Video Transcript
In this video, we will define and learn to measure the solubility of a substance and how the solubility of a substance is affected by the solvent in which it is dissolved, the temperature, and the pH.
Let’s start by asking, what is solubility? When a solute such as sugar dissolves in a solvent such as water, a solution of dissolved solute is formed. If more solute can be dissolved, we say the solution is unsaturated. If more and more solute is added and dissolved, the solution will become more and more concentrated. A point will be reached where the solution cannot hold more dissolved solute. The solution is so concentrated that it contains the maximum amount of dissolved solute. We say the solution has become saturated.
A saturated solution contains the maximum amount of dissolved solute. Even if more solute is added, it will not dissolve but will separate out from the solution as undissolved solid solute. Now different substances dissolve to different degrees in different solvents, in other words, in different amounts. Some substances can dissolve in large amounts in a certain solvent. For example, a lot more sugar can dissolve in 100 milliliters of water compared to the amount of table salt that can dissolve in the same amount of water, specifically, about 200 grams of sugar versus about 36 grams of sodium chloride salt if the temperature is 20 degrees Celsius.
Sugar and salt are said to have different solubilities. Solubility is the maximum amount of solute in grams that can dissolve in a given amount of solvent, usually 100 milliliters or 100 grams, at a specific temperature. So solubility is for saturated solutions. In this video, we will focus on water as the solvent, although there are many solvents out there. And we will focus mainly on solid solutes, although solutes can be solid, liquid, or gas. The table shows the solubility values of sugar and table salt at different temperatures in terms of grams of solute per 100 grams of water. We can see that the amount or mass of a substance which can dissolve depends on the temperature.
In general, the higher the temperature is, the more solute can dissolve. But this is not always true, and we will discuss this a bit further later on. We can visualize the data from the table on a solubility graph, which has temperature on the 𝑥-axis and solubility on the 𝑦-axis. These graphs are useful for determining the solubility of a solute at a temperature that is not necessarily shown on a table, as well as for making predictions and for comparing different solutes. Now let’s look at an interesting solubility graph, which shows a big difference between solids and gases.
This graph shows a general rule for the difference in behavior between solids and liquids versus gases. Generally, for solids and liquids, the higher the temperature, the higher the solubility. In other words, a high temperature promotes the dissolving or dissolution of a solid or liquid. Generally, gases follow the opposite trend. A decrease in temperature increases the solubility. Or we could say, increasing the temperature decreases the solubility and hinders the dissolution of gas particles in solution.
Why this opposite trend? Regardless of whether the solute is solid, liquid, or gas, the higher the temperature, the more kinetic energy the solute particles will have. In the case of solid solutes, the more kinetic energy they have, the more likely the particles are to break free of the attractive forces holding them to other solute particles. And they can then move off and dissolve in solution. For gas particles, the more kinetic energy they have, the more likely they are to break free of the attractive forces from solvent molecules and escape into the atmosphere. In other words, for solids and liquids, a high temperature aids dissolving and for gases, a high temperature aids the particles leaving the solution.
Remember, these are just general rules. There are always exceptions to the rule. This general rule for gases is part of the explanation why an opened can of soda retains some of its fizziness and dissolved gas when it is cool, whereas a warm opened can goes flatter much more quickly. Now not all substances are soluble in water. Let’s have a look at the common ionic substances, which do and don’t dissolve in water.
Compounds of alkali metals are all soluble in water. These are those which contain lithium, sodium, potassium, rubidium, or cesium ions. Compounds which contain ammonium, nitrate, bicarbonate, chlorate, a halide, sulfate, or acetate are also soluble in water. Halide and sulfate compounds, however, do have some exceptions. For example, a compound containing silver ions and sulfate ions together is not soluble in water. Compounds which are insoluble in water are those which contain carbonate, phosphate, sulfide, or hydroxide ions, with their exceptions shown in blue. For example, a compound containing the calcium ion and the hydroxide ion, calcium hydroxide, does show some solubility in water.
It will be very useful to your chemistry studies to learn this kind of information. We now know which substances do and don’t dissolve in water. For those that do dissolve in water, let’s have a look at those factors which influence the rate at which they dissolve.
If a soluble substance is added to water, it will dissolve all on its own, provided that the amount of solute does not exceed the solubility value at that temperature. However, we can speed up the rate of dissolving or dissolution by decreasing the solute particle size by stirring or by increasing the temperature. Big, chunky pieces of solute will take some time to dissolve. If the big pieces are broken into smaller pieces, they will dissolve a bit faster and very small pieces of solute for, example, powdered solute, dissolve the fastest. What about stirring?
