Ideal solution

In chemistry, an ideal solution or ideal mixture is a solution in which the gas phase exhibits thermodynamic properties analogous to those of a mixture of ideal gases.[1] The enthalpy of mixing is zero[2] as is the volume change on mixing by definition; the closer to zero the enthalpy of mixing is, the more "ideal" the behaviour of the solution becomes. The vapor pressure of the solution obeys either Raoult's law or Henry's law (or both),[3] and the activity coefficient of each component (which measures deviation from ideality) is equal to one.[4]

The concept of an ideal solution is fundamental to chemical thermodynamics and its applications, such as the use of colligative properties.

Physical origin

Ideality of solutions is analogous to ideality for gases, with the important difference that intermolecular interactions in liquids are strong and cannot simply be neglected as they can for ideal gases. Instead we assume that the mean strength of the interactions are the same between all the molecules of the solution.

More formally, for a mix of molecules of A and B, then the interactions between unlike neighbors (U_AB) and like neighbors U_AA and U_BB must be of the same average strength, i.e., 2 U_AB = U_AA + U_BB and the longer-range interactions must be nil (or at least indistinguishable). If the molecular forces are the same between AA, AB and BB, i.e., U_AB = U_AA = U_BB, then the solution is automatically ideal.

If the molecules are almost identical chemically, e.g., 1-butanol and 2-butanol, then the solution will be almost ideal. Since the interaction energies between A and B are almost equal, it follows that there is a very small overall energy (enthalpy) change when the substances are mixed. The more dissimilar the nature of A and B, the more strongly the solution is expected to deviate from ideality.

Formal definition

Different related definitions of an ideal solution have been proposed. The simplest definition is that an ideal solution is a solution for which each component (i) obeys Raoult's law ${\displaystyle p_{i}=x_{i}p_{i}^{*}}$ for all compositions. Here ${\displaystyle p_{i}}$ is the vapor pressure of component i above the solution, ${\displaystyle x_{i}}$ is its mole fraction and ${\displaystyle p_{i}^{*}}$ is the vapor pressure of the pure substance i at the same temperature.[5][6][7]

This definition depends on vapor pressures which are a directly measurable property, at least for volatile components. The thermodynamic properties may then be obtained from the chemical potential μ (or partial molar Gibbs energy g) of each component, which is assumed to be given by the ideal gas formula

${\displaystyle \mu (T,p_{i})=g(T,p_{i})=g^{\mathrm {u} }(T,p^{u})+RT\ln {\frac {p_{i}}{p^{u}}}}$.

The reference pressure ${\displaystyle p^{u}}$ may be taken as ${\displaystyle P^{0}}$ = 1 bar, or as the pressure of the mix to ease operations.

On substituting the value of ${\displaystyle p_{i}}$ from Raoult's law,

${\displaystyle \mu (T,p_{i})=g^{\mathrm {u} }(T,p^{u})+RT\ln {\frac {p_{i}^{*}}{p^{u}}}+RT\ln x_{i}=\mu _{i}^{*}+RT\ln x_{i}}$.

This equation for the chemical potential can be used as an alternate definition for an ideal solution.

However, the vapor above the solution may not actually behave as a mixture of ideal gases. Some authors therefore define an ideal solution as one for which each component obeys the fugacity analogue of Raoult's law ${\displaystyle f_{i}=x_{i}f_{i}^{*}}$,

Here ${\displaystyle f_{i}}$ is the fugacity of component ${\displaystyle i}$ in solution and ${\displaystyle f_{i}^{*}}$ is the fugacity of ${\displaystyle i}$ as a pure substance.[8][9] Since the fugacity is defined by the equation

${\displaystyle \mu (T,P)=g(T,P)=g^{\mathrm {u} }(T,p^{u})+RT\ln {\frac {f_{i}}{p^{u}}}}$

this definition leads to ideal values of the chemical potential and other thermodynamic properties even when the component vapors above the solution are not ideal gases. An equivalent statement uses thermodynamic activity instead of fugacity.[10]

