Chapter 2
Chemistry Comes Alive
41
2
Te concept of strong and weak bases is more easily ex-
plained. Remember that bases are proton acceptors. Tus,
strong bases
are those, like hydroxides, that dissociate
easily in water and quickly tie up H
1
. On the other hand,
sodium bicarbonate (commonly known as baking soda)
ionizes incompletely and reversibly. Because it accepts rela-
tively few protons, its released bicarbonate ion is considered
a
weak base
.
Now let’s examine how one buffer system helps to maintain
pH homeostasis of the blood. Although there are other chemi-
cal blood buffers, the
carbonic acid–bicarbonate system
is a
major one. Carbonic acid (H
2
CO
3
) dissociates reversibly, releas-
ing bicarbonate ions (HCO
3
2
) and protons (H
1
):
Response to rise in pH
H
2
CO
3
HCO
3
2
1
H
1
H
1
donor
Response to drop in pH
H
1
acceptor
proton
(weak acid)
(weak base)
Te chemical equilibrium between carbonic acid (a weak acid)
and bicarbonate ion (a weak base) resists changes in blood
pH by shiFing to the right or leF as H
1
ions are added to or
removed from the blood. As blood pH rises (becomes more
alkaline due to the addition of a strong base), the equilibrium
shiFs to the right, forcing more carbonic acid to dissociate.
Similarly, as blood pH begins to drop (becomes more acidic
due to the addition of a strong acid), the equilibrium shiFs to
the leF as more bicarbonate ions begin to bind with protons.
As you can see, strong bases are replaced by a weak base (bi-
carbonate ion) and protons released by strong acids are tied
up in a weak one (carbonic acid). In either case, the blood pH
changes much less than it would in the absence of the buffer-
ing system. We discuss acid-base balance and buffers in more
detail in Chapter 26.
Check Your Understanding
16.
Water makes up 60–80% of living matter. What property
makes it an excellent solvent?
17.
Salts are electrolytes. What does that mean?
18.
Which ion is responsible for increased acidity?
19.
To minimize the sharp pH shift that occurs when a strong
acid is added to a solution, is it better to add a weak base or
a strong base? Why?
For answers, see Appendix H.
Organic Compounds
Describe and compare the building blocks, general structures,
and biological functions of carbohydrates and lipids.
Explain the role of dehydration synthesis and hydrolysis in
forming and breaking down organic molecules.
Molecules unique to living systems—carbohydrates, lipids
(fats), proteins, and nucleic acids—all contain carbon and hence
are organic compounds. Organic compounds are generally dis-
tinguished by the fact that they contain carbon, and inorganic
Tis type of reaction is called a
neutralization reaction
, be-
cause the joining of H
1
and OH
2
to form water neutralizes the
solution. Although the salt produced is written in molecular
form (NaCl), remember that it actually exists as dissociated so-
dium and chloride ions when dissolved in water.
Buffers
Living cells are extraordinarily sensitive to even slight changes
in the pH of the environment. In high concentrations, acids
and bases are extremely damaging to living tissue. Imagine what
would happen to all those hydrogen bonds in biological mole-
cules with large numbers of free H
1
running around. (Can’t you
just hear those molecules saying “Why share hydrogen when I
can have my own?”)
Homeostasis of acid-base balance is carefully regulated by the
kidneys and lungs and by chemical systems (proteins and other
types of molecules) called
buffers
. Buffers resist abrupt and large
swings in the pH of body fluids by releasing hydrogen ions (act-
ing as acids) when the pH begins to rise and by binding hydrogen
ions (acting as bases) when the pH drops. Because blood comes
into close contact with nearly every body cell, regulating its pH is
particularly critical. Normally, blood pH varies within a very nar-
row range (7.35 to 7.45). If blood pH varies from these limits by
more than a few tenths of a unit, it may be fatal.
±o comprehend how chemical buffer systems operate, you
must thoroughly understand strong and weak acids and bases.
Te first important concept is that the acidity of a solution re-
flects
only
the free hydrogen ions, not those still bound to anions.
Consequently, acids that dissociate completely and irreversibly
in water are called
strong acids
, because they can dramatically
change the pH of a solution. Examples are hydrochloric acid
and sulfuric acid. If we could count out 100 hydrochloric acid
molecules and place them in 1 milliliter (ml) of water, we could
expect to end up with 100 H
1
, 100 Cl
2
, and no undissociated
hydrochloric acid molecules in that solution.
Acids that do not dissociate completely, like carbonic acid
(H
2
CO
3
) and acetic acid (HAc), are
weak acids
. If we were to
place 100 acetic acid molecules in 1 ml of water, the reaction
would be something like this:
100 HAc
S
90 HAc
1
10 H
1
1
10 Ac
2
Because undissociated acids do not affect pH, the acetic acid
solution is much less acidic than the HCl solution. Weak acids
dissociate in a predictable way, and molecules of the intact acid
are in dynamic equilibrium with the dissociated ions. Conse-
quently, the dissociation of acetic acid may also be written as
HAc
÷
H
+
1
Ac
2
Tis viewpoint allows us to see that if H
1
(released by a
strong acid) is added to the acetic acid solution, the equilibrium
will shiF to the leF and some H
1
and Ac
2
will recombine to
form HAc. On the other hand, if a strong base is added and the
pH begins to rise, the equilibrium shiFs to the right and more
HAc molecules dissociate to release H
1
. Tis characteristic of
weak acids allows them to play important roles in the chemical
buffer systems of the body.
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