Percentage
Yield |
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.
- 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.
- 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.
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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.
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Solution
Chemistry |
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.
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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:
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| 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.
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| 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.)
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Solubility |
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).
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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.
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Very
Soluble
Ionic
Compounds |
- 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:
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Very
Insoluble
Ionic
Compounds
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- 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.
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There are at least two reasons that compounds will agree to dissolve
in water.
- 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.
- 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!
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| 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.
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Non-
electrolytes |
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.
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ACIDS
and
BASES |
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.
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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)?
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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.
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Strong
vs
Concentrated |
"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!
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Strong
Acids |
- 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.)
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Strong
Bases |
- 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)
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Common
Properties
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!)
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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.
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Indicators |
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 juice...no 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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