Why there is depression in freezing point when a non-volatile solute is dissolved in a volatile solvent?

Freezing point depression is a colligative property observed in solutions that results from the introduction of solute molecules to a solvent. The freezing points of solutions are all lower than that of the pure solvent and is directly proportional to the molality of the solute.

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Due to crystal defect
Laboratory uses
For concentrated solution

\[\Delta{T_f} = T_f(solvent) - T_f (solution) = K_f \times m\]

where $\Delta{T_f}$ is the freezing point depression, $T_f$ (solution) is the freezing point of the solution, $T_f$ (solvent) is the freezing point of the solvent, $K_f$ is the freezing point depression constant, and m is the molality.

Nonelectrolytes are substances with no ions, only molecules. Strong electrolytes, on the other hand, are composed mostly of ionic compounds, and essentially all soluble ionic compounds form electrolytes. Therefore, if we can establish that the substance that we are working with is uniform and is not ionic, it is safe to assume that we are working with a nonelectrolyte, and we may attempt to solve this problem using our formulas. This will most likely be the case for all problems you encounter related to freezing point depression and boiling point elevation in this course, but it is a good idea to keep an eye out for ions. It is worth mentioning that these equations work for both volatile and nonvolatile solutions. This means that for the sake of determining freezing point depression or boiling point elevation, the vapor pressure does not effect the change in temperature. Also, remember that a pure solvent is a solution that has had nothing extra added to it or dissolved in it. We will be comparing the properties of that pure solvent with its new properties when added to a solution.

Adding solutes to an ideal solution results in a positive ΔS, an increase in entropy. Because of this, the newly altered solution's chemical and physical properties will also change. The properties that undergo changes due to the addition of solutes to a solvent are known as colligative properties. These properties are dependent on the number of solutes added, not on their identity. Two examples of colligative properties are boiling point and freezing point: due to the addition of solutes, the boiling point tends to increase, and freezing point tends to decrease.

The freezing point and boiling point of a pure solvent can be changed when added to a solution. When this occurs, the freezing point of the pure solvent may become lower, and the boiling point may become higher. The extent to which these changes occur can be found using the formulas:

\[\Delta{T}_f = -K_f \times m\]

\[\Delta{T}_f = K_b \times m\]

where $m$ is the solute molality and $K$ values are proportionality constants; ($K_f$ and $K_b$ for freezing and boiling, respectively.

If solving for the proportionality constant is not the ultimate goal of the problem, these values will most likely be given. Some common values for $K_f$ and $K_b$ respectively, are:

Solvent	$K_f$	$K_b$
Water	1.86	.512
Acetic acid	3.90	3.07
Benzene	5.12	2.53
Phenol	7.27	3.56

Molality is defined as the number of moles of solute per kilogram solvent. Be careful not to use the mass of the entire solution. Often, the problem will give you the change in temperature and the proportionality constant, and you must find the molality first in order to get your final answer.

The solute, in order for it to exert any change on colligative properties, must fulfill two conditions. First, it must not contribute to the vapor pressure of the solution, and second, it must remain suspended in the solution even during phase changes. Because the solvent is no longer pure with the addition of solutes, we can say that the chemical potential of the solvent is lower. Chemical potential is the molar Gibb's energy that one mole of solvent is able to contribute to a mixture. The higher the chemical potential of a solvent is, the more it is able to drive the reaction forward. Consequently, solvents with higher chemical potentials will also have higher vapor pressures.

The boiling point is reached when the chemical potential of the pure solvent, a liquid, reaches that of the chemical potential of pure vapor. Because of the decrease in the chemical potential of mixed solvents and solutes, we observe this intersection at higher temperatures. In other words, the boiling point of the impure solvent will be at a higher temperature than that of the pure liquid solvent. Thus, boiling point elevation occurs with a temperature increase that is quantified using

\[\Delta{T_b} = K_b b_B\]

where

$K_b$ is known as the ebullioscopic constant and
$m$ is the molality of the solute.

Freezing point is reached when the chemical potential of the pure liquid solvent reaches that of the pure solid solvent. Again, since we are dealing with mixtures with decreased chemical potential, we expect the freezing point to change. Unlike the boiling point, the chemical potential of the impure solvent requires a colder temperature for it to reach the chemical potential of the pure solid solvent. Therefore, a freezing point depression is observed.

