In chemistry, an ideal solution or ideal mixture is a Solution (chemistry), solution that exhibits thermodynamic properties analogous to those of a mixture of ideal gases. The enthalpy of mixing is zero as is the volume change on mixing by definition; the closer to zero the enthalpy of mixing is, the more "ideal" the behavior of the solution becomes. The vapor pressures of the solvent and solute obey Raoult's law and Henry's law, respectively, and the activity coefficient (which measures deviation from ideality) is equal to one for each component. The concept of an ideal solution is fundamental to chemical thermodynamics and its applications, such as the explanation of colligative properties.

Physical origin

Ideality of solutions is analogous to ideal gas, 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 intermolecular force, 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., n-butanol, 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 only 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 obeys Raoult's law

p_i=x_ip_i^*

for all compositions. Here

p_i

is the vapor pressure of component

i

above the solution,

x_i

is its mole fraction and

p_i^*

is the vapor pressure of the pure substance

i

at the same temperature. This definition depends on vapor pressure, which is a directly measurable property, at least for volatile components. The thermodynamic properties may then be obtained from the chemical potential μ (which is the partial molar property, partial molar Gibbs energy ''g'') of each component. If the vapor is an ideal gas, :

\mu(T,p_i) = g(T,p_i)=g^\mathrm(T,p^u)+RT\ln .

The reference pressure

p^u

may be taken as

P^o

= 1 bar, or as the pressure of the mix, whichever is simpler. On substituting the value of

p_i

from Raoult's law, :

\mu(T,p_i) =g^\mathrm(T,p^u)+RT\ln  + 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

f_i = x_i f_i^*

. Here

f_i

is the fugacity of component

i

in solution and

f_i^*

is the fugacity of

i

as a pure substance. Since the fugacity is defined by the equation :

\mu(T,P) = g(T,P)=g^\mathrm(T,p^u)+RT\ln

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 (chemistry), activity instead of fugacity.P.A. Rock, ''Chemical Thermodynamics: Principles and Applications'' (Macmillan 1969), p.261

Thermodynamic properties

Volume

If we differentiate this last equation with respect to

p

T

constant we get: :

\left(\frac\right)_=RT\left(\frac\right)_.

Since we know from the Gibbs potential equation that: :

\left(\frac\right)_=v

with the molar volume

v

, these last two equations put together give: :

\left(\frac\right)_=\frac.

Since all this, done as a pure substance, is valid in an ideal mix just adding the subscript

i

to all the intensive variables and changing

v

\bar

, with optional overbar, standing for partial molar volume: :

\left(\frac\right)_=\frac.

Applying the first equation of this section to this last equation we find: :

v_i^* = \bar_i

which means that the partial molar volumes in an ideal mix are independent of composition. Consequently, the total volume is the sum of the volumes of the components in their pure forms: :

V = \sum_i V_i^*.

Enthalpy and heat capacity

Proceeding in a similar way but taking the derivative with respect to

T

we get a similar result for molar enthalpy, enthalpies: :

\frac=\ln\frac.

Remembering that

\left( \frac\right)_P=-\frac

we get: :

-\frac=-\frac

which in turn means that

\bar=h_i^*

and that the enthalpy of the mix is equal to the sum of its component enthalpies. Since

\bar=\bar-p\bar

and

u_i^* = h_i^* - p v_i^*

, similarly :

u_i^*=\bar.

It is also easily verifiable that :

C_^*=\bar.

Entropy of mixing

Finally since :

\bar=\mu _i=g_i^\mathrm+RT\ln \frac=g_i^\mathrm+RT\ln \frac+RT\ln x_i=\mu _i^*+ RT\ln x_i

we find that :

\Delta g_=RT\ln x_i.

Since the Gibbs free energy per mole of the mixture

G_m

G_m = \sum_i x_i

then :

\Delta G_\mathrm=RT\sum_i.

At last we can calculate the molar entropy of mixing since

g_i^*=h_i^*-Ts_i^*

and

\bar=\bar-T\bar

\Delta s_=-R\sum _i \ln x_i

\Delta S_\mathrm=-R\sum _i x_i\ln x_i.

Consequences

Solvent–solute interactions are the same as solute–solute and solvent–solvent interactions, on average. Consequently, the enthalpy of mixing (solution) is zero and 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 :

\Delta G_ = RT \sum_i x_i \ln x_i

or for a two-component ideal solution :

\Delta G_ = 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

x_i

is the mole fraction of component

i

. Note that this free energy of mixing is always negative (since each

x_i \in [0,1]

, each

\ln x_i

or its limit for

x_i \to 0

must be negative (infinite)), i.e., ''ideal solutions are miscible at any composition'' and no phase separation will occur. The equation above can be expressed in terms of chemical potentials of the individual components :

\Delta G_ = \sum_i x_i \Delta\mu_

where

\Delta\mu_=RT\ln x_i

is the change in chemical potential of

i

on mixing. If the chemical potential of pure liquid

i

is denoted

\mu_i^*

, then the chemical potential of

i

in an ideal solution is :

\mu_i = \mu_i^* + RT \ln x_i.

Any component

i

of an ideal solution obeys Raoult's Law over the entire composition range: :

\ p_=(p_)_\text x_i

where

(p_i)_\text

is the equilibrium vapor pressure of pure component

i

and

x_i\,

is the mole fraction of component

i

in solution.

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 solution, 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.

References

{{Chemical solutions Solutions Thermodynamics Chemical thermodynamics