universal gas constant
(noun)Definition of universal gas constant
the constant of proportionality that relates the energy scale in physics to the temperature scale
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Examples of universal gas constant in the following topics:

Expressing the Equilibrium Constant of a Gas in Terms of Pressure
 For gasspecific reactions, you can use the equilibrium concentration.
 The equilibrium constant for gases can therefore also be expressed in terms of partial pressures .
 We can show this relationship mathematically using the ideal gas equation: PV = nRT where P is pressure, V is volume, n is concentration, R is the universal gas constant, and T is temperature.
 Because "RT" is constant as long as the temperature is constant, the pressure is directly proportional to the concentration.
 The equilibrium expression can be expressed in terms of partial pressure for a gasphase reaction: $K_{p} = K_{c}(RT) ^{ \Delta n}$ where Kp is the equiibrium constant expressed in partial pressure, R is the universal gas constant, T is temperature, and \Delta n is the number of moles changing in the reaction.
 For gasspecific reactions, you can calculate the equilibrium constant using concentration or partial pressure.

The Nernst Equation
 The two (ultimately equivalent) equations for these two cases (halfcell, full cell) are as follows: Ered = Estd,red  RT/zF ln ared/aox (halfcell reduction potential) Ecell = Estd,cell  RT/zF ln Q (total cell potential) where: Ered is the halfcell reduction potential at the temperature of interest Estd,red is the standard halfcell reduction potential Ecell is the cell potential (electromotive force) Estd,cell is the standard cell potential at the temperature of interest R is the universal gas constant T is the absolute temperature a is the chemical activity for the relevant species F is the Faraday constant z is the number of moles transferred in the cell reaction or halfreaction Q is the reaction quotient

Van der Waals Equation
 This equation of state can be presented as: $(p + a/Vm2)(Vmb) = RT$ where p is the pressure, Vm is the molar volume, R is the universal gas constant, and T is the absolute temperature.
 The constants a and b have positive values and are specific to each gas.
 The term involving the constant a corrects for intermolecular attraction.
 The b term represents the excluded volume of the gas or the volume occupied by the gas particles.
 The van der Waals equation becomes the ideal gas law as these two correction terms approach zero.
 The van der Waals equation is a modification to the ideal gas law that corrects for nonzero gas volume and intermolecular interactions.

The Effect of Intermolecular Forces
 The ideal gas law is a convenient approximation for predicting the behavior of gasphase chemical reactions.
 The universal attractive force, or London disperson force, also generally increases with molecular weight. the London dispersion force is caused by correlated movements of the electrons in interacting molecules.
 Finally, as a gas is compressed and pressure increases, repulsive forces from the gas molecules oppose the decrease in volume.
 He corrected for the intermolecular attractive forces which hold the particles together and reduce the pressure on the container by replacing the pressure term with $P + (a/V_m^2)$ where Vm is the molar volume and a is a constant specific to each gas.
 The van der Waals equation of state can then be presented as: $(P+(a/V_m^2))(V_mb) = RT$ Where b is the excluded volume, R is the universal gas constant, and T is the absolute temperature.

The Arrhenius Equation
 The Arrhenius equation is a simple, but remarkably accurate, formula for the temperature dependence of the reaction rate constant, and therefore, the rate of a chemical reaction.
 A is the preexponential factor and R is the universal gas constant.
 What is "decaying" here is not the concentration of a reactant as a function of time, but the magnitude of the rate constant as a function of the exponent –Ea /RT.
 So, what would limit the rate constant if there were no activation energy requirements?
 They are sometime estimated by comparing the observed rate constant with the one in which A is assumed to be the same as Z.