Lecture #1
Introduction to Biochemistry,
Chemical Bonding, Water,
pH and Buffers
Slide 1. Introduction. Welcome to BioC 6011, Biochemistry for Dental
Students. The focus of the first half of this course is to review the basic
concepts of biochemistry, so that you can refresh your knowledge about this
subject and be prepared for your basic science board exams.
(In the process of reviewing biochemistry, we hope that you will become
wiser and more humane individuals, but that is largely up to you.)
In this first lecture set we would like to concentrate on the fundamental
interactions which occur within biological molecules and between these
biological molecules and water. Since all living cells are based on water as a
solvent, it is vitally important that we understand how interactions with
water determine the structure and function of biological molecules.
Slide 2. Molecular components of an E. coli cell. Here we see the
molecular components of an E. coli cell. The major cellular constituent is
water which comprises about 70% of the total cell weight. That high value
again emphasizes the critical role that water plays in living organisms.
Proteins constitute about 15% of the total weight (half of the dry weight of
the cell). The remaining 15% consists of the other biopolymers (the lipids,
carbohydrates, and nucleic acids) as well as low molecular weight organic
metabolites and salts. All of these molecules are present in a water based
system and their structures and metabolic functions are determined by
interaction with water.
Slide 3. Ionic attractions in Sodium Chloride. Ions are charged particles
that are formed from atoms by the gain or loss of electrons. Neutral atoms
are converted in cations by the loss of one or more electrons and into anions
by the gain of one or more electrons. Ionic compounds are composed of
cations and anions held together by ionic bonds. These bonds result from
the attraction of opposite charges. Ionic compounds are generally soluble in
water and they conduct electricity when dissolved in water. In addition they
usually have relatively high melting and boiling points.
In the example used here sodium has a single electron in its outer valence
shell and can be converted to a positively charged cation by the loss of an
electron. Chlorine can fill its valence shell by gaining one electron to form a
negatively charged anion. Sodium chloride is highly soluble in water.
Slide 4. Covalent bonding of hydrogen atoms. Covalent bonds are
formed when two atoms come together and electrical interactions occur.
Some interactions are repulsive—the two positively charged nuclei repel
each other and the negatively charged electrons from the two atoms repel
each other. Other interactions are attractive—each nucleus attracts electrons
from both atoms and each electron attracts both nuclei. Because the
attractive forces are stronger than the repulsive forces a covalent bond is
formed. A covalent single bond (such as H-H) occurs when two atoms
share a pair of electrons.
Slide 5. Covalent bonds in water. Bonds between atoms are polar covalent
if the electrons are not shared equally between the two bonded atoms. In a
polar covalent bond the more electronegative atom carries a partial negative
charge and the more electropositive atom has a partial positive charge. In a
water molecule the oxygen atom is more electronegative and attracts
electrons so it carries a partially negative charge. The two hydrogen atoms
are less electronegative and lose some of their electronic atmosphere so they
carry a partial positive charge.
Slide 6. The occurrence of hydrogen bonding. Hydrogen bonding occurs
when a hydrogen atom, which is covalently bonded to an electronegative
atom such as oxygen, nitrogen, or sulfur, becomes electrostatically attracted
to another electronegative atom. The partial positive charge on the hydrogen
atom attracts the partial negative charge on the electronegative atom of the
oxygen, nitrogen, or sulfur. Hydrogen bonds are relatively weak and often
transient, but the collective effect of multiple hydrogen bonds can result in a
strong attractive force. This is particularly true in macromolecules such as
DNA.
Slide 7. Hydrogen bonding in water. We have seen that in a water
molecule the oxygen is partially negative and the hydrogen atoms are
partially positive. Because of the high polarity the water molecules form an
extensive series of hydrogen bonds with other water molecules. Each water
molecule can form a maximum of four hydrogen bonds.
Slide 8. Hydrogen bonding in ice. In ice each water molecule is bonded to
four other molecules—one bond from each hydrogen atom and two bonds
from each oxygen atom. Much of this bonding persists when ice melts. The
resulting liquid water is fluid but still retains a high percentage of the
hydrogen bonds. This ability to form extensive hydrogen bonds is
responsible for the high thermal constants of water and for the wide spacing
between the freezing point and boiling point of water.
Slide 9. Interaction of water with a hydrophopic molecule. Most
biological molecules are somewhat amphipathic, that is, they have both
polar and non-polar regions. When such compounds are suspended in water
the polar areas (show here in blue) tend to orient to the outside where they
interact with water. In contrast the non-polar, hydrophobic regions (show in
black) generally cluster in the interior of the molecule where they do not
have to interact with the water. These effects occur because the entire
system is at its lowest energy state under those conditions.
Slide 10. Solvation of ions by water. In contrast to non-polar molecules,
small ionic compounds tend to be very stable in water, and they generally
will dissolve readily to form highly concentrated solutions. In the example
show here the negative anions are surrounded by the partially positive
hydrogen atoms of water. The positive cations are in contact with the
partially positive oxygen atoms.
Slide 11. Properties of water. Water exhibits a number of unusual
characteristics that reflect its polarity, its high ability to hydrogen bond and
its expanded crystalline structure.
-It exhibits a very high surface tension which among other things
contributes to its ability to hydrate plants via capillary action.
-It high specific heat helps maintain a constant temperature in higher
animals.
-The high heat of vaporization also contributes to the maintenance of
constant temperature in animal species.
-The characteristic of expansion upon freezing reflects the fact that the
crystalline structure is somewhat expanded in comparison the liquid state.
That is the molecules occupy more space in ice than they do in liquid water.