Thermodynamic properties

Volume

If we differentiate this last equation with respect to ${\displaystyle P}$ at ${\displaystyle T}$ constant we get:

${\displaystyle \left({\frac {\partial g(T,P)}{\partial P}}\right)_{T}=RT\left({\frac {\partial \ln f}{\partial P}}\right)_{T}}$

but we know from the Gibbs potential equation that:

${\displaystyle \left({\frac {\partial g(T,P)}{\partial P}}\right)_{T}=v}$

These last two equations put together give:

${\displaystyle \left({\frac {\partial \ln f}{\partial P}}\right)_{T}={\frac {v}{RT}}}$

Since all this, done as a pure substance is valid in a mix just adding the subscript ${\displaystyle i}$ to all the intensive variables and changing ${\displaystyle v}$ to ${\displaystyle {\bar {v_{i}}}}$, standing for Partial molar volume.

${\displaystyle \left({\frac {\partial \ln f_{i}}{\partial P}}\right)_{T,x_{i}}={\frac {\bar {v_{i}}}{RT}}}$

Applying the first equation of this section to this last equation we get

${\displaystyle v_{i}^{*}={\bar {v_{i}}}}$

which means that in an ideal mix the volume is the addition of the volumes of its components:

${\displaystyle V=\sum V_{i}}$

Enthalpy and heat capacity

Proceeding in a similar way but derivative with respect of ${\displaystyle T}$ we get to a similar result with enthalpies

${\displaystyle {\frac {g(T,P)-g^{\mathrm {gas} }(T,p^{u})}{RT}}=\ln {\frac {f}{p^{u}}}}$

derivative with respect to T and remembering that ${\displaystyle \left({\frac {\partial {\frac {g}{T}}}{\partial T}}\right)_{P}=-{\frac {h}{T^{2}}}}$ we get:

${\displaystyle -{\frac {{\bar {h_{i}}}-h_{i}^{\mathrm {gas} }}{R}}=-{\frac {h_{i}^{*}-h_{i}^{\mathrm {gas} }}{R}}}$

which in turn is ${\displaystyle {\bar {h_{i}}}=h_{i}^{*}}$.

Meaning that the enthalpy of the mix is equal to the sum of its components.

Since ${\displaystyle {\bar {u_{i}}}={\bar {h_{i}}}-p{\bar {v_{i}}}}$ and ${\displaystyle u_{i}^{*}=h_{i}^{*}-pv_{i}^{*}}$:

${\displaystyle u_{i}^{*}={\bar {u_{i}}}}$

It is also easily verifiable that

${\displaystyle C_{pi}^{*}={\bar {C_{pi}}}}$

Entropy of mixing

Finally since

${\displaystyle {\bar {g_{i}}}=\mu _{i}=g_{i}^{\mathrm {gas} }+RT\ln {\frac {f_{i}}{p^{u}}}=g_{i}^{\mathrm {gas} }+RT\ln {\frac {f_{i}^{*}}{p^{u}}}+RT\ln x_{i}=\mu _{i}^{*}+RT\ln x_{i}}$

Which means that

${\displaystyle \Delta g_{i,\mathrm {mix} }=RT\ln x_{i}}$

and since the Gibbs free energy per mole of the mixture ${\displaystyle G_{m}}$ is

${\displaystyle G_{m}=\sum _{i}x_{i}{g_{i}}}$

then

${\displaystyle \Delta G_{\mathrm {m,mix} }=RT\sum _{i}{x_{i}\ln x_{i}}}$

At last we can calculate the molar entropy of mixing since ${\displaystyle g_{i}^{*}=h_{i}^{*}-Ts_{i}^{*}}$ and ${\displaystyle {\bar {g_{i}}}={\bar {h_{i}}}-T{\bar {s_{i}}}}$

${\displaystyle \Delta s_{i,\mathrm {mix} }=-R\sum _{i}\ln x_{i}}$

${\displaystyle \Delta S_{\mathrm {m,mix} }=-R\sum _{i}x_{i}\ln x_{i}}$

Consequences

Solvent-solute interactions are similar to solute-solute and solvent-solvent interactions

