The dissolving process involves a consideration of the relative strength of three intermolecular attractive forces. The type of forces between solute-solute molecules and solvent-solvent molecules must be considered. These intermolecular attractions must be broken before new solute-solvent attractive forces can become effective. Perhaps the bond breaking and bond forming processes take place simultaneously.
A solute will dissolve in a solvent if the solute-solvent forces of attraction are great enough to overcome the solute-solute and solvent-solvent forces of attraction. Obviously a solute will not dissolve if the solute-solvent forces of attraction are weaker than individual solute and solvent intermolecular attractions. Generally, if all three of the intermolecular forces of attraction are roughly equal, the substances will be soluble in each other.
As a general solubility rule, like solvents dissolve like solutes. This means that ionic or polar solutes dissolve in polar solvents. Non-polar solutes dissolve in non-polar solvents. Polar and ionic solutes do not dissolve in non-polar solvents and vice versa. Remember that when applying the solubility rule: “Likes Dissolve Likes”, that there are no absolutes and there are exceptions with a small amount of solubility possible. The rule is most useful when making comparisons between a series of compounds.
For example, polar ammonia molecules dissolve in polar water molecules. These molecules mix readily because both types of molecules engage in hydrogen bonding. Since the intermolecular attractions are roughly equal, the molecules can break away from each other and form new solute (NH3), solvent (H2O) hydrogen bonds. The -OH group on alcohol is polar and mixes with the polar water through the formation of hydrogen bonds. A wide variety of solutions are in this category such as sugar in water, alcohol in water, acetic and hydrochloric acids.
When an ionic crystal such as NaCl is placed in water, a dissolving reaction will occur. Initially, the positive and negative ions are only attracted to each other. The water molecules are hydrogen bonded to each other. If the crystal is to dissolve, these bonds must be broken. Negative chloride ions on the surface are attracted by neighboring positive sodium ions and by the partially positive hydrogen atom in the polar water molecule. Similarly, the positive sodium ions are attracted by both chloride ions and the partially negative oxygen atom in the polar water molecule. A “tug-of-war” occurs for the positive and negative ions between the other ions in the crystal and the water molecules. Several water molecules are attracted to each of the ions. Whether the crystal dissolves is determined by which attractive force is stronger. If the internal ionic forces in the crystal are the strongest, the crystal does not dissolve. This is the situation in reactions where precipitates form. If the attractions for the ions by the polar water molecules are the strongest, the crystal will dissolve. This is the situation in sodium chloride.
Once the ions are released from the crystals, the ions are completely surrounded by water molecules. Note that the proper atom in the water molecule must “point” toward the correct ion. The charge principle and the partial charges in the polar molecule determine the correct orientation. Partially negative oxygen atoms in the water molecule interact with the positive sodium ion. Partially positive hydrogen atoms in the water molecule interact with the negative chloride ion.
The solubility of solutes is dependent on temperature. When a solid dissolves in a liquid a change in the physical state of the solid analogous to melting takes place. Heat is required to break the bonds holding the molecules in the solid together. At the same time, heat is given off during the formation of new solute — solvent bonds. If the heat given off in the dissolving process is greater than the heat required to break apart the solid, the net dissolving reaction is exothermic (energy given off). The addition of more heat (increases temperature) inhibits the dissolving reaction since excess heat is already being produced by the reaction. This situation is not very common where an increase in temperature produces a decrease in solubility.
If the heat given off in the dissolving reaction is less than the heat required to break apart the solid, the net dissolving reaction is endothermic (energy required). The addition of more heat facilitates the dissolving reaction by providing energy to break bonds in the solid. This is the most common situation where an increase in temperature produces an increase in solubility for solids. The use of first-aid instant cold packs is an application of this solubility principle. A salt such as ammonium nitrate is dissolved in water after a sharp blow breaks the containers for each. The dissolving reaction is endothermic – requires heat. Therefore the heat is drawn from the surroundings, the pack feels cold.
As for as the variation of solubility for a gas with temperature is concerned the solubility of a gas decreases as the temperature increases. More gas is present in a solution with a lower temperature compared to a solution with a higher temperature. The reason for this gas solubility relationship with temperature is very similar to the reason that vapor pressure increases with temperature. Increased temperature causes an increase in kinetic energy. The higher kinetic energy causes more motion in molecules which break intermolecular bonds and escape from solution.
This gas solubility relationship can be remembered if you think about what happens to a “soda pop” as it stands around for awhile at room temperature. The taste is very “flat” since more of the “tangy” carbon dioxide bubbles have escaped. Boiled water also tastes “flat” because all of the oxygen gas has been removed by heating.
Liquids and solids exhibit practically no change of solubility with changes in pressure. Gases as might be expected increase in solubility with an increase in pressure. Henry’s Law states that: The solubility of a gas in a liquid is directly proportional to the pressure of that gas above the surface of the solution. If the pressure is increased, the gas molecules are “forced” into the solution since this will best relieve the pressure that has been applied. The number of gas molecules is decreased. The number of gas molecules
dissolved in solution has increased as shown in the graphic on the left.
Carbonated beverages provide the best example of these phenomena. All carbonated beverages are bottled under pressure to increase the carbon dioxide dissolved in solution. When the bottle is opened, the pressure above the solution decreases. As a result, the solution effervesces and some of the carbon dioxide bubbles off. Deep sea divers may experience a condition called the “bends” if they do not readjust slowly to the lower pressure at the surface. As a result of breathing compressed air and being subjected to high pressures caused by water depth, the amount of nitrogen dissolved in blood and other tissues increases. If the diver returns to the surface too rapidly, the nitrogen forms bubbles in the blood as it becomes less soluble due to a decrease in pressure. The nitrogen bubbles can cause great pain and possibly death.
To alleviate this problem somewhat, artificial breathing mixtures of oxygen and helium are used. Helium is only one-fifth as soluble in blood as nitrogen. As a result, there is less dissolved gas to form bubbles. Another application of Henry’s Law is in the administration of anesthetic gases. If the partial pressure of the anesthetic gas is increased, the anesthetic solubility increases in the blood.