SRS Unit 1.1 ATAR Physics - Heating Processes

Kinetic Particle Model

• A model stating that all matter is made up of tiny particles that are constantly moving.

Solid

• A state of matter where particles are close together in a regular pattern and vibrate on the spot.

• Low energy.

Liquid

• A state of matter where particles are close together in a random arrangement and move around each other.

• Medium energy.

Gas

• A state of matter where particles are far apart in a random arrangement and move quickly in all directions.

• High energy

Thermal Energy (U)

• The internal energy present in a system due to its temperature.

• Measured in Joules (J).

Temperature

• The degree or intensity of heat present in a substance or object, especially as expressed according to a comparative scale, e.g. the Celsius temperature scale.

• It is related to the average kinetic energy of atoms and molecules in a system.

Kinetic Energy (KE)

• The energy resulting from the movement of an object.

• When associated with temperature, it is associated with the motion of particles in a substance.

Heat

• The energy transferred from one system to another because of a difference in temperature, usually a transfer of thermal energy.

• A transfer or change in energy is known as Work (Q)

Internal Energy

• The total potential energy and kinetic energy of the particles in a system.

Conduction

• Heat transferred when neighbouring atoms or molecules are in contact.

Convection

• Heat transferred by movement of a heated fluid such as air or water.

Radiation

• Heat transferred in electromagnetic/light waves emitted from a source that may be absorbed, reflected, or transmitted through a colder body.

Absolute zero (0 K)

• The temperature at which there is no molecular motion

• Nothing can reach it because it breaks a lot of physics.

Celsius (°C)

• Temperature relative to real life

• Represented by T_C

• 0°C = 273K

Kelvin (K)

• Absolute temperature, SI unit

• Represented by T_K

• 0K = -273°C

Q ∝ ∆T

• The change in temperature of a substance is dependent on the heat source and energy transferred.

Q ∝ ∆m

• The mass of the substance influences the rate of heating and cooling.

Specific Heat Capacity (c)

• The amount of energy (heat) required to change the temperature of 1kg of the substance by 1K or 10C.

• The units are Jkg^-1K^-1.

Q = mc∆T

• Q is the heat energy transferred (J)

• m is the mass of the substance (kg)

• c is the Specific Heat Capacity of the substance (Jkg^-1K^-1)

• ∆T is the change in temperature (K or °C)

Phase change

• When a substance changes state.

• The temperature doesn’t change because all the work is going into either breaking the molecules apart (heating) or putting then together (cooling).

• This energy used is called Latent Heat

Latent Heat (L)

• The amount of energy needed to change a substance’s state.

• Its units are Jkg^-1

Latent Heat of Fusion (L_F)

• Latent Heat for the solid ⇔ liquid phase change

Latent Heat of Vapourisation

• Latent Heat for the liquid ⇔ gas phase change

Q = mL

• Q is the heat energy transferred (J)

• m is the mass of the substance (kg)

• L is the Latent Heat of the substance (Jkg^-1)

Heating curve

• A graphical representation of the temperature and phase changes of a substance.

• The horizontal sections of the curve represent the phase changes (latent heat) and the diagonal sections represent changing temperatures in each state (specific heat capacity)

Laws of Thermodynamics

• Laws that define fundamental physical quantities (temperature, energy, and entropy) that characterize thermodynamic systems.

Zeroth Law of Thermodynamics

• If two thermodynamic systems are each in thermal equilibrium with a third system, then they are in thermal equilibrium with each other.

Thermal equilibrium

• When two or more objects are at the same temperature and their average kinetic energies are equal.

• Heat is transferred from the hotter substance to the cooler one until they reach thermal equilibrium and the transfer stops. The heat lost by the hotter substance is gained by the cooler one.

First Law of Thermodynamics/Law of Conservation of Energy

• Energy in an isolated system can neither be created nor destroyed; only transferred or changed from one form to another.

• An example of this is when two or more objects achieve thermal equilibrium.

Calorimetry

• The science or act of measuring the thermal energy in a substance. It is an application of the Laws of Thermodynamics and has many uses.

• Equations Q_lost = Q_gained and mc₁∆T = mc₂∆T are used

Mechanical Work / Work (W)

• Energy transferred to or from an object via the application of a force.

∆U = Q - W

• ∆U is the change in internal energy (J)

• Q is the energy added by heating (+Q) or removed by cooling (-Q)

• W is the energy added by work done on the system (-W) or removed by work done by the system (+W)

• If W > Q, then the total internal energy goes down (-∆U)

• If Q > W, then the total internal energy goes up (+∆U)

Efficiency (η)

• A way to measure how well energy is transferred or transformed in mechanical systems through comparing used and lost energy

• η% = ( energy output / energy input ) × 100%

• η% = ( energy used / total energy ) × 100%