Kinetic Particle Model: States and Changes

Kinetic Particle Model of Matter: States of Matter

• Solids possess a definite shape and volume; their particles are tightly packed and cannot flow or be compressed.
• Liquids have a definite volume but take the shape of their container; their particles can flow but are not easily compressed.
• Gases lack a definite shape or volume, filling the container they occupy; their particles can flow and are highly compressible.

Changes of State

• Altering a substance's state doesn't change the number of molecules or mass, only its energy.
• State changes are physical and reversible.
• Melting is the process of a solid turning into a liquid.
• Freezing is when a liquid transforms into a solid.
• Boiling occurs when a liquid becomes a gas, also known as evaporation.
• Condensation is the process of a gas turning into a liquid.

Molecular Matter: Arrangement and Motion of Particles

• In solids, molecules are very close together in a regular pattern and vibrate around fixed positions.
• In liquids, molecules are close but not in a regular pattern and can slide past each other.
• In gases, molecules are widely separated, move randomly at high speeds, and are about 10 times further apart compared to solids or liquids.
• Solids exhibit high density with molecules in a regular pattern, liquids have medium density with randomly arranged molecules, and gases have low density with randomly arranged molecules.

Intermolecular Forces and Motion of Particles

• Intermolecular forces affect a substance's state by influencing the distances and motion of particles, affecting its ability to change shape, volume, and flow.
• Matter consists of atoms, molecules, ions, and electrons.
• Solid molecules are held by strong intermolecular forces and only vibrate in place, giving them a rigid shape and fixed volume.
• Liquid molecules have enough energy to overcome some intermolecular forces, remaining close but able to move around, allowing liquids to change shape and flow.
• Gas molecules possess more energy and move randomly at high speeds, overcoming the forces that hold them close, allowing gases to be compressed, expanded, and flow freely.

Temperature and Energy of Particles

• Gas pressure depends on temperature; an increase results in higher particle kinetic energy.
• Lowering gas temperature reduces pressure on the container.
• In 1848, Lord Kelvin theorized absolute zero, the temperature at which gas particles exert no pressure and stop moving; this is equal to -273°C.
• At absolute zero, particles have no net movement, making a lower temperature impossible.
• The unit Kelvin (K) is used to measure absolute temperature, without degree notation (°).

Temperature & Pressure: Motion of Particles in a Gas

• Gas molecules move randomly at high speeds, changing direction upon collision with container walls or other molecules.
• Gas pressure results from collisions with the container's surface (walls).
• Gases fill their container, and pressure is defined as force per unit area.
• Gases at higher pressure collide more frequently with container walls and exert greater force per unit area.

Brownian Motion

• The Kinetic Theory of Matter: all matter comprises tiny particles.
• Robert Brown observed pollen grains' random motion in water under a microscope.
• Brownian motion: the erratic movement of particles in a liquid or gas caused by numerous collisions can be observed as they move under a microscope.
• Light, fast-moving molecules collide with larger particles, giving them a small nudge.

Gases & Absolute Temperature Formulas

• The Kelvin scale starts at absolute zero.
• 0 K equals -273 °C.
• Kelvin scale: an increase of 1 K is the same as an increase of 1 °C.
• It isn't possible to have a temperature lower than 0 K.
• Temperature in Kelvin wil NEVER be a negative value.
• Formulas to convert between Celsius and Kelvin: °C = K - 273, K = °C + 273

The Gas Laws: Pressure & Volume (Constant Temperature)

• If temperature remains constant: compressing a gas decreases the volume which then increases the pressure, and expanding a gas increases the volume which then decreases the pressure.
• A change in pressure leads to a change in volume.
• Boyle's Law: at constant temperature, pressure and volume are inversely proportional
• Where: P1V1 = P2V2
  • P1 = initial pressure (Pa)
  • P2 = final pressure (Pa)
  • V1 = initial volume (m^3)
  • V2 = final volume (m^3)

Thermal Properties & Temperature: Thermal Expansion

• When heated, materials expands because molecules move faster, causing them to push each other apart.
• Thermal expansion occurs in solids, liquids, and gases.
• With increased temperature(at constant pressure), solids expand the least, gases expand the most, and liquids' expansion falls between solids and gases.
• Thermal expansion terms: molecules don't expand, but the space in between them does.
• When solids, liquids, and gases heat up: solids expand slightly because the strong bonds that hold the molecules together, liquids expand more because of weaker bonds between the molecules, and gases expand significantly because there is an absence of bonds holding the molecules together.

Uses & Consequences of Thermal Expansion: Applications and Consequences

• Thermal expansion has applications and can have undesirable consequences.
• Applications consist of thermometers which rely on liquid expansion to measure temperature.
• Bimetallic strip is used in temperature-activated switches and consists of two metals with different rates of expansion.
• Expansion in solid materials causes them to buckle if they get too hot. This could include metal railway tracks, road surfaces, and bridges.
• Structures prone to buckling have built-in gaps to allow space for expansion without damage.

Specific Heat Capacity: Internal Energy

• An object's temperature increase raises internal energy.
• Internal energy arises from an increase in the average speed of particles.
• The amount of energy required to heat a liquid depends on the mass of the substance, the type of material, and how much thermal energy is transferred into the system.
• Specific heat capacity formula: the amount of energy required to raise the temperature of 1 kg of the substance by 1 °C
• Internal energy; the total energy stored inside a system by the particles that make up the system due to their motion and positions.
• Motion of particles affects how kinetic energy is stored
• Positions of the particles relative to each other affects their potential energy
• Together, these two points make up the internal energy of the system.

