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Lowering the Freezing Point of Water

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Published: Mon, 5 Dec 2016

It is common knowledge that the freezing point of pure water is at 0 degrees centigrade or 32 degrees ferinheight. However, is it possible to keep water in its liquid state below that freezing point? It is indeed possible, and people have been using this principle for centuries! Traveling back to the 1600s we find King Charles I of England dining with his lords and ladies. The final course is the epicurean delight of ice cream. It is doubtful that King Charles I understood the scientific principle of depressing the freezing point of a solution; nevertheless, at that time it was impossible to make ice cream without freezing the crème by depressing the freezing point of water below 0 degrees centigrade (Zinger, 2005). Today, principalities spread salt on icy roads in order to “melt” the ice. In actuality, the salt is merely depressing the freezing point of the water, allowing the roads to remain ice free even while the temperatures are below 0 degrees centigrade.

To comprehend freezing point depression, you must first understand freezing point. Simply put, it is considered the temperature at which a liquid changes into its solid phase. However, it can also be thought of as the temperature at which the liquid and solid phases are at equilibrium with the atmospheric, or vapor, pressure around it. Freezing occurs as water molecules become ordered into a crystalline lattice.

Scientists have long known about the phenomenon that when you add a solute to a solvent, the freezing point lowers, or depresses. Freezing point depression is a colligative property. Colligative properties are the properties of solutions that depend on the number of molecules in a solvent. It does not depend on the properties of the individual molecules in the solution (Prentice-Hall, 1972). As an example, when you create a solution by adding sodium chloride as the solute, to the solvent of water, the freezing temperature of the solution decreases. The increase of the number of solute particles of the solution interferes with the development of the crystalline structure, therefore the freezing process is delayed (Newton, 1999). Freezing point depression can be expressed mathematically as: ΔT = i Kf m. The ΔT equals change in temperature, i equals the number of particles into which the solute dissociates, m equals the moles of the solute per kilogram of solvent, and Kf equals the molal freeing point constant (for water, Kf = 1.853 C/m) (thinkquest, 2010).

As discussed, solutes interfere with the shifting of a liquid to a solid state. The colligative properties relate to the number of solute particles in a solution. The greater the solute particles there are in a solution, the greater the decrease in freezing temperature. If 10 grams of sodium chloride were added to 100 grams of water, the freezing point would drop to -5.9 degrees centigrade. However if 10 grams of sucrose were added to 100 grams of water, the water solution’s freezing point would only drop to -0.56 degrees centigrade. Why the dramatic difference between the two? After all, the same amount of sucrose and sodium chloride was added to the same amount of water. The answer lies in the number of particles in each solute. There are more particles in 10 grams of sodium chloride then there are in 10 grams of sucrose. Sucrose, C12 H22 O11, has a molecular weight of 342.3 grams per mole. Sodium Chloride on the other hand, has a molecular weight of 58.44 grams per mole. Sodium Chloride has almost six times as many particles than sucrose has in the same number of grams. Therefore, the sodium chloride solution has a lower freezing point than the sucrose solution (Chemistry Explained, 2010).

Not only is it possible to quantify the depression of the freezing point of a solution, it is possible to predict how far the freezing point will be decreased. According to the principles of the colligative properties, it doesn’t matter what the physical properties of the solute added to the solution may be. The only determining factor is the number of particles in the solution. Therefore, if you double the amount of sodium chloride in a solution, the depression of the freezing point will be double the original solution.

The original question of, “is it possible to keep water in its liquid state below that freezing point?”, has most assuredly been answered with a resounding yes. Not only can it be lowered, that lowering can be understood, quantified and predicted. In the experiment phase of this project, the scientific method will be used to assess the validity of this research. King Charles I of England would be surprised to know that his epicurean delight of ice cream paved the way for the discoveries of colligative properties and lowering the freezing point of water.

Bibliography

A Brief History of Ice Cream, http://www.zingersicecream.com/history.htm

Colligative Properties, http://www.chemistryexplained.com/Ce-Co/Colligative-Properties.html

Solutions and Colligative Properties: Colligative Properties, http://library.thinkquest.org/C006669/data/Chem/colligative/colligative.html

W.J. Moore, Physical Chemistry Prentice-Hall 1972


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