Just Ask Antoine!
Glossary: Energy and chemical change
- adiabat. adiabatic line. Compare with adiabatic.
- A line on an indicator diagram that represents an adiabatic process.
- adiabatic. adiabatic process; isentropic process.
- A process that neither absorbs nor releases energy into the surroundings. For example, a chemical reaction taking place in a closed thermos bottle can be considered adiabatic. Very fast processes can often be considered adiabatic with respect to heat exchange with the surroundings, because heat exchange is not instantaneous.
- average bond enthalpy. Compare with bond enthalpy.
- Average enthalpy change per mole when the same type of bond is broken in the gas phase for many similar substances
- Boltzmann constant. (k) Boltzmann's constant.
- A fundamental constant equal to the ideal gas law constant divided by Avogadro's number, equal to 1.3805 × 10-23 J K-1.
- Boltzmann equation.
- A statistical definition of entropy, given by S = k ln W, where S and k are the entropy and Boltzmann's constant, respectively, and W is the probability of finding the system in a particular state.
- bond energy. Compare with bond enthalpy.
- Energy change per mole when a bond is broken in the gas phase for a particular substance.
- bond enthalpy. Compare with average bond enthalpy.
- Enthalpy change per mole when a bond is broken in the gas phase for a particular substance.
- The amount of heat required to raise the temperature of 1 g of water at 14.5°C to 15.5°C. One calorie is equivalent to exactly 4.184 J.
- An insulated vessel for measuring the amount of heat absorbed or released by a chemical or physical change.
- Experimental determination of heat absorbed or released by a chemical or physical change.
- empirical temperature.
- A property that is the same for any two systems that are in thermodynamic equilibrium with each other.
- endothermic. endothermic reaction; endothermic process. Compare with exothermic.
- A process that absorbs heat. The enthalpy change for an endothermic process has a positive sign.
- energy. Compare with heat and work.
- Energy is an abstract property associated with the capacity to do work.
- enthalpy. (H) enthalpy change. Compare with heat.
- Enthalpy (H) is defined so that changes in enthalpy (H) are equal to the heat absorbed or released by a process running at constant pressure. While changes in enthalpy can be measured using calorimetry, absolute values of enthalpy usually cannot be determined. Enthalpy is formally defined as H = U + PV, where U is the internal energy, P is the pressure, and V is the volume.
- enthalpy of atomization. (Hat) atomization enthalpy; heat of atomization.
- The change in enthalpy that occurs when one mole of a compound is converted into gaseous atoms. All bonds in the compound are broken in atomization and none are formed, so enthalpies of atomization are always positive.
- enthalpy of combustion. (Hc) heat of combustion.
- The change in enthalpy when one mole of compound is completely combusted. All carbon in the compound is converted to CO2(g), all hydrogen to H2O(), all sulfur to SO2(g), and all nitrogen to N2(g).
- enthalpy of fusion. (Hfus) heat of fusion; molar heat of fusion; molar enthalpy of fusion.
- The change in enthalpy when one mole of solid melts to form one mole of liquid. Enthalpies of fusion are always positive because melting involves overcoming some of the intermolecular attractions in the solid.
- enthalpy of hydration. (Hhyd) hydration enthalpy; heat of hydration.
- The change in enthalpy for the process
A(g)A(aq)where the concentration of A in the aqueous solution approaches zero. Enthalpies of hydration for ions are always negative because strong ion-water attractions are formed when the gas-phase ion is surrounded by water.
- enthalpy of reaction. (Hrxn) heat of reaction.
- The heat absorbed or released by a chemical reaction running at constant pressure.
- enthalpy of sublimation. (Hsub) heat of sublimation.
- The change in enthalpy when one mole of solid vaporizes to form one mole of gas. Enthalpies of sublimation are always positive because vaporization involves overcoming most of the intermolecular attractions in the sublimation.
- enthalpy of vaporization. (Hvap) heat of vaporization.
- The change in enthalpy when one mole of liquid evaporates to form one mole of gas. Enthalpies of vaporization are always positive because vaporization involves overcoming most of the intermolecular attractions in the liquid.
- entropy. (S)
- Entropy is a measure of energy dispersal. Any spontaneous change disperses energy and increases entropy overall. For example, when water evaporates, the internal energy of the water is dispersed with the water vapor produced, corresponding to an increase in entropy.
- exothermic. exothermic reaction; exothermic process. Compare with endothermic.
- A process that releases heat. The enthalpy change for an exothermic process is negative. Examples of exothermic processes are combustion reactions and neutralization reactions.
- first law. first law of thermodynamics.
