Temperature & Heat

Temperature as a measure of hotness

Temperature is a relative measure, or indication, of the hotness or coldness of a body. A hot utensil is said to have a high temperature and an ice cube a low temperature; an object at a higher temperature than another is said to be hotter. As NCERT stresses, hot and cold are relative terms, like tall and short — they describe a comparison, not an absolute amount.

We can perceive temperature directly by touch, but that sense is unreliable and its range is far too limited for scientific work. The same lukewarm water can feel warm to a hand just removed from ice and cold to a hand just removed from a flame. A quantitative, reproducible measure requires a thermometer and a temperature scale, and ultimately the absolute scale built from ideal-gas thermometry.

Heat as energy in transit

A glass of ice-cold water left on a table on a hot summer day eventually warms up, while a cup of hot tea on the same table cools down. In both cases the body and its surroundings start at different temperatures, and something flows between them until the two reach the same temperature. That something is heat.

NCERT defines it sharply: heat is the form of energy transferred between two or more systems, or between a system and its surroundings, by virtue of a temperature difference. The crucial word is transferred. Heat is energy caught in the act of crossing a boundary — energy in transit. The moment it has crossed and settled into the receiving body, it stops being heat and simply becomes part of that body's stored energy.

SI units — joule and calorie

Because heat is a form of energy, it is measured in the same unit as every other energy: the joule (J). NCERT states the SI unit of heat energy transferred is the joule, while the SI unit of temperature is the kelvin (K), with degree Celsius (°C) commonly used. The two quantities therefore carry entirely different units — a reliable signal that they are different physical entities.

An older, non-SI unit of heat is the calorie, defined as the heat required to raise the temperature of 1 g of water by 1 °C. The numerical bridge to the SI unit, the mechanical equivalent of heat, is \(1~\text{cal} = 4.186~\text{J}\). The kilocalorie used in nutrition is 1000 of these calories.

Direction of heat flow

Heat does not flow at random. When a system and its surroundings are at different temperatures, heat is transferred until both reach the same temperature — and it always travels from the body at higher temperature to the body at lower temperature. The ice-cold water absorbs heat from the warmer environment; the hot tea releases heat to the cooler environment. The direction is set by temperature, never by the amount of energy a body happens to store.

This one-way, spontaneous tendency — heat flowing of its own accord only from hot to cold — is the everyday face of the Second law of thermodynamics. Heat can be driven the other way, from cold to hot, but only by doing external work on the system, as a refrigerator or air-conditioner does. Left alone, the flow never reverses.

Thermal equilibrium

When heat flow between two bodies in contact finally ceases, the bodies have reached the same temperature and we say they are in thermal equilibrium. Equilibrium is therefore the end-state of the hot-to-cold process: there is no longer any temperature difference to drive a net transfer, so no net heat flows in either direction.

The defining feature of thermal equilibrium is equal temperature, and nothing else. Two bodies in thermal equilibrium need not be the same size, the same material, or hold the same amount of energy — they need only share a common temperature. This makes "same temperature" and "in thermal equilibrium" two ways of saying the same thing, which is exactly the foothold the Zeroth law builds on.

The Zeroth law of thermodynamics

The Zeroth law gives temperature its formal definition. It states: if two systems A and B are each separately in thermal equilibrium with a third system C, then A and B are in thermal equilibrium with each other. The property the three systems share when they are in mutual thermal equilibrium is precisely what we call temperature.

The law looks almost trivial, but it is what makes thermometry possible. Let system C be a thermometer. Bring it into equilibrium with body A and note its reading; bring the same thermometer into equilibrium with body B and note its reading. If the two readings agree, A and B are at the same temperature and would exchange no net heat if placed in contact — without ever having to touch A to B. The thermometer is the third body C, and the Zeroth law guarantees the inference.

Heat vs temperature vs internal energy

NEET repeatedly tests whether you can keep three closely related quantities apart. They share a tight web of relationships — heat changes temperature, temperature governs internal energy, internal energy is what heat transfer adds to — yet each is conceptually distinct.

The link is causal but not interchangeable. Supplying heat to a body usually raises its temperature, which reflects an increase in the average kinetic energy of its molecules and hence in its internal energy. But a large warm object can hold more internal energy than a small very hot one, and during a change of state heat flows in while the temperature does not change at all — the energy goes into latent heat, not into raising temperature. How much a body's temperature rises for a given heat input is controlled by its specific heat capacity.