By Jason Doub

A solution is defined as a homogeneous mixture or a mixture where the components are uniformly mingled. Normally we thing of a solution is the liquid state, but this is a common misconception; a solution can be in any state. For example, air is a solution of oxygen, nitrogen, and a variety of other gases all in the gas state. Steel is also solution of carbon and iron, but this solution exists in the solid state. However, the most important solutions in chemistry involve water, so this section will focus of aqueous solution.

Before we begin discussing the properties of solutions, we must define ways to describe the concentrations of various mixtures:

Molarity (M): the number of moles per liter (moles/Liter)

Mole fraction (c ): the ratio of the number of moles of a compound to the number of moles of solution (c A = moles A / (Moles A + Moles B))

Molality (m): the number of moles solute1 per kilogram solvent (m = moles solute/ kilograms solvent)

Mass percent or weight percent: the percent by mass of a solute1 in a solution (Mass percent = mass solute / mass solution * 100)

1 A solute is the compound being dissolved in medium. The solvent is the medium in which the compound in being dissolved.

Forming a solution

Dissolving compound in liquids, such as water, is very common; so it would be valuable to understand how and why compound form solutions. To address these questions we must examine what happens when you add a compound to a liquid:

Step 1: the solute particles separate (overcoming the intermolecular attractive forces).

Step 2: the solvent particles separate (overcoming the intermolecular attractive forces).

Step 3: the solute and solvent interact to form the solution.

Each of these steps results in a change of energy. We can sum these energy changes to find the energy change of the solution, called enthalpy (heat) of solution (D Hsoln). Therefore, the D Hsoln is:

D Hsoln =D H1 +D H2 + D H3

The first two steps usually result in a large positive D H, but the third step is usually negative. If the first two enthalpies are greater that the third, the overall D Hsoln will be positive and vice versa. It is important to note, that process with a large positive D Hsoln tend not to occur; however, process with small positive D Hsoln may occur because nature strives for disorder (entropy), and a solution is more disordered than its individual components.





The above concepts can be applied to understand structure can effect solubility; let’s use water mixed with oil for an example. The first two steps from above result in a positive D H, because it requires energy to break water’s hydrogen bonding, and oil’s dispersion forces. The third step does not counter this with a large negative D H, because there is little interaction between oil and water. Thus the D Hsoln is large and positive, so the reaction is not favored. Salts do dissolve in water, because they have a significant interaction with water resulting in a low D Hsoln. Salts are describes as hydrophilic (water loving) where oils are called hydrophobic (water fearing).

Pressure can also effect the solubility of gasses; the relationship between pressure and the concentration of dissolved gas is described in Henry’s law; the amount of gas dissolved is directly proportional to the pressure of the gas above the solution:


Temperature effects the solubility of solutions. One normally thinks that by heating the solution, more solute can be dissolved: this is a misconception. There is come correlation between D Hosoln of a solution and the solubility (see Le Chatelier’s Principle), but the temperature dependence can only truly be seen by experiments.






Gases dissolved in a solution follow a simple rule; the higher the temperature, the less gas dissolved.

Effects of solutions

When a nonvolatile solute is added to a solvent, the vapor pressure of the solvent is lowered. Fancois M. Raoult studied phenomenon, and his results are desribes by Raoult’s Law:

Psoln = C solventPosolvent

One is able to use a simple model to understand the phenomenon; think of the dissolved solute molecules evenly distributed in the solution, even at the surface. Then understand that this caused less molecules of volatile solvent to be at the surface and less molecules evaporating.

This decrease in vapor pressure as a result of the solution can help to explain the following colligative properties, properties that depend only on the number of molecules not he molecules themselves.

When a nonvolatile solute is added to a solvent, the boiling point of the solution is higher than that of the pure solvent. This can be explained by the decreased vapor pressure of the solution; the solution boils when the vapor pressure matches the outer pressure (usually 1 atm), but since the vapor pressure is lowered, a higher temperature is needed to match that outer pressure. This change is given by the equation:

D T = Kbmsolute

Similarly, when a nonvolatile solute is added to a solvent, the freezing point of the solution is lower than that of the pure solvent. This is also explained by the decreased vapor pressure; the solution freezes when the vapor pressure for the liquid and solid phases is the same. This occurs at a lower temperature because of the decreased vapor pressure. This change is given by the equation;

D T = Kfmsolute







The osmotic pressure of a solution is also a colligative property. Osmotic pressure (p ) is the pressure that must be applied to stop osmosis, or the flow of solvent down a concentration gradient. The solvent attempts to move to an area where there is a lesser concentration; the driving force of this movement is an increase in disorder. Experimentally, the osmotic pressure has been show to follow the equation:

p = MRT

where p is the osmotic pressure in atmospheres, M is the molarity of the solute, R is the universal gas law constant, and T is the temperature.

All these colligative properties must be slightly changed for electrolytes, which breakup into multiple pieces. The van’t Hoff factor i must be inserted into the equations. The van’t Hoff factor I is:

i = moles of particle in solution / moles of solute dissolved

Therefore, the new modified freezing point and boiling point equation is:

D T = iKm

The modifies osmotic pressure equation is:

p = iMRT

One should note that ion pairing may result is slightly lowered results than calculated. For example I for NaCl is 2, but the observed i for a .10 m solution is 1.87. This is caused by a few ions pairing together making a fewer number of molecules.


A colloid is similar to a solution, except that a colloid is a suspension of single large molecule or clumps of small molecules. Mud is an example of a colloid; the mud particles stick together and remain suspended in water for an extremely long period of time. The reason these particles remain in solution is that they have layers of oppositely charged articles. When the charges on the outside of these particles are the same, they experience electrostatic repulsion and remain suspended.







Colloids can be destroyed, called coagulation, by heating or adding an electrolyte.


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