- The first law states that energy cannot be created or destroyed. Many equivalent statements are possible, including: Internal energy changes depend only on the initial and final states of the system, not on the path taken. The work done during an adiabatic process depends only on the initial and final states of the system, and not on the path taken. The internal energy change for any cyclic process is zero.
- formation. formation reaction.
- A reaction that forms one mole of a compound from its elements in their most stable forms. For example, the formation reaction for water is H2(g) + ½O2 H2O().
- free energy.
- Energy that is actually available to do useful work. A decrease in free energy accompanies any spontaneous process. Free energy does not change for systems that are at equilibrium.
- Gibbs free energy. (G) Gibbs' free energy.
- A thermodynamic property devised by Josiah Willard Gibbs in 1876 to predict whether a process will occur spontaneously at constant pressure and temperature. Gibbs free energy G is defined as G = H - TS where H, T and S are the enthalpy, temperature, and entropy. Changes in G correspond to changes in free energy for processes occuring at constant temperature and pressure; the Gibbs free energy change corresponds to the maximum nonexpansion work that can be obtained under these conditions. The sign of DeltaG is negative for all spontaneous processes and zero for processes at equilibrium.
- Gibbs free energy of formation. (Gf) Gibbs' free energy of formation.
- The change in Gibbs free energy that accompanies the formation of one mole of a compound from its elements in their most stable form.
- heat capacity. Compare with molar heat capacity and specific heat.
- The heat required to raise the temperature of an object by 1°C is called the heat capacity of the object. Heat capacity is an extensive property with units of J K-1.
- heat. Compare with work, energy, enthalpy, and temperature.
- Heat is a transfer of energy that occurs when objects with different temperatures are placed into contact. Heat is a process, not a property of a material.
- Helmholtz free energy. (A) Arbeitfunktion.
- A thermodynamic property that can be used to predict whether a process will occur spontaneously at constant volume and temperature. Helmholtz free energy A is defined as A = U - TS where U, T and S are the internal energy, temperature, and entropy. Changes in A correspond to changes in free energy for processes occuring at constant temperature and volume. The sign of DeltaA is negative for spontaneous processes and zero for processes at equilibrium.
- Hess's law. law of constant heat summation; Hess's law of heat summation.
- The heat released or absorbed by a process is the same no matter how many steps the process takes. For example, given a reaction A B, Hess's law says that H for the reaction is the same whether the reaction is written as A C B or as A B. This is the same as writing that H(A B) = H(A C) + H(C B).
- indicator diagram. PV diagram.
- A plot of pressure vs. volume. Lines or curves on the indicator diagram represent processes. The areas under curves on the indicator diagram are equal to the work released by the process.
- internal energy. (U, E) Compare with enthalpy and energy.
- Internal energy (U) is defined so that changes in internal energy (U) are equal to the heat absorbed or released by a process running at constant volume. While changes in internal energy can be measured using calorimetry, absolute values of internal energy usually cannot be determined. Changes in internal energy are equal to the heat transferred plus the work done for any process.
- isobar. Compare with isotope.
- 1. A contour line that corresponds to values measured at identical pressures. For example, curves on a plot of gas volumes measured at different temperatures in an open container are isobars. 2. Nuclides that have the same isotopic mass but different atomic number.
- Having constant pressure.
- A contour line that corresponds to values measured at identical volumes. For example, a curve on a plot of gas pressure measured at different temperatures in a rigid container is an isochore.
- Having constant volume.
- A contour line that corresponds to values measured at identical temperatures. For example, curves on a plot of gas pressure measured at different volumes in a constant temperature bath are isotherms.
- Having constant temperature.
- joule. (J)
- The SI unit of energy, equal to the work required to move a 1 kg mass against an opposing force of 1 newton. 1 J = 1 kg m2 s-2 = 4.184 calories.
- kinetic energy. Compare with potential energy.
- The energy an object possesses by virtue of its motion. An object of mass m moving at velocity v has a kinetic energy of ½mv2.
- latent heat.
- Heat that is absorbed without causing a rise in temperature. For example, "latent heat of vaporization" refers to the amount of heat required to convert a liquid to vapor at a particular temperature.
- molar heat capacity. atomic heat capacity. Compare with molar heat capacity and specific heat.
- The heat required to raise the temperature of one mole of a substance by 1°C is called the molar heat capacity of the substance. Molar heat capacity is an intensive property with SI system units of J mol-1 K-1. The molar heat capacity of elements is sometimes called the "atomic heat capacity".
- perfect crystal.
