![]() It does not represent a natural zero point of any scale. It represents a non-zero standard deliberately chosen to be so (otherwise c/c o does not exist). ![]() In any case, the nature of o must be specified. (Obviously one might have to convert if comparing with someone else's numbers that are based on a more obvious choice like everything in molality). This does lead to a legitimate choice of standard state and a consistent set of values. If more than one concentration is to be taken into account, it is in principle possible to do one in molality, the other in mole fraction etc. When working in molalities this becomes P=1 bar, T= temperature of interest, concentration = 1 mol/kg and in mole fractions P=1 bar, T= temperature of interest, pure compound. It also specifies what standard state o is taken: P=1 bar, T= temperature of interest, concentration = 1 mol/l. l -1 and is present to cancel the dimensions of the formula.If we (somewhat unwisely) choose to work in molarities the C o factor is equal to 1 mol Where γ i is the coefficient of activity of the species i, C i is concentration in one of its many measures. It is equal to 1 bar and is present to cancel the dimensions of the formula. The value p 0 is the standard atmospheric pressure. The quantity f i has the dimension of pressure and is fugacity of the gas: for a pure Ideal gas, the fugacity coefficient is equal to one. Where γ i is the dimensionless fugacity coefficient ( fugacity) of the species i, x i the mole fraction of the species in the gaseous mixture and p the pressure expressed in bar. Molalities do not depend on volume, molarities do. Molalities are often preferred in this type of calculation because the volumes of non-ideal mixtures are strictly spoken not additive. This implies that the value of μ o will be different under these different reference systems. We can safely use such definitions but the meaning of the plimsoll superscript o (by lack of better represented as a Θ above) will differ, as we choose a different reference point. The same holds if we use molarities or molalities as concentration measures rather than mole fractions. However, the definition leads to a consistent system of values as long as we do not make x 1 too large. It represents the extrapolation of the Henry line all the way to x 1 =1 and this is a virtual state: a state of the system that can never be achieved. Raoult: P solvent = P* solvx 2 Henry: P solute = K solux 1Ī 2 ≡ P solvent/ P* solvent a 1 ≡ P solute/ K soluteĪs K solute ≠ P* solute this implies the choice of a different reference point. It is often convenient to use these laws as a basis for a definition: For dilute solutions usually the solvent follows Raoult's law, but the solute follows Henry's law. ![]() However there are many alternative schemes to define activity. We can also define an activity coefficient γ: We can use this formula to define an activity a, by insisting that the same formulism continues to hold in a non-ideal case. In ideal mixtures we can write the dependency of the thermodynamic potential on the composition (written as mole fraction x) as: ![]()
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