Ideal mixture

Ideal solution

In chemistry, an ideal solution or ideal mixture is a solution in which the enthalpy of solution is zero; the closer to zero the enthalpy of solution, the more "ideal" the behavior of the solution becomes. Equivalently, an ideal mixture is one in which the activity coefficients (which measure deviation from ideality) are equal to one.

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 can not 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, the interactions between unlike neighbors (U_AB) and like neighbors U_AA and U_BB must be of the same average strength i.e. 2U_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 ideal. Since the interaction energies between A and B are the same, it follows that there is no 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

An ideal mix is defined as a mix that satisfies:

f_i=x_if_i^*

where

f_i

is the fugacity of component

i

and

f_i^*

is the fugacity of

i

as a pure substance.

Since the definition of fugacity in a pure substance is:

g(T,P)=g^mathrm{gas}(T,p^u)+RTln {frac{f}{p^u}}

Where

g^mathrm{gas}(T,p^u)

is the molar free energy of an ideal gas at a temperature

T

and a reference presure

P^u

which might be taken as

P^0

or the presure of the mix to ease operations.

If we derivative this last equation with respect to $P$ at $T$ constant we get:

left(frac{partial g(T,P)}{partial P}right)_{T}=RTleft(frac{partial ln f}{partial P}right)_{T}

but we know from the Gibbs potential equation that:

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

These last two equations put together give:

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 $i$ to all the intensive variables and changing $v$ to $bar{v_i}$ , standing for Partial molar volume.

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

v_i^*=bar{v_i}

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

Prociding in a similar way but derivative with respect of $T$ we get to a similar result with enthalpies

frac{g(T,P)-g^mathrm{gas}(T,p^u)}{RT}=lnfrac{f}{p^u}

derivative with respect to T ang remembering that

left(frac{partial frac{g}{T}}{partial T}right)_P=-frac{h}{T^2}

we get:

-frac{bar{h_i}-h_i^mathrm{gas}}{R}=-frac{h_i^*-h_i^mathrm{gas}}{R}

whitch in turn is

bar{h_i}=h_i^*

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

Since $bar{u_i}=bar{h_i}-pbar{v_i}$ and $u_i^*=h_i^*-pv_i^*$ :

u_i^*=bar{u_i}

It is also easily verifiable that

C_{pi}^*=bar{C_{pi}}

Finaly since

bar{g_i}=mu _i=g_i^mathrm{gas}+RTln frac{f_i}{p^u}=g_i^mathrm{gas}+RTln frac{f_i^*}{P^u}+RTln x_i=mu _i^*+ RTln x_i

Which means that

Delta g_{i,mathrm{mix}}=RTln x_i

and since

$G=sum_i x_i{g_i}$

then

Delta G_mathrm{mix}=RTsum_i{x_iln x_i}

At last we can calculate the entropy of mixing since $g_i^*=h_i^*-Ts_i^*$ and $bar{g_i}=bar{h_i}-Tbar{s_i}$

Delta s_{i,mathrm{mix}}=-Rsum _i ln x_i

Delta S_mathrm{mix}=-Rsum _i x_iln x_i

Consequences

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

Delta G_{mathrm{m,mix}} = RT sum_i x_i ln x_i

or for a two component solution

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

x_i

is the mole fraction of component

i

Note that this free energy of mixing is always negative (since each $x_i$ is positive and each $ln x_i$ must be negative) i.e. ideal solutions are always completely miscible.

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

Delta G_{mathrm{m,mix}} = sum_i x_i Deltamu_{i,mathrm{mix}}

where

Deltamu_{i,mathrm{mix}}=RTln 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^* + Delta mu_{i,mathrm{mix}} = mu_i^* + RT ln x_i

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

P_{i}=(P_{i})_{pure} x_i

where

(P_i)_{pure},

is the equilibrium vapor pressure of the pure component

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.