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23 Cards in this Set

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In physics and other sciences, energy (from the Greek ενεργός, energos, "active, working")[1] is a scalar physical quantity that is a property of objects and systems of objects which is conserved by nature. Several different forms, such as kinetic, potential, thermal, electromagnetic, chemical, nuclear, and mass have been defined to explain all known natural phenomena.
Law Of Conservation Of Energy
In physics, the conservation of energy states that the total amount of energy in any closed system remains constant but can be recreated, although it may change forms, e.g. friction turns kinetic energy into thermal energy. In thermodynamics, the first law of thermodynamics is a statement of the conservation of energy for thermodynamic systems, and is the more encompassing version of the conservation of energy. In short, the law of conservation of energy states that energy can not be created or destroyed, it can only be changed from one form to another.
Potential Energy
Potential energy can be thought of as energy stored within a physical system. This energy can be released or converted into other forms of energy, including kinetic energy. It is called potential energy because it has the potential to change the states of objects in the system when the energy is released.
Kinetic Energy
The kinetic energy of an object is the extra energy which it possesses due to its motion. It is defined as the work needed to accelerate a body of a given mass from rest to its current velocity. Having gained this energy during its acceleration, the body maintains this kinetic energy unless its speed changes. Negative work of the same magnitude would be required to return the body to a state of rest from that velocity.
In physics, heat, symbolized by Q, is energy transferred from one body or system to another due to a difference in temperature.[1][2] In thermodynamics, the quantity TdS is used as a representative measure of heat, which is the absolute temperature of an object multiplied by the differential quantity of a system's entropy measured at the boundary of the object.
In thermodynamics, work is the quantity of energy transferred from one system to another without an accompanying transfer of entropy. It is a generalization of the concept of mechanical work in mechanics. In the SI system of measurement, work is measured in joules (symbol: J). The rate at which work is performed is power.
State Function
In thermodynamics, a state function, or state quantity, is a property of a system that depends only on the current state of the system, not on the way in which the system got to that state. A state function describes the equilibrium state of a system. For example, internal energy, enthalpy and entropy are state quantities because they describe quantitatively an equilibrium state of thermodynamic systems. At the same time, mechanical work and heat are process quantities because they describe quantitatively the transition between equilibrium states of thermodynamic systems
System (from Latin systēma, in turn from Greek σύστημα systēma) is a set of interacting or interdependent entities, real or abstract, forming an integrated whole.
Surroundings are the area around a given physical or geographical point or place. The exact definition depends on the field. Surroundings can also be used in geography and mathematics, as well as philosophy, with the literal or metaphorically extended definition.

However, in thermodynamics, the term as a slightly more exact use: surroundings are seen as a medium for the exchange of materials or energy, and can be considered as an open system.
In thermodynamics, the word exothermic describes a process or reaction that releases energy in the form of heat. Its etymology stems from the Greek prefix ex-, meaning “outside” and the Greek word thermein, meaning “to heat”.
In thermodynamics, the word endothermic describes a process or reaction that absorbs energy in the form of heat. Its etymology stems from the Greek prefix endo-, meaning “inside” and the Greek suffix –thermic, meaning “to heat”
Thermodynamics (from the Greek θερμη, therme, meaning "heat" and δυναμις, dynamis, meaning "power") is a branch of physics that studies the effects of changes in temperature, pressure, and volume on physical systems at the macroscopic scale by analyzing the collective motion of their particles using statistics.
First Law Of Thermodynamics
In thermodynamics, the first law of thermodynamics is an expression of the more universal physical law of the conservation of energy.
Succinctly, the first law of thermodynamics states:

The increase in the internal energy of a system is equal to the amount of energy added by heating the system, minus the amount lost as a result of the work done by the system on its surroundings.
Internal Energy
In thermodynamics, the internal energy of a thermodynamic system, or a body with well-defined boundaries, denoted by U, or sometimes E, is the total of the kinetic energy due to the motion of molecules (translational, rotational, vibrational) and the potential energy associated with the vibrational and electric energy of atoms within molecules or crystals.
In thermodynamics and molecular chemistry, the enthalpy or heat content (denoted as H or ΔH, or rarely as χ) is a quotient or description of thermodynamic potential of a system, which can be used to calculate the "useful" work obtainable from a closed thermodynamic system under constant pressure.
A calorimeter is a device used for calorimetry, the science of measuring the heat of chemical reactions or physical changes as well as heat capacity.
Calorimetry is the science of measuring the heat of chemical reactions or physical changes. Calorimetry involves the use of a calorimeter.
Heat capacity
Heat capacity (symbol: Cp) — as distinct from specific heat capacity — is the measure of the heat energy required to increase the temperature of an object by a certain temperature interval. Heat capacity is an extensive property because its value is proportional to the amount of material in the object; for example, a bathtub of water has a greater heat capacity than a cup of water.
Specific heat capacity
Specific heat capacity, also known simply as specific heat, is the measure of the heat energy required to increase the temperature of a unit quantity of a substance by a certain temperature interval.
Molar heat capacity
When measuring specific heat capacity in science and engineering, the unit quantity of a substance is often in terms of mass: either the gram or kilogram, both of which are an SI unit. Especially in chemistry though, the unit quantity of specific heat capacity may also be the mole, which is a certain number of molecules or atoms. When the unit quantity is the mole, the term molar heat capacity may also be used to more explicitly describe the measure
Hess's law
Hess's Law is a law of physical chemistry named for Germain Hess's expansion of the Hess Cycle and used to predict the enthalpy change and conservation of energy (denoted as state function ΔH) regardless of the path through which it is to be determined
Standard enthalpy of formation
The standard enthalpy of formation or "standard heat of formation" of a compound is the change of enthalpy that accompanies the formation of 1 mole of a substance in its standard state from its constituent elements in their standard states (the most stable form of the element at 100 kPa of pressure and the specified temperature, usually 298 K or 25 degrees Celsius). Its symbol is ΔHfO
Standard state
In chemistry, the standard state of a material is its state at 1 bar (100 kilopascals exactly). This pressure was changed from 1 atm (101.325 kilopascals) by IUPAC in 1990.[1] The standard state of a material can be defined at any given temperature, most commonly 25 degrees Celsius, although quite a few texts (especially in related disciplines such as physics and engineering) use 0 degrees Celsius for Standard Temperature and Pressure (STP).