1. Heat, Temperature and Thermal Energy
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1.1 What is Heat?
Every substance is made up of tiny particles — atoms and molecules — that are in constant, random motion. This motion may be vibrational (as in solids), or translational and rotational (as in liquids and gases). The total kinetic energy and potential energy possessed by all these particles due to their motion and mutual interactions is called the internal energy or thermal energy of the body. Heat is defined as the thermal energy that flows from one body (or one part of a body) to another because of a temperature difference between them. Heat is therefore not something a body 'contains' — it is energy in transit. Once heat is transferred into a body, it becomes part of that body's internal energy; it is inaccurate to say a hot cup of tea 'contains a lot of heat' — it contains a large amount of internal energy, some of which will flow out as heat if placed in cooler surroundings.
Heat, being a form of energy, is measured in the SI unit joule (J). An older, still commonly used unit is the calorie (cal): one calorie is the amount of heat required to raise the temperature of 1 gram of water by 1°C (specifically from 14.5°C to 15.5°C). The relationship between the two units, established through Joule's mechanical equivalent of heat experiments, is: 1 calorie = 4.184 joules (commonly rounded to 4.2 J). This equivalence was historically important because it proved that heat is a form of energy and not, as earlier believed, an invisible weightless fluid called 'caloric'.
1.2 What is Temperature?
Temperature is a measure of the degree of hotness or coldness of a body. More precisely, in kinetic theory terms, temperature is a measure of the average kinetic energy of the molecules of a substance due to their random motion. A body at a higher temperature has molecules moving, on average, faster (or vibrating more vigorously) than a body at a lower temperature. Temperature is a scalar quantity and, unlike heat, it is not a form of energy — it is a physical property that determines the direction of heat flow: heat always flows spontaneously from a body at a higher temperature to one at a lower temperature, until thermal equilibrium is reached (that is, until both bodies attain the same temperature).
1.3 Key Distinctions
Aspect | Heat | Temperature |
|---|---|---|
Definition | Energy transferred due to temperature difference | Degree of hotness or coldness |
Nature | Form of energy (extensive property) | Not energy; an intensive property |
SI Unit | Joule (J) | Kelvin (K) |
Depends on | Mass, substance, and temperature change | Independent of mass or amount of substance |
Measuring device | Calorimeter | Thermometer |
Example | Amount of energy needed to boil a litre of water | The reading of 100°C at which water boils |
Exam tip: A common SSC/RRB trap: two bodies of different mass at the same temperature contain different amounts of heat energy (a large pot of water at 60°C has more internal energy than a teaspoon of water at 60°C), even though both are equally 'hot' as measured by a thermometer. Temperature does not depend on the quantity of matter; heat content does.