Example $\PageIndex{1}$

2.00 g of some unknown compound reduces the freezing point of 75.00 g of benzene from 5.53 to 4.90 $^{\circ}C$. What is the molar mass of the compound?

Solution

First we must compute the molality of the benzene solution, which will allow us to find the number of moles of solute dissolved.

\[ \begin{align*} m &= \dfrac{\Delta{T}_f}{-K_f} \\[4pt] &= \dfrac{(4.90 - 5.53)^{\circ}C}{-5.12^{\circ}C / m} \\[4pt] &= 0.123 m \end{align*}\]

\[ \begin{align*} \text{Amount Solute} &= 0.07500 \; kg \; benzene \times \dfrac{0.123 \; m}{1 \; kg \; benzene} \\[4pt] &= 0.00923 \; m \; solute \end{align*}\]

We can now find the molecular weight of the unknown compound:

\[ \begin{align*} \text{Molecular Weight} =& \dfrac{2.00 \; g \; unknown}{0.00923 \; mol} \\[4pt] &= 216.80 \; g/mol \end{align*}\]

The freezing point depression is especially vital to aquatic life. Since saltwater will freeze at colder temperatures, organisms can survive in these bodies of water.

Road salting takes advantage of this effect to lower the freezing point of the ice it is placed on. Lowering the freezing point allows the street ice to melt at lower temperatures. The maximum depression of the freezing point is about −18 °C (0 °F), so if the ambient temperature is lower, $\ce{NaCl}$ will be ineffective. Under these conditions, $\ce{CaCl_2}$ can be used since it dissolves to make three ions instead of two for $\ce{NaCl}$.

Figure $\PageIndex{1}$: Workers manually spreading salt from a salt truck in Milwaukee, Wisconsin. from Wikipedia

Benzophenone has a freezing point of 49.00oC. A 0.450 molal solution of urea in this solvent has a freezing point of 44.59oC. Find the freezing point depression constant for the solvent.(answ.: 9.80oC/m)

Freezing-point depression is a drop in the temperature at which a substance freezes, caused when a smaller amount of another, non-volatile substance is added. Examples include adding salt into water (used in ice cream makers and for de-icing roads), alcohol in water, ethylene or propylene glycol in water (used in antifreeze in cars), adding copper to molten silver (used to make solder that flows at a lower temperature than the silver pieces being joined), or the mixing of two solids such as impurities into a finely powdered drug.

Workers spreading salt from a salt truck for deicing the road.

Freezing point depression is responsible for keeping ice cream soft below 0°C.[1]

In all cases, the substance added/present in smaller amounts is considered the solute, while the original substance present in larger quantity is thought of as the solvent. The resulting liquid solution or solid-solid mixture has a lower freezing point than the pure solvent or solid because the chemical potential of the solvent in the mixture is lower than that of the pure solvent, the difference between the two being proportional to the natural logarithm of the mole fraction. In a similar manner, the chemical potential of the vapor above the solution is lower than that above a pure solvent, which results in boiling-point elevation. Freezing-point depression is what causes sea water (a mixture of salt and other compounds in water) to remain liquid at temperatures below 0 °C (32 °F), the freezing point of pure water.

The freezing point is the temperature at which the liquid solvent and solid solvent are at equilibrium, so that their vapour pressures are equal. When a non-volatile solute is added to a volatile liquid solvent, the solution vapour pressure will be lower than that of the pure solvent. As a result, the solid will reach equilibrium with the solution at a lower temperature than with the pure solvent.[2] This explanation in terms of vapor pressure is equivalent to the argument based on chemical potential, since the chemical potential of a vapor is logarithmically related to pressure. All of the colligative properties result from a lowering of the chemical potential of the solvent in the presence of a solute. This lowering is an entropy effect. The greater randomness of the solution (as compared to the pure solvent) acts in opposition to freezing, so that a lower temperature must be reached, over a broader range, before equilibrium between the liquid solution and solid solution phases is achieved. Melting point determinations are commonly exploited in organic chemistry to aid in identifying substances and to ascertain their purity.

