14. Temperature

When two regions have different temperatures, thermal energy is transferred from the hotter region to the colder until equilibrium is reached.

Temperature indicates the direction of net thermal transfer; the mechanism may be conduction, convection or radiation.

A hot metal block in cooler water loses thermal energy while the water gains it, even if the block contains less total energy.

Heat is energy in transfer, not a substance stored in an object, and transfer direction is not set simply by which object has more energy.

Two systems are in thermal equilibrium when their temperatures are equal, so there is no net transfer of thermal energy between them in contact.

Microscopic exchanges may still occur, but they balance on average. Thermal equilibrium is not the same as identical internal energy.

A thermometer left long enough in a liquid reaches the liquid’s temperature and is then in thermal equilibrium with it.

Equal temperature does not mean equal mass, energy content or particle number.

A thermometric property varies with temperature and can be calibrated to provide a temperature reading; examples include liquid volume, resistance or thermocouple emf.

A useful property should be monotonic over the range, reproducible and sensitive enough to resolve the required temperature change.

The resistance of a platinum sensor can be calibrated against fixed points and then used to infer an unknown temperature.

A property changing with temperature is not automatically a good thermometer; calibration and response range are essential.

The thermodynamic temperature scale is based on universal physical principles rather than a chosen material property; the kelvin is the SI unit.

A practical thermometer is calibrated to approximate this scale, but its raw property may be nonlinear or limited in range.

A gas, resistance and radiation thermometer can agree after calibration even though their measured properties differ.

The Celsius scale and a material’s expansion are convenient representations, not the fundamental definition of temperature.

Thermodynamic temperature T in kelvin relates to Celsius temperature θ by T=θ+273.15.

Kelvin is an absolute scale with the same degree size as Celsius but a different zero. Use kelvin in gas and thermodynamic equations unless instructed otherwise.

25 °C is 298.15 K; 0 °C is 273.15 K, not 0 K.

A temperature difference of 1 °C equals 1 K, but an absolute temperature of 1 °C is not 1 K.

Absolute zero is zero kelvin, the lowest limit of thermodynamic temperature; it corresponds to −273.15 °C.

It is a limiting state, not simply “no motion” in every quantum description. Use it as the zero of the absolute scale.

Cooling from 300 K to 150 K halves the absolute temperature even though Celsius readings do not behave as a ratio scale.

Negative Celsius temperatures can be physically valid, but temperatures below 0 K are not reached in the ordinary thermodynamic scale.

Specific heat capacity c is energy required to raise 1 kg of a substance by 1 K: ∆E=mc∆T.

Use temperature difference, not absolute temperature, and distinguish the substance mass from container or heater losses.

Heating 2.0 kg of water by 5 K requires about 42 kJ if c=4200 J kg⁻¹ K⁻¹.

A high specific heat capacity means more energy per kilogram per degree, not that the object always reaches a higher final temperature.

Specific latent heat L is energy required per kilogram to change state at constant temperature: E=mL. Fusion melts solid to liquid; vaporisation changes liquid to gas.

During the phase change, energy changes intermolecular potential-energy store rather than raising temperature.

Melting 0.20 kg of ice at its melting point requires E=mL_f, while its temperature remains constant during the melt.

Latent heat is not “hidden temperature”; it is energy transferred during a state change, and fusion/vaporisation have different values.