- A crystal with no defects or impurities, made of completely identical repeating subunits. Further, a perfect crystal has only one possible arrangement of subunits, with every subunit making exactly the same contribution to the total energy of the crystal.
- potential energy. Compare with kinetic energy.
- energy an object possesses by virtue of its position. For example, lifting a mass mby h meters increases its potential energy by mgh, where g is the acceleration due to gravity.
- second law. second law of thermodynamics.
- The second law states that every spontaneous process causes a net increase in the entropy of the universe. Many alternative statements are possible, including: Heat cannot be converted to work via an isothermal cycle. Heat cannot be converted to work with 100% efficiency. Heat cannot flow from a cold object to a warmer object without doing outside work.
- specific heat. Compare with heat capacity.
- The heat required to raise the temperature of 1 g of a substance by 1°C is called the specific heat of the substance. Specific heat is an intensive property with units of J g-1 K-1.
- spontaneous. spontaneity; spontaneous process; spontaneous reaction.
- A spontaneous process occurs because of internal forces; no external forces are required to keep the process going, although external forces may be required to get the process started. For example, the burning of wood is spontaneous once the fire is started. The combination of water and carbon dioxide to reform the wood and oxygen is NOT spontaneous!
- standard entropy of reaction. (Srxn°) entropy of reaction.
- A change in entropy associated with a reaction involving substances in their standard states. A superscript circle (°) distinguishes standard enthalpy changes from enthalpy changes which involve reactants and products that are not in their standard states.
- standard enthalpy change. (H°) standard enthalpy. Compare with enthalpy change.
- A change in enthalpy associated with a reaction or transformation involving substances in their standard states.
- standard enthalpy of formation. (Hf°) standard heat of formation; heat of formation; enthalpy of formation.
- The change in enthalpy when one mole of compound is formed from its elements in their most stable form and in their standard states.
- standard enthalpy of reaction. (Hrxn°) standard heat of reaction.
- A change in enthalpy associated with a reaction involving substances in their standard states.
- standard molar entropy. (S°)
- The entropy of one mole of a substance in its standard state.
- standard state. (° or
- A set of conditions defined to allow convenient comparison of thermodynamic properties. The standard state for a gas is the the state of the pure substance in the gaseous phase at the standard pressure, with the gas behaving ideally. The standard state for liquids and solids is the state of the most stable form of the substance at the standard pressure. Temperature is not included in the definition of standard state and must be specified, but when not given a temperature of 25°C is usually implied.
- standard pressure. (P° or P
- Standard pressure is a pressure of 1 bar. Before 1982, the standard pressure was 1 atm (1 atm = 1.01325 bar).
- state function.
- A property that depends only on the condition or "state" of the system, and not on the path used to obtain the current conditions. Energy, enthalpy, temperature, volume, pressure, and temperature are examples of state functions; heat and work are examples of non-state functions.
- temperature. Compare with heat and thermodynamic temperature.
- Temperature is an intensive property associated with the hotness or coldness of an object. It determines the direction of spontaneous heat flow (always from hot to cold).
- thermal energy.
- energy an object possesses by virtue of its temperature. For example, 1 g of water at 15°C has 4.184 J more energy than 1 g of water at 14°C.
- thermochemical equation.
- An compact equation representing a chemical reaction that describes both the stoichiometry and the energetics of the reaction. For example, the thermochemical equation CH4(g) + 2 O2(g) CO2(g) + 2 H2O(g), H = -2220 kJ means "When one mole of gaseous CH4 is burned in two moles of oxygen gas, one mole of CO2 gas and 2 moles of steam are produced, and 2220 kilojoules of heat are released."
- The study of heat absorbed or released during chemical changes.
- thermodynamic equilibrium.
- A system is at thermodynamic equilibrium if the energy it gains from its surroundings is exactly balanced by the energy it loses, no matter how much time is allowed to pass.
- thermodynamics. thermodynamic.
- The study of energy transfers and transformations.
- The science of temperature measurement.
- water gas. blue gas; synthesis gas.
- A fuel gas used in industrial synthesis of organic chemicals, and in welding, glassmaking, and other high-temperature industrial applications. Water gas made by passing steam over a bed of hot coal or coke. It consists mainly of of carbon monoxide (CO) and hydrogen (H2), contaminated with small amounts of CO2, N2, CH4, and O2.
- work. Compare with heat.
- Work is the energy required to move an object against an opposing force. Work is usually expressed as a force times a displacement. Dropping a stone from a window involves no work, because there is no force opposing the motion (unless you consider air friction...). Pushing against a stone wall involves no work, unless the stone wall actually moves.