Due to crystal defect

Salt prevents the water molecules from solidifying into ice crystals at 0 °C (32 °F), instead staying slushy at that temperature, before eventually freezing around −9 °C (16 °F).[3]

Consider the problem in which the solvent freezes to a very nearly pure crystal, regardless of the presence of the nonvolatile solute. This typically occurs simply because the solute molecules do not fit well in the crystal, i.e. substituting a solute for a solvent molecule in the crystal has high enthalpy. In this case, for low solute concentrations, the freezing point depression depends solely on the concentration of solute particles, not on their individual properties. The freezing point depression thus is called a colligative property.[4]

The explanation for the freezing point depression is then simply that as solvent molecules leave the liquid and join the solid, they leave behind a smaller volume of liquid in which the solute particles can roam. The resulting reduced entropy of the solute particles thus is independent of their properties. This approximation ceases to hold when the concentration becomes large enough for solute-solute interactions to become important. In that case, the freezing point depression depends on particular properties of the solute other than its concentration.[citation needed]

The phenomenon of freezing-point depression has many practical uses. The radiator fluid in an automobile is a mixture of water and ethylene glycol. The freezing-point depression prevents radiators from freezing in winter. Road salting takes advantage of this effect to lower the freezing point of the ice it is placed on. Lowering the freezing point allows the street ice to melt at lower temperatures, preventing the accumulation of dangerous, slippery ice. Commonly used sodium chloride can depress the freezing point of water to about −21 °C (−6 °F). If the road surface temperature is lower, NaCl becomes ineffective and other salts are used, such as calcium chloride, magnesium chloride or a mixture of many. These salts are somewhat aggressive to metals, especially iron, so in airports safer media such as sodium formate, potassium formate, sodium acetate, potassium acetate are used instead.

Pre-treating roads with salt relies on the warmer road surface to initially melt the snow and make a solution; Pre-treatment of bridges (which are colder than roads) does not typically work.[5]

Freezing-point depression is used by some organisms that live in extreme cold. Such creatures have evolved means through which they can produce a high concentration of various compounds such as sorbitol and glycerol. This elevated concentration of solute decreases the freezing point of the water inside them, preventing the organism from freezing solid even as the water around them freezes, or as the air around them becomes very cold. Examples of organisms that produce antifreeze compounds include some species of arctic-living fish such as the rainbow smelt, which produces glycerol and other molecules to survive in frozen-over estuaries during the winter months.[6] In other animals, such as the spring peeper frog (Pseudacris crucifer), the molality is increased temporarily as a reaction to cold temperatures. In the case of the peeper frog, freezing temperatures trigger a large-scale breakdown of glycogen in the frog's liver and subsequent release of massive amounts of glucose into the blood.[7]

Conifers have concentrated cell sap that also act like antifreeze in winters.[8]

With the formula below, freezing-point depression can be used to measure the degree of dissociation or the molar mass of the solute. This kind of measurement is called cryoscopy (Greek cryo = cold, scopos = observe; "observe the cold"[9]) and relies on exact measurement of the freezing point. The degree of dissociation is measured by determining the van 't Hoff factor i by first determining mB and then comparing it to msolute. In this case, the molar mass of the solute must be known. The molar mass of a solute is determined by comparing mB with the amount of solute dissolved. In this case, i must be known, and the procedure is primarily useful for organic compounds using a nonpolar solvent. Cryoscopy is no longer as common a measurement method as it once was, but it was included in textbooks at the turn of the 20th century. As an example, it was still taught as a useful analytic procedure in Cohen's Practical Organic Chemistry of 1910,[10] in which the molar mass of naphthalene is determined using a Beckmann freezing apparatus.

Laboratory uses

Freezing-point depression can also be used as a purity analysis tool when analyzed by differential scanning calorimetry. The results obtained are in mol%, but the method has its place, where other methods of analysis fail.

In the laboratory, lauric acid may be used to investigate the molar mass of an unknown substance via the freezing-point depression. The choice of lauric acid is convenient because the melting point of the pure compound is relatively high (43.8 °C). Its cryoscopic constant is 3.9 °C·kg/mol. By melting lauric acid with the unknown substance, allowing it to cool, and recording the temperature at which the mixture freezes, the molar mass of the unknown compound may be determined.[11][citation needed]

This is also the same principle acting in the melting-point depression observed when the melting point of an impure solid mixture is measured with a melting-point apparatus since melting and freezing points both refer to the liquid-solid phase transition (albeit in different directions).

