Solutions are No Problem

Chm 1311 Lecture for 21 June 1999

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Although we've assumed all chemical reactions proceed to completion with only the chemical species we've identified, that is rarely the case.

Instead two problems conspire to make actual product yields less than the theoretical yields we've been calculating the last couple of days.
  1. Reactions other than the one of our primary interest often take place to compete for our reactants or consume our products. These undesired competitors are often called side reactions. An example would be incomplete combustion in the elemental analysis of an earlier chapter; if carbon was burned only as far as CO, the CO2 found would be less than the theoretical amount.

  2. Products may take part in reverse reaction and reproduce the reactants. Since the newly-reformed reactants would then continue to react, the outcome would be a stalemate. But it's usually a stalemate that leaves one side more heavily represented than the other, and we hope that side is the product one. Nevertheless, when the dust settles (chemists would say "when equilibrium is reached"), only some fraction of the reactants would have been converted to products, a situation which would likewise reduce the expected (theoretical) yield. Reaction isn't complete.
Indeed, both of these spoilers usually happen at the same time! And the percentage yield is found by comparing the actual yield to the theoretical one:

% Yield = [(Actual Yield) / (Theoretical Yield)] × 100%

Although one would wish for % Yields near 100%, the products of modern chemistry are often arrived at as the result of sequential reactions, each one of which is incomplete. The overall fractional yield will be all the individual fractional yields multiplied together. Those can get small fast! Makes for (sometimes unavoidable) waste. With luck, a chemical manufacturer can reuse unreacted or inopportunely reacted materials or these may be sold to other companies for reuse. In an Era of awareness of the fragility of the environment, it's no wonder that chemical manufacturers strive to minimize their effluent wastes by maximizing their yields; that way they save twice.


Chemicals are by far easier to handle in solutions than in any other form. Solutions can be diluted to deliver inconceivably small amounts of their cargo, if that's required. Liquids take the shape of whatever container they're placed in, making their measurement easy against standardized volume markings. Solutions flow easily and can thus form a stream in industrial processes. Solutions often intermix completely, thus ensuring intimate contact of their contents.
In contrast, gases and solids are a pain to work with. Gases keep trying to escape from whatever contains them, sometimes explosively. Their reaction vessels must be pressurized and are more expensive than solution vessels. Solids are often recalcitrant about mixing well.

All in all, solutions are a real blessing.

That blessing comes in two parts:
Solute: The solute is the item of interest that we're attempting to put into solution; it's the stuff we're trying to dissolve. It need not be a solid, as you might guess. The solute could be a gas which will dissolve in the fluid phase, like the CO2 in soda. Or the solute could be another fluid that we're merely trying to dilute.
Solvent: The solvent is the fluid phase in which we intend to dissolve the solute. In almost all cases in this course, the solvent will be water. Our solutions of interest will be aqueous ones. But as soon as you hit Organic Chemistry, you'll become interested in nonaqueous solvents like benzene or carbon tetrachloride. And in Inorganic Chemistry, the solvents really get wild! (Like thionyl chloride, SOCl2, for instance.)


Some things are infinitely soluble (like ethanol in water), but most make some pretense, at any rate, of reluctance to be dissolved (like BaSO4 which saturates at about 0.2 moles per liter of cold water ... but 3 moles per liter of hot water). Other compounds completely refuse (like PbO2). Whatever the maximum solubility of a particular solute in a particular solvent, when it's reached, the solution is said to be saturated. Any additional solute settles out (if it's a solid or liquid) or bubbles out (if it's a gas).
There isn't any obvious dividing line between soluble and insoluble compounds since solubilities vary from lots to none pretty much continuously. Most books will count solutes which don't saturate until you've crammed at least one mole of 'em in a liter of solvent as very soluble. If you can't get more than 0.01 moles of solute to dissolve in a liter, that's very insoluble. That factor of 100 in between is the realm of the slightly soluble solutes.

It's handy to know what's likely to be very soluble since then you know what compounds are going to be most useful in stuffing various reactants into solution.
  • All compounds of Group I and the ammonium (NH4+) salts.
  • All acetates (C2H3O2-), nitrates (NO3-), chlorates (ClO3-), and perchlorates (ClO4-)
  • Almost all halides except Ag+, Hg22+, and Pb2+.
  • Almost all sulfates (SO42-) except Sr2+, Ba2+, and Pb2+.
For insoluble compounds, we're going to have to exclude all the compounds which qualified above! But with that in mind, the stubborn solutes are:
  • carbonates (CO32-), phosphates (PO43-), oxalates (C2O42-), and chromates (CrO42-)
  • oxides (O2-) and hydroxides (OH-) except Group II elements which are sparingly soluble as hydroxides
  • sulfides (S2-) except Group II elements.
There are at least two reasons that compounds will agree to dissolve in water.
  1. The compound is itself water-like in that it sports -OH bonds or -O- bonds which water adopts as close enough for Government work. That is, the forces which hold water together as a liquid at room temperature (when the heavier molecule below it in the Periodic Table, H2S, is still a gas) arise from (the misnomered) "hydrogen bonds" whereby electron rich O atoms in one water attract the electron poor H atoms in a neighbor. The same effect (hydrogen bonding) works for alcohols or carbohydrates (-OH groups) or ethers (-O- groups) for example.

