# universal gas constant

(noun)

## Definition of universal gas constant

the constant of proportionality that relates the energy scale in physics to the temperature scale

Source: Wikipedia - CC BY-SA 3.0

## Examples of universal gas constant in the following topics:

• ### Expressing the Equilibrium Constant of a Gas in Terms of Pressure

• For gas-specific reactions, you can use the equilibrium concentration.
• Increasing the pressure of a gas has exactly the same effect as increasing its concentration.
• The equilibrium constant for gases can therefore also be expressed in terms of partial pressures .The pressure of a given mass of gas is increased by squeezing it into a smaller volume.
• We can show this relationship mathematically using the ideal gas equation:PV = nRTwhere 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 gas-phase 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.Examples of Using Pressure to Shift EquilibriumConsider the reaction for the formation of NH3:${N}_{2}(g)+{3H}_{2}(g)\rightleftharpoons {2NH}_{3}(g)$The total number of moles of reactants is four, and the total number of moles of products is two.
• For gas-specific reactions, you can calculate the equilibrium constant using concentration or partial pressure.
• ### Van der Waals Equation

• This equation of state can be presented as: $(p + a/Vm2)(Vm-b) = RT$ 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 non-zero gas volume and intermolecular interactions.
• ### The Nernst Equation

• The two (ultimately equivalent) equations for these two cases (half-cell, full cell) are as follows: Ered = Estd,red - RT/zF ln ared/aox (half-cell reduction potential) Ecell = Estd,cell - RT/zF ln Q (total cell potential) where: Ered is the half-cell reduction potential at the temperature of interest Estd,red is the standard half-cell 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 half-reaction Q is the reaction quotient
• ### The Effect of Intermolecular Forces

• The ideal gas law is a convenient approximation for predicting the behavior of gas-phase 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: $(P+(a/V_m^2))(V_m-b) = RT$ Where b is the excluded volume, R is the universal gas constant, and T is the absolute temperature.