In principle, the boiling-point elevation and the freezing-point depression could be used interchangeably for this purpose. However, the cryoscopic constant is larger than the ebullioscopic constant, and the freezing point is often easier to measure with precision, which means measurements using the freezing-point depression are more precise.

This phenomenon is applicable in preparing a freezing mixture to make ice cream. For this purpose, NaCl or another salt is used to lower the melting point of ice.

FPD measurements are also used in the dairy industry to ensure that milk has not had extra water added. Milk with a FPD of over 0.509 °C is considered to be unadulterated.[12]

Freezing temperature of seawater at different pressures and some substances as a function of salinity. See image description for source.

If the solution is treated as an ideal solution, the extent of freezing-point depression depends only on the solute concentration that can be estimated by a simple linear relationship with the cryoscopic constant ("Blagden's Law").

Δ t ∝ Moles of dissolved species Weight of solvent {\displaystyle \Delta t\propto {\frac {\text{Moles of dissolved species}}{\text{Weight of solvent}}}}

Δ t = K f b i {\displaystyle \Delta t=K_{f}bi}

where:

$Δ t {\displaystyle \Delta t}$ is the decrease in freezing-point
$K f {\displaystyle K_{f}}$ , the cryoscopic constant, which is dependent on the properties of the solvent, not the solute. (Note: When conducting experiments, a higher k value makes it easier to observe larger drops in the freezing point.)
$b {\displaystyle b}$ is the molality (moles of solute per kilogram of solvent)
$i {\displaystyle i}$ is the van 't Hoff factor (number of ion particles per formula unit of solute, e.g. i = 2 for NaCl, 3 for BaCl2).

Some values of the cryoscopic constant Kf for selected solvents:[13]

Compound	Freezing point (°C)	Kf in K⋅kg/mol
Acetic acid	16.6	3.90
Benzene	5.5	5.12
Camphor	179.8	39.7
Carbon disulfide	−112	3.8
Carbon tetrachloride	−23	30
Chloroform	−63.5	4.68
Cyclohexane	6.4	20.2
Ethanol	−114.6	1.99
Ethyl ether	−116.2	1.79
Naphthalene	80.2	6.9
Phenol	41	7.27
Water	0	1.86[14]

For concentrated solution

The simple relation above doesn't consider the nature of the solute, so it is only effective in a diluted solution. For a more accurate calculation at a higher concentration, for ionic solutes, Ge and Wang (2010)[15][16] proposed a new equation:

Δ T F = Δ H T F fus − 2 R T F ⋅ ln ⁡ a liq − 2 Δ C p fus T F 2 R ⋅ ln ⁡ a liq + ( Δ H T F fus ) 2 2 ( Δ H T F fus T F + Δ C p fus 2 − R ⋅ ln ⁡ a liq ) . {\displaystyle \Delta T_{\text{F}}={\frac {\Delta H_{T_{\text{F}}}^{\text{fus}}-2RT_{\text{F}}\cdot \ln a_{\text{liq}}-{\sqrt {2\Delta C_{p}^{\text{fus}}T_{\text{F}}^{2}R\cdot \ln a_{\text{liq}}+(\Delta H_{T_{\text{F}}}^{\text{fus}})^{2}}}}{2\left({\frac {\Delta H_{T_{\text{F}}}^{\text{fus}}}{T_{\text{F}}}}+{\frac {\Delta C_{p}^{\text{fus}}}{2}}-R\cdot \ln a_{\text{liq}}\right)}}.}

In the above equation, TF is the normal freezing point of the pure solvent (273 K for water, for example); aliq is the activity of the solvent in the solution (water activity for aqueous solution); ΔHfusTF is the enthalpy change of fusion of the pure solvent at TF, which is 333.6 J/g for water at 273 K; ΔCfusp is the difference between the heat capacities of the liquid and solid phases at TF, which is 2.11 J/(g·K) for water.

The solvent activity can be calculated from the Pitzer model or modified TCPC model, which typically requires 3 adjustable parameters. For the TCPC model, these parameters are available[17][18][19][20] for many single salts.