If we took two beakers with the same amount of solvent, the same mass of solute, and the solute is in the same form, for example, powdered sugar, the first we just left to sit by itself and the second we stirred vigorously, we’d notice that in the first beaker, the particles would dissolve. But this would occur very, very slowly. In the second beaker, the particles would be rapidly dispersed by the stirring and would rapidly dissolve. Lastly, what about temperature? The higher the temperature, the more kinetic energy the particles have and the faster they dissolve. We’ve probably all seen this in our daily lives when making a cup of tea or coffee or trying to dissolve some sugar. Using boiling water will cause coffee or sugar to dissolve much faster than, say, ice water.
We now know the factors which affect the rate of dissolution and that temperature affects solubility. Let’s have a look at two other factors which affect solubility.
The choice of solvent will affect the amount of solute that will dissolve in it at a given temperature. Solvents are generally described as being nonpolar or polar. Nonpolar solvents include hexane, oil, and carbon tetrachloride, which is CCl4. And two polar solvents are water and ethanol. Nonpolar solvents dissolve nonpolar solutes and polar solvents dissolve polar or ionic solutes. Two examples of nonpolar solutes are bromine and chlorine, while NaCl or sodium chloride is an ionic solute. And HCl hydrogen chloride, found in hydrochloric acid, is a polar solute. We say “like dissolves like” because a nonpolar solvent dissolves a nonpolar solute and a polar solvent dissolves a polar solute.
If we added a nonpolar solute like bromine and a polar, or in this case ionic, solute such as sodium chloride into a test tube containing hexane and water, which are nonpolar and polar solvents, respectively, and we give it all a good mixing, most of the nonpolar bromine will dissolve in the nonpolar hexane. And most of the polar or ionic sodium chloride will dissolve in the polar water. Again, this is because of the rule like dissolves like. In other words, the solubility of a nonpolar solute will be very high in a nonpolar solvent. And the solubility of a polar or ionic solute will be very high in a polar solvent.
What about pH effect on solubility? Imagine a saturated solution of an acid, let’s call it HA, which contains some undissolved solid HA. The acid and its ions would be in equilibrium with each other with the following equilibrium equation. Now, if we change the pH of the solution by adding acid, in other words, by adding H+ ions, this would increase the total amount of H+ ions in solution. And in turn, this would shift the equilibrium to the left according to Le Chatelier’s principle, decreasing the amount of H+ and A− ions in solution and increasing the amount of solid acid. In this way, adding more acid to the system decreased the solubility of the initial acid or we could say decreased its concentration in solution.
Conversely, if the pH was changed by adding base to the system, in other words, adding OH− ions or hydroxide ions, some of the hydroxide ions would react with some of the H+ ions in solution forming water, decreasing the amount of H+ ions in solution, in other words, removing them from solution. Though there is a decrease in H+ ions in solution, there would not be a decrease in A− ions in solution. Instead, the equilibrium would shift to the right, according to Le Chatelier’s principle, decreasing the amount of solid acid with more of it dissolving into solution. And this would increase the amount of H+ ions and A− ions in solution again. In other words, there’d be an increased solubility in the acid or a higher concentration of ions in solution.
Now we understand the basics of how a solvent and pH affect solubility. Sometimes, it is possible to change the solubility of two solutions by mixing them. Let’s investigate this further. Sometimes, when two aqueous solutions are mixed together, a reaction occurs. Sometimes, the reaction forms an insoluble solid product called a precipitate. A precipitate is an insoluble solid that forms or is deposited out from a solution. So two dissolved solutes reacted to give an undissolved solute. The precipitate may be a compound or an element. An example of this is when aqueous solutions of potassium chloride, KCl, and silver nitrate, AgNO3, react. One of the products is aqueous — KNO3, potassium nitrate — and one of the products is not aqueous but an insoluble solid precipitate — in this case, AgCl or silver chloride.
If we were given this equation, we would be able to identify that silver chloride is a precipitate using the solubility rules that we learned just now. The formation of a precipitate decreases the solubility of some of the ions in solution. In this case, the concentration of silver ions and chloride ions dissolved in solution would both decrease since they join together to form the insoluble precipitate.
Let’s summarize everything that we have learned. We learned that solubility is the maximum amount of a solute in grams that can dissolve in a given amount of solvent, usually 100 milliliters or 100 grams, at a specific temperature. We learned that increasing the temperature increases the solubility of a solid solute and that increasing the temperature decreases the solubility of a gas. But this is just a general rule.
We used a table to look at the general rules for the solubility of common ionic substances in water and learned that we can use this table to predict whether the product of the reaction between two aqueous solutions will be a precipitate or an aqueous product. We saw that solute particle size, stirring, and temperature affect the rate at which a solute dissolves and that the choice of solvent and pH of a solution influence the solubility of a solute.