Since the enthalpy of mixing (solution) is zero, the change in Gibbs free energy on mixing is determined solely by the entropy of mixing. Hence the molar Gibbs free energy of mixing is

${\displaystyle \Delta G_{\mathrm {m,mix} }=RT\sum _{i}x_{i}\ln x_{i}}$

or for a two component solution

${\displaystyle \Delta G_{\mathrm {m,mix} }=RT(x_{A}\ln x_{A}+x_{B}\ln x_{B})}$

where m denotes molar, i.e., change in Gibbs free energy per mole of solution, and ${\displaystyle x_{i}}$ is the mole fraction of component ${\displaystyle i}$.

Note that this free energy of mixing is always negative (since each ${\displaystyle x_{i}\in [0,1]}$, each ${\displaystyle \ln x_{i}}$ or its limit for ${\displaystyle x_{i}\rightarrow 0}$ must be negative (infinite)), i.e., ideal solutions are always completely miscible.

The equation above can be expressed in terms of chemical potentials of the individual components

${\displaystyle \Delta G_{\mathrm {m,mix} }=\sum _{i}x_{i}\Delta \mu _{i,\mathrm {mix} }}$

where ${\displaystyle \Delta \mu _{i,\mathrm {mix} }=RT\ln x_{i}}$ is the change in chemical potential of ${\displaystyle i}$ on mixing.

If the chemical potential of pure liquid ${\displaystyle i}$ is denoted ${\displaystyle \mu _{i}^{*}}$, then the chemical potential of ${\displaystyle i}$ in an ideal solution is

${\displaystyle \mu _{i}=\mu _{i}^{*}+RT\ln x_{i}}$

Any component ${\displaystyle i}$ of an ideal solution obeys Raoult's Law over the entire composition range:

${\displaystyle \ P_{i}=(P_{i})_{pure}x_{i}}$

where

${\displaystyle (P_{i})_{pure}\,}$ is the equilibrium vapor pressure of the pure component

${\displaystyle x_{i}\,}$ is the mole fraction of the component in solution

It can also be shown that volumes are strictly additive for ideal solutions.

Non-ideality

Deviations from ideality can be described by the use of Margules functions or activity coefficients. A single Margules parameter may be sufficient to describe the properties of the solution if the deviations from ideality are modest; such solutions are termed regular.

In contrast to ideal solutions, where volumes are strictly additive and mixing is always complete, the volume of a non-ideal solution is not, in general, the simple sum of the volumes of the component pure liquids and solubility is not guaranteed over the whole composition range. By measurement of densities thermodynamic activity of components can be determined.

See also

• Activity coefficient
• Entropy of mixing
• Margules function
• Regular solution
• Coil-globule transition
• Apparent molar property
• Dilution equation
• Virial coefficient

References

1. Felder, Richard M.; Rousseau, Ronald W.; Bullard, Lisa G. (2005). Elementary Principles of Chemical Processes. Wiley. p. 293. ISBN 978-0471687573.
2. A to Z of Thermodynamics Pierre Perrot ISBN 0-19-856556-9
3. Felder, Richard M.; Rousseau, Ronald W.; Bullard, Lisa G. Elementary Principles of Chemical Processes. Wiley. p. 293. ISBN 978-0471687573.
4. IUPAC, Compendium of Chemical Terminology, 2nd ed. (the "Gold Book") (1997). Online corrected version:  (2006–) "ideal mixture". doi:10.1351/goldbook.I02938
5. P. Atkins and J. de Paula, Atkins’ Physical Chemistry (8th edn, W.H.Freeman 2006), p.144
6. T. Engel and P. Reid Physical Chemistry (Pearson 2006), p.194
7. K.J. Laidler and J.H. Meiser Physical Chemistry (Benjamin-Cummings 1982), p.180
8. R.S. Berry, S.A. Rice and J. Ross, Physical Chemistry (Wiley 1980) p.750
9. I.M. Klotz, Chemical Thermodynamics (Benjamin 1964) p.322
10. P.A. Rock, Chemical Thermodynamics: Principles and Applications (Macmillan 1969), p.261