At the top the water molecules are tightly bound to each other by hydrogen bonds and the Ca2+ ions and Cl– ions are held together by strong ionic bonds. To form a solution these bonds must be broken, as shown in the middle. Finally, the ions form bonds with water molecules as shown at the bottom.
Solutions are homogeneous mixtures of a solute in a solvent. In order for particles of a solute to dissolve they have to break the bonds that hold them together. This is endothermic and consumes energy. Likewise, when particles of the solvent separate to allow solute particles to come in between them they must break bonds holding them together. Of course this, too, is endothermic. If this were the whole story then the formation of solutions would be very unlikely to happen spontaneously. But it is not the whole story. The solute particles form new bonds to the solvent particles and new bond formation is exothermic. This pushes the formation of a solution toward spontaneity. There is also another important factor. Particles have more ways to be arranged when they are in a solution than when they are pure. Since situations with more ways to be arranged are more likely, solutions are likely to form.
Most solutions that we encounter in our daily lives have a nearly perfect balance of energy absorbed and energy given off so that solutions form without any obvious drama. Salt, sugar, maple syrup, honey, instant coffee, and powdered detergents all dissolve without any hint that heat is given off or absorbed. But some chemicals do clearly absorb or give off heat when they dissolve. For example, it is a common laboratory experiment to measure the temperature change in water when potassium nitrate (KNO3) is dissolved in it. When this salt dissolves it absorbs energy and the temperature of the water drops. Another example is anhydrous calcium chloride (CaCl2). This salt is sometimes used in pickling vegetables but in northern climates is also sold in large bags for the purpose of melting ice. This capacity of salt to encourage the melting of ice is the subject of another demonstration. It is an example of the application of a colligative property.
Calcium chloride releases heat when it dissolves, provided any water of hydration is missing from its crystal lattice. The release of heat can be quite impressive under the right circumstances. When the ions of the salt are taken apart they have to break very strong ionic bonds. And when the water molecules are pulled apart to accommodate the ions they must break the weaker but still powerful hydrogen bonds that hold them together. But as the water molecules surround the ions (usually six water molecules per ion) they form medium-strength ion-dipole bonds. The water molecules are said to hydrate the ions. So many of these hydration bonds form that the net result is the release of quite a lot of heat. The heat of solution of calcium chloride is –81.3 kJ/mol.
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The calcium chloride must be anhydrous for the demonstration to work. It is common for a lab supply of this salt to be a dihydrate. This dihydrate often absorbs more moisture from the air to become a mixture of the dihydrate, the tetrahydrate, and the hexahydrate. I found that I could dehydrate about 19 g of the dihydrate in a crucible heated gently over a bunsen burner for about 15 minutes. When it was done the material had fused together but a little water around the edges got it loose. Then some hard work with a mortar and pestle ground it down to produce a yield of about 11.5 g. One gram of this material mixed with about 5 mL of water in a small (10 mL) test tube produced a temperature change of almost 30°C from 22°C to 51°C. This is very hot and students should not touch it for long, though it is safe to touch the outside of the test tube. A couple of test tubes in various locations around the room will help the presentation. Allow students to add the water themselves and to see the thermometer’s rise.
To do the demonstration prepare a few test tubes (one per three or four students) with 1.0 g of anhydrous calcium chloride. Give the test tubes to students along with a thermometer and a squirt bottle of tap or distilled water. Tell students to note the initial temperature, which should be the same as room temperature for a bottle of water that's been sitting on the table. Then tell them to add about 5 mL of water and to stir with the thermometer, noting the final temperature.
Questions
Describe the demonstration set-up and the results of the demonstration.
How strong are the bonds between ions in an ionic compounds such as CaCl2?
How strong are the bonds between water molecules in liquid water?
Both types of bonds have to be broken when a solution of CaCl2 forms, and these steps are endothermic. What bonds are forming when the solution is made? (Hint: Look at the illustration).
Based on the results of the demonstration, which has a larger absolute value:
the sum of the energy required to breaks bonds between both water molecules and ions in solid CaCl2? or
the energy released when bonds form between the ions and the water molecules?
Justify your answer.
In your own words, why is the formation of a solution of CaCl2 exothermic?
In the demonstration 1.0 g of CaCl2 is dissolved in 5.0 mL of water. Given that the water starts at a temperature of 22°C, has a heat capacity of 4.184 J/g·°C, and assuming no heat is lost to the surroundings, what should the final temperature of the water be? Remember from the text above that the heat of solution of CaCl2 is 81.3 kJ/mol.
What was the final temperature you measured? Account for any difference between your calculation and your measurement, thinking carefully about assumptions you made.
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1.0 g of CaCl2 is 9.01 × 10–3 mol so 732 J of heat should be produced by dissolving this much. For 6.0 g of material (assuming 1 g/mL and the same heat capacity as water) the change in temperature is 29.2°C, which makes the final temperature 51°C.
