The solubility of n-butanol in water is 77.0 g/L. This means that up to 77.0 g of n-butanol will dissolve in one liter of water. When an n-butanol molecule dissolves in water, instead of being surrounded by other n-butanol molecules, it is surrounded by water molecules. These two different configurations represent two different states.
In a mixture of n-butanol and water, n-butanol molecules are constantly breaking away from the pure n-butanol state and entering a “dissolved-in-water” state. And n-butanol molecules in the dissolved-in-water state are constantly returning to the pure n-butanol state. The rate at which an n-butanol molecule changes state depends on the intermolecular attraction it has to water molecules and to other n-butanol molecules. If an n-butanol molecule is strongly attracted to water molecules and only weakly attracted to other n-butanol molecules, then it will dissolve easily. However, if an n-butanol molecule is strongly attracted to other n-butanol molecules and only weakly attracted to water molecules, then it will barely dissolve at all. These processes are very similar to the processes of evaporation and condensation we studied earlier.
Table salt is sodium chloride (NaCl)… sodium and chlorine ions chemically bonded together in crystals as a solid. The chemical bonds holding the sodium and chlorine ions together are ionic bonds. Basically, the positively charged sodium ions are electrostatically attracted to the negatively charged chlorine ions. This attraction is what makes sodium chloride a stable configuration. However, there is nothing specific about the attraction between sodium and chlorine ions. As a positively charged ion, sodium will be just as strongly attracted to any other negatively charged ion. And as a negatively charged ion, chlorine will be just as strongly attracted to any other positively charged ion.
Like all atoms and molecules, the sodium and chlorine ions are in constant motion, and some of those ions will be moving much faster than other ions. A sodium ion on the surface of a sodium chloride crystal has a chance to break away and enter a gas state in the air. But for a sodium ion, breaking away and going off on its own puts it in a very unstable configuration. It can and does happen, but it is extremely unlikely. This is why sodium chloride has a high melting point (801 °C) and why you do not hear about salt crystals evaporating (or sublimating) like water does.
But put that same sodium chloride crystal in liquid water, and the stability of a sodium ion breaking away and going off on its own is suddenly much greater. Because water is a polar molecule with a slightly negative charge on the oxygen atom, water molecules can balance the positive charge on the sodium ion by surrounding it with negative dipoles (δ-). Negative dipoles are much weaker than the negative charge on a chlorine ion, so it takes many water molecules to take the place of a single chlorine ion. And while the configuration of a sodium ion surrounded by polar water molecules is still less stable than a sodium ion in a crystal structure surrounded by chlorine ions, it is close enough to enable some sodium and chlorine ions to dissolve in water. (A negative chlorine ion dissolved in water is surrounded by the positive dipoles on water molecules.)
As more and more sodium and chlorine ions break away and enter a dissolved state, the concentration of sodium chloride in the water increases. And as this concentration increases, the rate of “phase joining” (the rate of sodium and chlorine ions returning to the pure sodium chloride crystal state) also increases. This happens simply because the more sodium and chlorine ions there are in the dissolved state, the more frequently those ions will come into contact with the salt crystal and rejoin it. This dynamic system will be in equilibrium when the opposing rates of dissolution and phase joining are balanced.
For sodium chloride, equilibrium is achieved when the concentration of sodium chloride is 359 g per liter of water. At this concentration, there is one dissolved sodium ion and one dissolved chlorine ion for every nine water molecules… and those dissolved sodium and chlorine ions are rejoining the sodium chloride crystal state just as quickly as sodium and chlorine ions in the sodium chloride crystal state are breaking away and dissolving in water. Because this is the maximum amount of sodium chloride that can dissolve in one liter of water, we say that the solubility of sodium chloride in water is 359 g/L, and that a salt water solution at that concentration is “saturated.”
Besides the relative strength of the intermolecular attraction between the solute and solvent, the solubility of a substance also depends on the temperature and pressure. While 359 g of sodium chloride will dissolve in one liter of water at 20 °C, only 357 g will dissolve at 0 °C and up to 390 g will dissolve at 100 °C. For sodium chloride, raising the temperature (and increasing the speed of molecules and ions) tips the scale towards a faster rate of dissolution. At higher temperatures, more sodium and chlorine ions will break away from the sodium chloride crystal state and dissolve in water. In general (but not always), the solubility of a solid increases as the temperature increases. This effect is even more noticeable for solids such as potassium chloride (KCl) and sugar (sucrose).
In this simulation, we have added 600 g of potassium chloride (KCl) to one liter of water. At 20 °C, only 342 g of KCl will dissolve in the water, and the other 258 g of KCl will settle on the bottom of the beaker as a solid. However, as you increase the temperature of the KCl and water mixture, the amount of KCl that will dissolve in the water increases and the amount of pure KCl sitting on the bottom of the beaker decreases. Decreasing the temperature of the mixture reverses this process. As you decrease the temperature, the amount of KCl that will dissolve in the water decreases and the amount of pure KCl sitting on the bottom of the beaker increases. For solid solutes, we call the process of phase joining “precipitation.” When solute molecules leave the dissolved state and return to the pure solid state, we say that the solute is precipitating out of the solution.
Like solid solutes, the solubility of gas solutes is also affected by temperature and pressure. If you leave a glass of water sitting out on a counter, some of the gases in the air will dissolve in the water. One of the more important dissolved gases in water is oxygen (O2). Many marine organisms breathe dissolved oxygen for cellular respiration.
Unlike solid solutes, the solubility of a gas solute tends to (but not always) decrease as the temperature is increased. At higher temperatures, more oxygen molecules will break away from the dissolved liquid state and re-enter the gas state. While 0.043 g of oxygen will dissolve in one liter of water at 20 °C, up to 0.070 g will dissolve at 0 °C and only 0.027 g will dissolve at 50 °C.
|temperature (°C)||solubility in water (g/L)|
The solubility of a gas solute is directly proportional to the pressure of the gas. Doubling the pressure doubles the amount of oxygen that will dissolve in water. Halving the pressure halves the amount of oxygen that will dissolve in water. This is known as Henry’s law. Basically, increasing the pressure of a gas increases the density of the gas molecules above a liquid surface, which increases the number of gas molecules that come into contact with and can be dissolved in the liquid solvent. If the density of the gas above the liquid surface is doubled, the concentration of gas molecules in the liquid has to double as well to stay in equilibrium.
A room temperature can of soda has an internal pressure of about 3.5 atm (354.6 kPa, or 3.5 times standard atmospheric pressure). At that pressure, the solubility of CO2 in water is 6.23 g/L. Once you open the can, the pressure drops to 1.0 atm, and at that pressure, the solubility of CO2 in water is only 1.78 g/L. The little bubbles you see when you open a can of soda are not dissolved CO2. When CO2 dissolves in water, the CO2 molecules actually enter the liquid state and you cannot see them in the water. The bubbles in soda water are pure CO2 molecules in the gas state that have left the dissolved state because of the drop in pressure. Over time, those bubbles of pure CO2 gas will rise to the surface and escape into the air.