Temperature is a measure of the average kinetic energy of particles in a sample of matter. In physics, temperature is defined as a property of objects and systems that determines the direction of heat flow between them when they are in contact. The SI unit for temperature is the kelvin (K). Temperature is often measured with a thermometer, which may be calibrated to any convenient scale such as the Celsius scale or Fahrenheit scale.
The zeroth law of thermodynamics states that if two systems are in thermal equilibrium with each other, then they must have the same temperature. This law allows us to define an absolute scale for temperature, independent of any particular reference point. The most common scales in use today are the Celsius (°C) and Kelvin (K) scales. On the Celsius scale, water freezes at 0 °C and boils at 100 °C; on the Kelvin scale, these same temperatures correspond to 273.15 K and 373.15 K, respectively.
The first law of thermodynamics states that when energy passes into or out of a system as heat, there must be an accompanying change in the system’s internal energy. This principle provides us with a means of defining work and heat independently from each other: work can be thought of as changes in energy due solely to external forces acting on the system boundary, while heat refers only to those changes in internal energy that result from thermal interactions between different parts of the system itself. Heat always flows from hotter objects to cooler ones until both reach thermal equilibrium; this process transfers entropy from one place to another, resulting in an overall increase in entropy over time. The second law of thermodynamics quantifies this behavior by stating that it is impossible for any process to occur whose sole effect is to transfer heat from a colder body to a hotter one without also doing some work along the way; this statement implies that it is also impossible for entropy ever decrease over time within an isolated system.