Melting-point depression
Boiling-point elevation
Colligative properties
Deicing
Eutectic point
Frigorific mixture
List of boiling and freezing information of solvents
Snow removal

^ "Controlling the hardness of ice cream, gelato and similar frozen desserts". Food Science and Technology. 2021-03-18. doi:10.1002/fsat.3510_3.x. ISSN 1475-3324.
^ Petrucci, Ralph H.; Harwood, William S.; Herring, F. Geoffrey (2002). General Chemistry (8th ed.). Prentice-Hall. pp. 557–558. ISBN 0-13-014329-4.
^ Pollock, Julie. "Salt Doesn't Melt Ice—Here's How It Makes Winter Streets Safer". Scientific American.
^ Atkins, Peter (2006). Atkins' Physical Chemistry. Oxford University Press. pp. 150–153. ISBN 0198700725.
^ Pollock, Julie. "Salt Doesn't Melt Ice—Here's How It Makes Winter Streets Safer". Scientific American.
^ Treberg, J. R.; Wilson, C. E.; Richards, R. C.; Ewart, K. V.; Driedzic, W. R. (2002). "The freeze-avoidance response of smelt Osmerus mordax: initiation and subsequent suppression 6353". The Journal of Experimental Biology. 205 (Pt 10): 1419–1427. doi:10.1242/jeb.205.10.1419.
^ L. Sherwood et al., Animal Physiology: From Genes to Organisms, 2005, Thomson Brooks/Cole, Belmont, CA, ISBN 0-534-55404-0, p. 691–692.
^ Ray, C. Claiborne (2002-02-05). "Q & A". The New York Times. ISSN 0362-4331. Retrieved 2022-02-10.
^ BIOETYMOLOGY – Biomedical Terms of Greek Origin. cryoscopy. bioetymology.blogspot.com.
^ Cohen, Julius B. (1910). Practical Organic Chemistry. London: MacMillan and Co.
^ "Archived copy" (PDF). Archived from the original (PDF) on 2020-08-03. Retrieved 2019-07-08.{{cite web}}: CS1 maint: archived copy as title (link)
^ "Freezing Point Depression of Milk". Dairy UK. 2014. Archived from the original on 2014-02-23.
^ P. W. Atkins, Physical Chemistry, 4th Ed., p. C17 (Table 7.2)
^ Aylward, Gordon; Findlay, Tristan (2002), SI Chemical Data 5th ed. (5 ed.), Sweden: John Wiley & Sons, p. 202, ISBN 0-470-80044-5
^ Ge, Xinlei; Wang, Xidong (2009). "Estimation of Freezing Point Depression, Boiling Point Elevation, and Vaporization Enthalpies of Electrolyte Solutions". Industrial & Engineering Chemistry Research. 48 (10): 5123. doi:10.1021/ie900434h. ISSN 0888-5885.
^ Ge, Xinlei; Wang, Xidong (2009). "Calculations of Freezing Point Depression, Boiling Point Elevation, Vapor Pressure and Enthalpies of Vaporization of Electrolyte Solutions by a Modified Three-Characteristic Parameter Correlation Model". Journal of Solution Chemistry. 38 (9): 1097–1117. doi:10.1007/s10953-009-9433-0. ISSN 0095-9782. S2CID 96186176.
^ Ge, Xinlei; Wang, Xidong; Zhang, Mei; Seetharaman, Seshadri (2007). "Correlation and Prediction of Activity and Osmotic Coefficients of Aqueous Electrolytes at 298.15 K by the Modified TCPC Model". Journal of Chemical & Engineering Data. 52 (2): 538–547. doi:10.1021/je060451k. ISSN 0021-9568.
^ Ge, Xinlei; Zhang, Mei; Guo, Min; Wang, Xidong (2008). "Correlation and Prediction of Thermodynamic Properties of Some Complex Aqueous Electrolytes by the Modified Three-Characteristic-Parameter Correlation Model". Journal of Chemical & Engineering Data. 53 (4): 950–958. doi:10.1021/je7006499. ISSN 0021-9568.
^ Ge, Xinlei; Zhang, Mei; Guo, Min; Wang, Xidong (2008). "Correlation and Prediction of Thermodynamic Properties of Nonaqueous Electrolytes by the Modified TCPC Model". Journal of Chemical & Engineering Data. 53 (1): 149–159. doi:10.1021/je700446q. ISSN 0021-9568.
^ Ge, Xinlei; Wang, Xidong (2009). "A Simple Two-Parameter Correlation Model for Aqueous Electrolyte Solutions across a Wide Range of Temperatures†". Journal of Chemical & Engineering Data. 54 (2): 179–186. doi:10.1021/je800483q. ISSN 0021-9568.