  2. The compound can ionize in solution yielding cations that love to be surrounded by the O end of waters and anions that get cradled by the H end of water molecules. Water's incredibly high dielectric constant ensures that the attractions between the water-caged ions is over 70× weaker than it would be in a vacuum. So the ions have little motivation to recombine as ionic solids until you boil almost all the water off!
Electrolytes Solutes of the 2nd kind have the free ions necessary to conduct electricity; so their solutions are "electrolytes." Of course a solute could dissolve in water due to both causes above. Then some of it may be dissolved in bound form (I'd say "molecular," but Brady's persnickety about reserving the word molecular for non-ionic compounds...pooh) while some of it is found in dissociated form as solvated ions. Such electrolytes would be weaker than those which ionized completely.

Of course you also could imagine compounds which made stronger than normal electrolytes by falling apart into many ions as would K3PO4, for example, since it produces 4 ions compared to, say, NaCl which makes only 2 ions.
Of course if a species remained molecular in solution and didn't ionize at all, it wouldn't conduct electricity (since pure water is an electrical insulator). By the way, that doesn't mean you should drop electric toasters in your bath tub! Your soaps produce copious Na+ ions as do the salts of your perspiration; so you'll still fry with electrical devices within reach as you shower or bathe.


An excellent example of a weak electrolyte would be a weak acid which ionized reluctantly to yield far less than 100% of its acidic protons as H3O+ ions. Likewise a strong acid would ionize completely and conduct electricity well. But it's not the electrolytic properties of acids which make them so interesting and useful. It is instead their often violently corrosive properties that make them so handy.
It's easy to imagine why acids are de rigeur in Chemistry. We know that atoms hide their protons deep in inaccessible nuclei; so only their electrons are exposed to other atoms. So it is the electrons which effect the chemistry, and the more effective you are at pushing electrons around, the more chemistry you can do! What would be more appropriate as a Big Stick to whack electrons than the steep electrical gradient of a proton on the loose? Even one which has suborned a water molecule as H3O+(aq)?
Well, instead of pulling on them you might push. So the OH-(aq) produced by bases is also likely to influence chemistry. Both acids and bases are important reagents. And the stronger they are (that is, the greater the extent of their ionization to H3O+ and OH-, respectively), the more important they're likely to be.

It's vital that you understand the distinction between the chemist's use of the words strong and concentrated. Common usage treats them as the same; Chemistry does not.
"Strong" is reserved for those species (acids or electrolytes) that ionize completely in solution. "Concentrated" is instead an indication that the solute, ionized or not, is present in large amounts in solution. So, for example, a strong base could easily be prepared in low concentration, but all of it in the solution would be in the form of ions.

Fortunately for chemistry students, almost all acids and bases are weak. So if you're going to memorize strengths of acids and bases, for goodness sake, just memorize the few strong ones!
  • HX, the hydrogen halide acids except HF.
  • Nitric acid, HNO3
  • Chloric and perchloric acids, HClO3 and HClO4 (well, truth to tell, all the higher halogen equivalent acids are strong too...HBrO3, HIO4, etc., etc.)
  • Sulfuric acid, H2SO4, but only the first proton to leave! (HSO4- is a weak acid.)
  • Group I hydroxides (LiOH through CsOH)
  • Group II hydroxides; strong despite the low solubility of Ca(OH)2 (again drawing a distinction between strong and concentrated)

of Acids
and Bases

Common acid properties stem, of course, from the free H3O+(aq) just as common base properties stem from the prevalence of OH-(aq). Acids are sour (lemon juice, say), a fact well-known to 18-19th century German chemists (the best in the world then) who regularly described the "schmeken" (taste) of chemicals. Perhaps that permits you to understand why I said in lab that chemists have actuarily short lives! No one should taste any chemical not known to be completely harmless! (Another reason we don't permit food or drink in the labs.)

But those (short-lived) German chemists understood a common theme among the strong acids: except for the hydrohalic acids, a ponderance of oxygens seemed to promote acidity! HNO3 or HClO4, for example. Indeed their very name for the element oxygen is "sauerstoff", the stuff which makes (sour) acids. (Many other German element names are telling like hydrogen's "wasserstoff", the stuff that makes water, or nitrogen's "stickstoff", the stuff that asphixiates!)

Bases feel slimy because they're dissolving your skin oils. They're bitter [Phillip's Milk of Magnesia, Mg(OH)2] and sometimes metallic tasting since many are metal hydroxides after all.


But the property that permits the easiest identification of an acid or bases is a relatively high concentration of H3O+(aq) or OH-(aq), respectively, Fortunately, these ions can facilitate the attachment or detachment of protons from many molecules, some of which change colors as they gain or lose protons! Such molecules are called indicators since they indicate the presence of an acid or base. One universal indicator is cabbage lie. It turns sort of reddish in the presence of acids and bluish in the presence of bases.

A similar color change occurs with common litmus paper, and this property of being an indictor has elevated the word "litmus" to Washington vernacular. Who hasn't heard of some contentious issue, like abortion, being used as a "litmus test" for judicial candidates?

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Last modified 18 June 1999. Chris Parr