Chapter 11:  Solutions and Their Properties

At the time of writing (Jan. 2002), we are changing over from WebCT to BlackBoard.  This is also an excellent platform but has some major differences from WebCT.  Most of these are in how the various pages appear and in the organization of the contents.  None of this should be disturbing, but it will take a little getting used to--especially when it comes to finding materials.  I am going to be using a different writing style to take care of this, but you must bear with me as I relearn how to do some of these things.  It is hoped that these pages will remain useful and that you will enjoy them and profit from them.

--MDJ  (2002.01.09)

Solutions are homogeneous mixtures of two more substances.

Every solution has a solvent, which is usually the most dominant species--however there are exceptions such as 95% ethanol in water; this is called an aqueous solution with water as the solvent even though the alcohol is the predominant component.  Most of the terminology here is, thus, arbitrary--but usually this presents no problem.

The other components of a solution are called solutes.  Often there is just one of these (e.g., sucrose in water) but there can be more than one.

Solutions are called aqueous (water the solvent) or nonaqueous (any other solvent other that water).

Short Summary:

Know these terms: homogeneous, solute, solvent, aqueous, nonaqueous.

An alternate classification:

We can also class solutions by the physical states of the solute and solvent (i.e., solid, liquid, or gas).  The following table shows the various types of solutions which are possible (along with some examples):
 
 

The Possible Types of Solutions

Solute	Solvent	Example
Gas	Gas	Air (gases mix in all proportions)
Gas	Liquid	Soda water (CO2 in water)
Gas	Solid	Hydrogen (H2) in Pd or Pt metal
Liquid	Liquid	Gasoline (mixture of hydrocarbons); whiskey
Liquid	Solid	Hg in Ag (an "amalgam")
Solid	Liquid	Sea water (NaCl + other salts in water)
Solid	Solid	Alloys (bronze, brass, sterling silver)

A special note about liquid/liquid mixtures:

Two liquids which can form a solution are said to be miscible.  If they mix in all proportions (e.g., ethanol/water) they are said to be infinitely miscible.

Liquids which do not mix at all are called immiscible.

Suggestion:

 Stop now and be sure that you understand the above ideas!

Substances held together by similar intermolecular forces tend to dissolve in each other.  For instance, polar compounds dissolve in others.

Good example:  NaCl dissolves in water (ions and a polar compound, such as water get along famously).  Similarly, ethanol and methanol dissolve readily in water because they hydrogen bond readily.

Exception:  Some salts such as AgCl are nearly insoluble because their lattice energies are too large to allow the hydration energy of the water molcules about the Ag⁺ and Cl^- ions to overcome the lattice energy.

Hydrocarbons such as cyclohexane and heptane mix easily because they have nearly identical dispersion forces and the molecules are about the same size.

Exception: Motor oil is not readily soluble in gasoline (if it were, you'd have major problems) because the hydrocarbon molecules in motor oil are very much larger than the molecules in the gasoline.

The solution process may be understood in two ways:

 

1)	Intermolecular forces:  If the intermolecular forces in a solution are more favorable than having two separate phases, then the solution forms.
2)	Thermodynamics:  If the free energy is negative for the dissolving process, then dissolution occurs spontaneously.  Remember that DG = DH - TDS.

The following figure, stolen from the text shows the process whereby NaCl dissolves in water.  The lattice energy is not as strong as the hydration energy and, hence, NaCl is soluble in water (if you don't believe, go out and stick your head into Tampa Bay...).

This is a fairly straighforward presentation of how intermolecular (or, if you prefer, "interparticle") forces act to cause the dissolving of a salt.  Similar ideas hold for other types of solutions.
 

The thermodynamic point of view is possibly best presented by thinking of DH representing the various forces involved and DS being something which always favors solution formation.  In general, if a mixture forms, S must always increase since a mixture is, by definition, more random than a pure substance.  (A possible exception here is a gas dissolving into a solid--we'll leave that as something for you to ponder over and argue about!)

The following figure pretty much takes care of making this a little clearer:

The following table shows some typical values of enthalpies and entropies of solution (at 25^oC). It is left as an exercise to get the DG values here. If solution formation is spontaneous, then DG < 0 for a given substance.  Are these all soluble in water?  (You can look up solubilities in handbooks; these are usually expressed as "g(solute)/100 mL of solution"--or something similar such as "g(solute)/100 g(solution)."  Anyway, here is a typical table of solution enthalpies and entropies:

It is interesting to note that solution can be either exothermic (DH < 0) or endothermic (DH > 0) and still be spontaneous!  This sort of thing is often exploited with "heat packs" and "cold packs."

Thinking in terms of intermolecular forces from the thermodynamic point of view, we can sequester interactions into

• solute-solvent,
• solute-solute, and
• solvent-solvent

interactions.  How these interact is easily understood from the following figure which shows cases for exothermic and endothermic dissolving, respectively.

A similar sort of analysis can be made with DG values (and that is left to you to ponder).  In any event, this should be clear--if not, don your dunce cap!  We need not consider this stuff too much more at this point--the basic concepts are there and the detailed understanding will come when we start solving problems.

Molality:

This is simply "moles solute per kg of solvent." For dilute solutions--especially aqueous--this is a very useful expression of concentration. It is independent of temperature (as are all concentration units, except for molarity). Units are "moles solute/kg solvent."

Mole fraction:

This is the fraction of moles of solute divided by the total moles present for all substances in a solution or mixture. This unit is most useful when discussing things such as mixing in gases and in writing laws for vapor pressures of liquid mixtures. This way of expressing concentration has no units. Mole fraction is independent of temperature.

Weight Percent (or Mass Percent):

Sometimes molecular masses (weights) are not known for one or more components of a solution. However, masses are known if the proper weighings are carried out. In such cases, we simply express the concentration as a "mass-%" of the total mass of the mixture. There are no units here, per se, but you should realize that percentages are 100 times the fractional value by mass. This quantity is independent of temperature.

Molarity:

This is probably the most important expression of concentration in chemistry and is simply the "moles solute/unit volume of solution." Since volume depends on temperature (in nearly any case one can imagine), molarity does also. When expressing molarities, it is necessary to give at least the temperature. Better yet, one should report the temperature and the solution density. Molarity is useful because it contains an expression of volume. Volume, in turn, is a useful way of recording inter-particle distances; remember that volume is proportional to the cube of the distance between particles!

Molarity vs. Molality:

In dilute aqueous solutions molarity and molality are nearly the same in value. This is a useful shortcut in much work involving aqueous solutions. However, it should be noted that this is not true of most solvents (unless they have a density of 1 g/mL). And, it should be noted that the concept of dilute is very important here. A good operational definition of a dilute solution is, for most purposes, 0.1 M (i. e., 0.1 moles/Liter) or less.