# Enthalpy of mixing

The enthalpy of mixing (or heat of mixing) is the enthalpy liberated or absorbed from a substance upon mixing.[1] When a substance or compound is combined with any other substance or compound the enthalpy of mixing is the consequence of the new interactions between the two substances or compounds.[1] This enthalpy if released exothermically can in an extreme case cause an explosion.

Enthalpy of mixing can often be ignored in calculations for mixtures where other heat terms exist, or in cases where the mixture is ideal.[2] The sign convention is the same as for enthalpy of reaction: when the enthalpy of mixing is positive, mixing is endothermic while negative enthalpy of mixing signifies exothermic mixing. In ideal mixtures the enthalpy of mixing is null. In non-ideal mixtures the thermodynamic activity of each component is different from its concentration by multiplying with the activity coefficient.

One approximation for calculating the heat of mixing is Flory–Huggins solution theory for polymer solutions.

## Formal Definition

For a liquid, enthalpy of mixing can be defined as follows[2]

${\displaystyle H_{mixture}=\Delta H_{mix}+\sum X_{i}H_{i}}$

Where:

• Hmixture is the total enthalpy of the system after mixing
• ΔHmix is the enthalpy of mixing
• xi is the mole fraction of component i in the system
• Hi is the enthalpy of pure i

Enthalpy of mixing can also be defined using Gibbs free energy of mixing

${\displaystyle \Delta G_{mix}=\Delta H_{mix}-T*\Delta S_{mix}}$

However, Gibbs free energy of mixing and entropy of mixing tend to be more difficult to determine experimentally.[3] As such, enthalpy of mixing tends to be determined experimentally in order to calculate entropy of mixing, rather than the reverse.

It should be noted that enthalpy of mixing is defined exclusively for the continuum regime, which excludes molecular-scale effects (However, first-principles calculations have been made for some metal-alloy systems such as Al-Co-Cr[4] or β-Ti[5]).

When two substances are mixed the resulting enthalpy is not an addition of the pure component enthalpies, unless the substances form an ideal mixture.[6] The interactions between each set of molecules determines the final change in enthalpy. For example, when compound “x” has a strong attractive interaction with compound “y” the resulting enthalpy is exothermic.[6] In the case of alcohol and its interactions with a hydrocarbon, the alcohol molecule participates in hydrogen bonding with other alcohol molecules, and these hydrogen bonding interactions are much stronger than alcohol-hydrocarbon interactions, which results in an endothermic heat of mixing.[7]

## Calculations

Enthalpy of mixing is often calculated experimentally using calorimetry methods. A bomb calorimeter is created to be an isolated system. With an insulated frame and a reaction chamber, a bomb calorimeter is used to transfer heat of a reaction or mixing into surrounding water which is then calculated for temperature. A typical solution would use the equation ${\displaystyle H_{mixture}=\Delta H_{mix}+\sum X_{i}H_{i}}$ (derived from the definition above) in conjunction experimentally determined total-mixture enthalpies and tabulated pure species enthalpies, the difference being equal to enthalpy of mixing.

More complex models, such as the Flory-Huggins and UNIFAC models, allow prediction of enthalpies of mixing. Flory-Huggins is useful in calculating enthalpies of mixing for polymeric mixtures and considers a system from a multiplicity perspective.

Calculations of organic enthalpies of mixing can be made by modifying UNIFAC using the equations[8]

• ${\displaystyle \Delta H_{mix}=\sum x_{i}{\overline {\Delta H_{i}}}}$
• ${\displaystyle {\overline {\Delta H_{i}}}=\sum _{k}N_{ki}(H_{k}-H_{ki}^{*})}$
• ${\displaystyle {H_{k} \over {RT^{2}}}=Q_{k}{\biggl (}{\sum _{m}{\theta \psi '_{mk}} \over {\sum _{m}{\theta \psi _{mk}}}}-{\biggl (}\sum _{m}{{\theta _{m}psi'_{k}m} \over {\sum _{n}\theta _{n}\psi _{nm}}}-{\theta _{m}\psi _{km}(\sum _{n}\theta _{n}\psi '_{nm}) \over {(\sum _{n}\theta _{n}\psi _{nm})^{2}}}{\biggr )}{\biggr )}}$

Where:

• ${\displaystyle x_{i}}$ = liquid mole fraction of i
• ${\displaystyle {\overline {\Delta H_{i}}}}$ = partial molar excess enthalpy of i
• ${\displaystyle N_{ki}}$ = number of groups of type k in i
• ${\displaystyle H_{k}}$ = excess enthalpy of group k
• ${\displaystyle H_{ki}^{*}}$ = excess enthalpy of group k in pure i
• ${\displaystyle Q_{k}}$ = area parameter of group k
• ${\displaystyle \theta _{m}={Q_{m}X_{m} \over \sum _{n}Q_{n}X_{n}}}$ = area fraction of group m
• ${\displaystyle X_{m}={\sum _{i}x_{i}N_{mi} \over \sum _{i}x_{i}\sum _{k}N_{ki}}}$ = mole fraction of group m in the mixture
• ${\displaystyle \psi _{mn}=exp{\biggl (}-{Za_{mn} \over 2T}{\biggr )}}$
• ${\displaystyle \psi _{mn}^{*}={\delta \over \delta T}({\psi _{m}n})}$
• ${\displaystyle Z=35.2-0.1272T+0.00014T^{2}}$ = Temperature dependant coordination number

It can be seen that prediction of enthalpy of mixing is incredibly complex and requires a plethora of system variables to be known. This explains why enthalpy of mixing is typically experimentally determined.

## Ideal and Regular Mixtures

An ideal mixture is any in which the arithmetic mean (with respect to mole fraction) of the two pure substances is the same as that of the final mixture. Among other important thermodynamic simplifications, this means that enthalpy of mixing is zero: ${\displaystyle \Delta H_{mix,ideal}=0}$. Any gas that follows the ideal gas law can be assumed to mix ideally, as can hydrocarbons and liquids with similar molecular interactions and properties.[2]

A regular solution or mixture has a non-zero enthalpy of mixing with an ideal entropy of mixing.[9][10] Under this assumption, ${\displaystyle \Delta H_{mix}}$ scales linearly with ${\displaystyle X_{1}X_{2}}$, and is equivalent to the excess internal energy.[11]

## Intermolecular Forces

Intermolecular forces are the main constituent of changes in the enthalpy of a mixture. Stronger attractive forces between the mixed molecules, such as hydrogen-bonding, induced-dipole, and dipole-dipole interactions result in a lower enthalpy of the mixture and a release of heat.[6] If strong interactions only exist between like-molecules, such as H-bonds between water in a water-hexane solution, the mixture will have a higher total enthalpy and absorb heat.

## References

1. ^ a b Carlson, Phillip (2002). Hazardous Chemicals Handbook (2nd edition). Elsevier. p. 52. ISBN 978-0-7506-4888-2.
2. ^ a b c Sinnot, Ray K (2009). Chemical Engineering Design - SI Edition (5th Edition). Elsevier. p. 95. ISBN 978-0-7506-8551-1 – via Knovel.
3. ^ Lin, Shu-Kun (1996). "Gibbs paradox of entropy of mixing: experimental facts, its rejection and the theoretical consequences" (PDF). ELECTRONIC JOURNAL OF THEORETICAL CHEMISTRY. 1: 135–150.
4. ^ Liu, Xuan L.; Gheno, Thomas; Lindahl, Bonnie B.; Lindwall, Greta; Gleeson, Brian; Liu, Zi-Kui (2015-04-13). "First-Principles Calculations, Experimental Study, and Thermodynamic Modeling of the Al-Co-Cr System". PLOS ONE. 10 (4): e0121386. doi:10.1371/journal.pone.0121386. ISSN 1932-6203. PMC . PMID 25875037.
5. ^ Chandran, Mahesh; Subramanian, P. R.; Gigliotti, Michael F. (2013-02-15). "First principles calculation of mixing enthalpy of β-Ti with transition elements". Journal of Alloys and Compounds. 550: 501–508. doi:10.1016/j.jallcom.2012.10.141.
6. ^ a b c CB,378-2590,224-2707, Richard Rowley,350. "Heat_of_Mixing". www.et.byu.edu. Retrieved 2017-02-22.
7. ^ Savini, C. G.; Winterhalter, D. R.; Kovach, L. H.; Van Ness, H. C. (1966-01-01). "Endothermic Heats of Mixing by Isothermal Dilution Calorimetry.". Journal of Chemical & Engineering Data. 11 (1): 40–43. doi:10.1021/je60028a009. ISSN 0021-9568.
8. ^ Dang, Dinh; Tassios, Dimitrios P. (1986-01-01). "Prediction of enthalpies of mixing with a UNIFAC model". Industrial & Engineering Chemistry Process Design and Development. 25 (1): 22–31. doi:10.1021/i200032a004. ISSN 0196-4305.
9. ^ P. Atkins and J. de Paula, Atkins' Physical Chemistry (8th ed. W.H. Freeman 2006) p.149
10. ^ P.A. Rock, Chemical Thermodynamics. Principles and Applications (Macmillan 1969) p.263
11. ^ Vidal, Jean (2003). Thermodynamics - Applications in Chemical Engineering and the Petroleum Industry. Editions Technip. p. 232. ISBN 978-2-7108-0800-8.