Calculating Specific Heat Capacity and Values

• Specific heat capacity (c) measures a substance's ability to store thermal energy.
• The specific heat capacity of a substance is defined as: the amount of energy required to raise the temperature of 1 kg of the substance by 1 °C.
• Substances with low specific heat capacity heat up and cool down quickly; substances with a high specific heat capacity heat up and cool down slowly.
• To calculate the specific heat capacity: ΔQ = mcΔT
  • ΔQ = change in thermal energy, in joules (J)
  • m = mass, in kilograms (kg)
  • c = specific heat capacity, in joules per kilogram per degree Celsius (J/kg °C)
  • ΔT = change in temperature, in degrees Celsius (°C)
• Heat transfer is not as efficient in insulators, so they're used to reduce heat loss.

Investigating Specific Heat Capacity: Aims and System Variables

• Aim: determining specific heat capacity by linking energy store decrease to temperature/thermal energy increase.
• Independent variable: time (t)
• Dependent variable: temperature (θ)
• Control variables:
  • Material of block
  • Current supplied (I)
  • Potential difference supplied (V)
• Formula to calculate the thermal energy supplied to a block: E = IVt, where:
  • ΔE = change in thermal energy, in joules (J)
  • m = mass, in kilograms (kg)
  • c = specific heat capacity, in joules per kilogram per degree Celsius (J/kg °C)
  • ∆θ = change in temperature, in degrees Celsius (°C)

Investigating Specific Heat Capacity: Resolution/Possible Errors

• Resolution of equipment: thermometer = 1 °C, and stopwatch = 0.01 s
• Make sure the voltmeter and ammeter are initially set to zero, to avoid zero error
• Random errors: make sure the block is fully insulated so not all heat energy from the heater is transferred to the block, and to reduce the surroundings
• The measured value of specific heat capacity is likely higher than it actually is
• You can use a joulemeter and read value at eye level to avoid parallax error

Melting & Boiling

• The melting and boiling points of pure water are known as fixed points.
• Ice melts at 0 °C
• Pure water boils at 100 °C.
• These are the accepted values for water at atmospheric pressure

Understanding States and Processes

• At melting/freezing or boiling/condensing, the state is changing.
• During any form of state change, a substance doesn't change temperature even with external factors.
• When liquid water is heated, temperature rises until water boils.
• At boiling point, liquid water does not get hotter, even with added thermal energy.
• Additional thermal energy is used to overcome intermolecular forces.
• This is called evaporation or vaporisation, and the water becomes gas.
• The process is reversible: when energy moves away, gas turns back into liquid through condensation.
• When solid water is heated, the ice melts: at the melting point, even with added thermal energy, solid water does not get warmer.
• In melting, thermal energy is used to overcome intermolecular forces and as the forces are decreased, the solid becomes liquid.
• A liquid turns back into a solid through freezing

Graphing Heating and Cooling

• Heating and cooling graphs illustrate temperature changes and state changes as energy is transferred into or out of a substance: The particles lose kinetic energy, in the particles get closer together until they only have enough energy to either flow-over one and another or vibrate
• Heating: energy transferred to increase Kinetic energy,
• Cooling: energy transferred from the molecule to decrease their kinetic energy.
• Condensation: gas loses heat, reducing kinetic energy, leading to closer particles with energy now sufficient to flow over each other, transitioning into a liquid.
• Solidification: when a liquid loses heat, its particles have reduced kinetic energy, bringing them closer, with energy limited to vibration at a fixed position, resulting in a transition to the solid phase.
• A change in temperature indicates heat transfer.

Evaporation

• Evaporation is the change in state of a liquid to gas
• Evaporation occurs: at any temperature, from the surface of the liquid.
• Molecules in a liquid possess various energies; this average energy is related to the temperature of the liquid.
• Evaporation is when more energetic molecules have the energy to escape. Therefore the average energy of the molecules will decrease as its kinetic energy decreases.
• Liquids are therefore cooled by evaporation
• Higher temperature leads to higher rate of evaporation. Molecules only escape the attraction at the surface.
• Larger surface area leads to a higher rate of evaporation.
• Air movement; therefore increasing air movement increases the rate of evaporation.
• Evaporation can be used to cool things down.

Thermal Energy Transfer: Conduction

• Good thermal conductors are solids which easily transfer heat
• Bad thermal conductors (also called insulators) are solids which do not transfer heat well
• If the foot touches the tiles, heat is transferred away from the foot, making it feel cold, therefore heat is transferred by conduction
• Where the paper touched by the metal, heat was transferred from the paper into the metal, which prevents the paper from getting hot. Whereas, in the paper touching the wood, the heat was not transferred from the paper and paper started to burn.
• If something has metals, the answer should be related to conduction
• Conduction in solids requires close particles for vibrations to pass along.
• Liquids have closer particles but that slide past each other.
• Cases particles are widely spared and won't naturally touch each other.
• For conduction to occur in thermal conditions, the particles must be close together. Liquids are poor conductors of heat and gases are poor conductors of heat.
• Conductors tend to be metals as their metals have decolonised electrons with can transfer energy, meaning there is a wide range of thermal conductivity.

Ways that Heat Travels

• Convection is the main way heat travels through liquids and gases. It only works in fluid states. Heat rises and replaces it with a convection current.
• Liquids tend to cause convection.
• Radiation happens as heat emitting and the temperature of these bodies must depend on: The body of the temperature, the area of the body, and how much area is being emitted. Black tends to be a good observer.
• Convection happens when heat energy transferals in the form of liquid and gas. It travels through liquid and gas then moves with convection so the current happens, only in liquids.
• Sources placed in the bottom generate convection currents.