms are the smallest units of elements; molecules are bonded groups of atoms.

Fast-moving molecules collide with microscopic particles, causing random motion known as Brownian motion, which supports the kinetic theory of matter.

• Microscopic particles are influenced by molecular collisions.
• Atoms and molecules are specific types of microscopic particles.
• Collisions cause observable random motion of microscopic particles.