standard state

(noun)

In chemistry, a reference point used to calculate a material's (pure substance, mixture, or solution) properties under different conditions.

Related Terms

Examples of standard state in the following topics:

  • Standard States and Standard Enthalpy Changes

    • In principle, the choice of standard state is arbitrary, although the International Union of Pure and Applied Chemistry (IUPAC) recommends a conventional set of standard states for general use.
    • Strictly speaking, temperature is not part of the definition of a standard state; the standard state of a gas is conventionally chosen to be 1 bar for an ideal gas, regardless of the temperature.
    • Standard states for atomic elements are given in terms of the most stable allotrope for each element.
    • The standard state should not be confused with standard temperature and pressure (STP) for gases, or with the standard solutions used in analytical chemistry.
    • Standard states are often indicated in textbooks by a circle with a horizontal bar $H^\ominus_f$.

  • Standard Free Energy Changes

    • The concept of standard states is especially important in the case of free energy, so take a moment to review it.
    • For most practical purposes, the following definitions of standard states are acceptable:
    • Don't confuse these thermodynamic standard states with the "standard temperature and pressure" (STP) widely employed in gas law calculations.
    • Recall that the symbol ° refers to the standard state of a substance measured under the conditions of 1 atm pressure or an effective concentration of 1 Molar and a temperature of 298K.
    • The standard Gibbs free energy of formation of a compound is the change of Gibbs free energy that accompanies the formation of 1 mole of that substance from its component elements, at their standard states.

  • Standard Enthalpy of Reaction

    • The change in enthalpy does not depend upon the particular pathway of a reaction, but only upon the overall energy level of the products and reactants; enthalpy is a state function, and as such, it is additive.
    • In order to calculate the standard enthalpy of a reaction, we can sum up the standard enthalpies of formation of the reactants and subtract this from the sum of the standard enthalpies of formation of the products.
    • Stated mathematically, this gives us:
    • Note that because it exists in its standard state, the standard enthalpy of formation for oxygen gas is 0 kJ/mol.
    • A calculation of standard enthalpy of reaction (∆H°rxn) from standard heats of formation (∆H°f)

  • Thermodynamics of Redox Reactions

    • The thermodynamics of redox reactions can be determined using their standard reduction potentials and the Nernst equation.
    • The Nernst equation allows the reduction potential to be calculated at any temperature and concentration of reactants and products; the standard reaction potential must be measured at 298K and with each solution at 1M.
    • This equation allows the equilibrium constant to be calculated just from the standard reduction potential and the number of electrons transferred in the reaction.
    • The relationship between the Gibbs free energy change and the standard reaction potential is:
    • Translate between the equilibrium constant/reaction quotient, the standard reduction potential, and the Gibbs free energy change for a given redox reaction

  • Standard Entropy

    • The standard entropy of a substance is its entropy at 1 atm pressure.
    • Some typical standard entropy values for gaseous substances include:
    • Scientists conventionally set the energies of formation of elements in their standard states to zero.
    • Entropy, however, measures not energy itself, but its dispersal among the various quantum states available to accept it, and these exist even in pure elements.
    • The standard entropy of reaction helps determine whether the reaction will take place spontaneously.

  • Standard Reduction Potentials

    • Standard reduction potentials provide a systematic measurement for different molecules' tendency to be reduced.
    • The standard reduction potential is defined relative to a standard hydrogen electrode (SHE) reference electrode, which is arbitrarily given a potential of 0.00 volts.
    • The values below in parentheses are standard reduction potentials for half-reactions measured at 25 °C, 1 atmosphere, and with a pH of 7 in aqueous solution.
    • Historically, many countries, including the United States and Canada, used standard oxidation potentials rather than reduction potentials in their calculations.
    • These are simply the negative of standard reduction potentials, so it is not a difficult conversion in practice.

  • Free Energy and Cell Potential

    • In electrochemistry, the standard electrode potential, abbreviated E°, is the measure of the individual potential of a reversible electrode at standard state, which is with solutes at an effective concentration of 1 M, and gases at a pressure of 1 atm.
    • Since the standard electrode potentials are given in their ability to be reduced, the bigger the standard reduction potentials, the easier they are to be reduced; in other words, they are simply better oxidizing agents.
    • In the example of Zn2+, whose standard reduction potential is -0.76 V, it can be oxidized by any other electrode whose standard reduction potential is greater than -0.76 V and can be reduced by any electrode with standard reduction potential less than -0.76 V.

  • Hess's Law

    • This law states that if a reaction takes place in several steps, then the standard reaction enthalpy for the overall reaction is equal to the sum of the standard enthalpies of the intermediate reaction steps, assuming each step takes place at the same temperature.
    • Since enthalpy is a state function, the change in enthalpy between products and reactants in a chemical system is independent of the pathway taken from the initial to the final state of the system.
    • However, because we know the standard enthalpy change for the oxidation for these two substances, it is possible to calculate the enthalpy change for this reaction using Hess's law.
    • First it looks at combining reactions according to Hess's law and their heats of reaction, and then it discusses using standard heats of formation of the reactants and products to find the overall heat of reaction.

  • Transition State Theory

    • Transition state theory (TST) describes a hypothetical "transition state" that occurs in the space between the reactants and the products in a chemical reaction.
    • The species that is formed during the transition state is known as the activated complex.
    • If the rate constant for a reaction is known, TST can be used successfully to calculate the standard enthalpy of activation, the standard entropy of activation, and the standard Gibbs energy of activation.
    • According to transition state theory, between the state in which molecules exist as reactants and the state in which they exist as products, there is an intermediate state known as the transition state.
    • The species that forms during the transition state is a higher-energy species known as the activated complex.

  • Standard Units (SI Units)

    • The imperial system is used for "everyday" measurements in a few places, such as the United States.
    • A light source made to produce 20 cd will be the same regardless of whether it is made in the United States, in the UK, or anywhere else.
    • The second (s) was originally based on a "standard day" of 24 hours, with each hour divided in 60 minutes and each minute divided in 60 seconds.
    • Therefore, a second is now defined as the duration of 9,192,631,770 periods of the radiation corresponding to the transition between the two hyperfine levels of the ground state of the cesium-133 atom.
    • The candela (cd) was so named to refer to "candlepower" back in the days when candles were the most common source of illumination (because so many people used candles, their properties were standardized).