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Heat is the energy transferred from one body to another as the result of a difference in temperature. Heat flows from a hotter body to a colder body when the two bodies are brought together. This transfer of energy usually results in an increase in the temperature of the colder body and a decrease in that of the hotter body. A substance may absorb heat without an increase in temperature as it changes from one phase to another--that is, when it melts or boils. The distinction between heat (a form of energy) and temperature (a measure of the amount of energy) was clarified in the 19th century by such scientists as Fourier, Kirchhoff, and Boltzmann.
Measure of temperature is expressed in terms of any of several arbitrary scales, such as Fahrenheit, Celsius, or Kelvin. Heat flows from a hotter body to a colder one and continues to do so until both are at the same temperature. Temperature is a measure of the average energy of the molecules of a body, whereas heat is a measure of the total amount of thermal energy in a body. For example, whereas the temperature of a cup of boiling water is the same as that of a large pot of boiling water (212°F, or 100°C), the large pot has more heat, or thermal energy, and it takes more energy to boil a pot of water than a cup of water. The most common temperature scales are based on arbitrarily defined fixed points. The Fahrenheit scale sets 32° as the freezing point of water and 212° as the boiling point of water (at standard atmospheric pressure). The Celsius scale defines the triple point of water (at which all three phases, solid, liquid, and gas, coexist in equilibrium) at 0.01° and the boiling point at 100°. The Kelvin scale, used primarily for scientific and engineering purposes, sets the zero point at absolute zero and uses a degree the same size as those of the Celsius scale. So in this scheme Temperature and Heat seems to be defined in terms of each other, however Temperature is not treated as more fundamental than heat. That is why it even has its own fundamental unit Kelvin in the SI system.
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- Calorimetry: Calorimetry is the science of measurement or calculation of the amount of heat evolved or absorbed in a chemical reaction, change of state, or formation of a solution.
- Cryogenics: Cryogenics is the study and use of low-temperature phenomena. The cryogenic temperature range is from -238°F (-150°C) to absolute zero. At low temperatures, matter has unusual properties. Substances that are naturally gases can be liquefied at low temperatures, and metals lose electrical resistance as they get colder (see superconductivity). Cryogenics dates from 1877, when oxygen was first cooled to the point at which it became a liquid (-297°F, or -183°C); superconductivity was discovered in 1911. Applications of cryogenics include the storage and transport of liquefied gases, food preservation, cryosurgery, rocket fuels, and superconducting electromagnets.
Absolute Zero is the temperature at which a thermodynamic system has the lowest energy, 0 kelvin (K). It corresponds to -459.67°F (-273.15°C) and is the lowest possible temperature theoretically achievable by a system. A gas at constant pressure contracts as the temperature is decreased. A perfect gas would reach zero volume at absolute zero. However, a real gas condenses to a liquid or a solid at a temperature higher than absolute zero. At absolute zero, the system's molecular energy is minimal and none is available for transfer to other systems. Near absolute zero temperatures are routinely achievable in labs but attainment of absolute zero still remains controversial.
- Heat Capacities of Gases: Ratio of heat absorbed by a material to the change in temperature. It is usually expressed as calories per degree in terms of the amount of the material being considered. Heat capacity and its temperature variation depend on differences in energy levels for atoms. Heat capacities are measured with a calorimeter and are important as a means of determining the entropies of materials. Heat capacities of gases are different when they are measured in isothermal or isobaric conditions.
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Thermodynamics is concerned with the relationships between the state coordinates (properties, such as pressure, temperature, or chemical composition) of a system established by studying the relationships among heat, work, temperature, and energy. Any physical system will spontaneously approach an equilibrium that can be described by specifying its state coordinates. If external constraints are allowed to change, these properties generally change. The three laws of thermodynamics describe these changes and predict the equilibrium state of the system. The first law states that whenever energy is converted from one form to another, the total quantity of energy remains the same. The second law states that, in a closed system, the entropy of the system does not decrease. The third law states that, as a system approaches absolute zero, further extraction of energy becomes more and more difficult, eventually becoming theoretically impossible.
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