This means that ice floats, a property that promotes the survival of aquatic
organisms under freezing conditions.
-Water is versatile as a solvent--ionic, polar and amphipathic molecules are
able to form stable solutions in water.
Slide 12. Ionization of water. There is a slight tendency of water to
dissociate into hydrogen and hydroxyl ions. In pure water at 23o C the
concentrations of the hydrogen and hydroxyl ions are both 10-7 M whereas
the concentration of the undissociated water molecules is 55 M.
Slide 13. The ion product of water. A formal dissociation constant can be
calculated for water. That is Keq = (H+)(OH-)/H2O = (H+)(OH-)/55M.
This equation can be simplified by incorporating the almost constant value
for water (55M) with the Keq to create a new constant, Kw. The Kw =
(H+)(OH-) = 10-14 . What this tells us is that the product of (H+)(OH-) is a
constant. That means that if the concentration of (H+) rises the concentration
of (OH-) will fall to an equal extent. The product of the two ionic species
will always equal 10-14 . In practical terms that means that it is impossible to
raise the concentration of both ionic species at the same time.
This inverse relationship (between the hydrogen ion concentration and the
hydroxyl ion concentration) has a consequence for biochemical reactions,
which often are catalyzed by both of these ions. In the absence of an
enzyme, the rate of one component of a chemical reaction often can be
increased by increasing the hydrogen ion concentration. However, if the
overall reaction also requires hydroxyl ion, that latter component of the
reaction will be slowed down. The genius of enzyme catalysis is that an
enzyme can use its functional groups to simultaneously provide (or replace)
both the hydrogen ions and the hydroxyl ions. We will discuss this in more
detail when we consider enzyme catalysis.
Slide 14. Definition of pH. The pH is defined as the negative log of the
hydrogen ion concentration. The examples shown here illustrate the utility
of the pH concept. It provides a very compact way of designating the
hydrogen ion concentration. Scientist, being intrinsically lazy, cling to the
pH concept like limpets, because it helps to prevent writers cramp.
Slide 15. Two caveats concerning the use of pH. There are a couple of
things that you should keep in mind when dealing with pH.
The first is that for each pH change of one unit (for example from pH 3 to
pH 4) the hydrogen ion concentration is changing ten-fold. So, if a pH
change of one unit equals a ten-fold change in hydrogen ion concentration, a
change of two units gives a hundred-fold change in hydrogen ion
concentration, and a pH change of three units gives a thousand-fold change.
A pH change across the entire commonly used range (pH 1-14) equals a
hydrogen ion change of 1014.
The second thing to remember is that as the numbers get smaller the
hydrogen ion concentration is actually increasing. That is a consequence
of taking the negative value of the logarithm. Thus pH 3 (10-3) has a greater
hydrogen ion concentration than pH 4 (10-4).
Slide 16. Strong and weak acids. Now I would like to compare strong and
weak acids. Hydrochloric acid (HCl) is a strong acid that exhibits
essentially complete dissociation in water. In contrast, weak acids are only
partially dissociated in water. Acetic acid serves as a good example of a
weak acid. You may also think of acetic acid with its partially dissociated
carboxyl group as a prototype for any of the carboxyl groups that occur in
biological compounds. Notice that the arrow for the dissociation of HCl
goes only to the right signifying a high degree of dissociation. In contrast
there are reciprocal arrows for the dissociation of acetic acid, and the arrow
pointed to the left is bolder and longer. In this contest size matters, and it
signifies that at equilibrium most of the acetic acid would remain in the
undissociated form.
Right here it might be appropriate to give you a word to the wise. In this
context the word strong refers to an acid that is highly dissociated. The
term has nothing to do with the concentration of the parent acid in solution.
A strong acid can be in a very dilute solution. Conversely, a weak acid can
be highly concentrated.
When we have a weak acid we can write a dissociation equation. (This is
never done for a strong acid.) The KD or dissociation constant for the
reaction equals the molar concentration of the products divided by the molar
concentrations of the reactants. In this case it is the acetate ion
concentration times the hydronium ion concentration divided by the
concentration of the undissociated acetic acid. That value is 1.8 x 10-5 M.
Slide 17. Dissociation constants document the relative strength of weak
acids. Some weak acids are highly dissociated, whereas other are mostly
undissociated. The dissociation constant is a measure of the tendency of a
weak acid to dissociate. A strong weak acid with a large dissociation
constant (the example here is Ka = 10-2) would dissociate to a significant
extent. In contrast a weak weak acid (eg Ka = 10-6) would be largely
undissociated. In any case you can use the dissociation constant to calculate
the pH of a weak acid solution. (nb. Strong acids do not have dissociation
constants—at our level of sophistication, we assume that strong acids are
completely dissociated in water, so the acid concentration added to solution
is taken to equal the hydrogen ion concentration.)
Slide 18. Bronsted acid/base theory. The Bronsted theory of acid-base
behavior will prove to be useful as we consider the behavior of biomolecules
in solution. On the left we have HA, which signifies a generic weak acid.
The acid (HA) can undergo dissociation to release a proton (actually in
water a hydronium ion is formed) plus the conjugate base (A-). The
conjugate base is the part of the molecule that is left over after the acid
donates a proton. You notice that the arrows point in both directions
designating an equilibrium situation in which the reaction is proceeding in
both directions simultaneously. The Bronsted acid (HA) is defined as the
molecule that can serve as a proton donor. When that proton is donated and
the reaction proceeds from left to right the form that is left over (A-) is
referred to as the conjugate base. The conjugate base can then serve as a
proton acceptor because as the reaction goes from the right to the left a
proton is added to form the weak acid.
Slide 19. Conjugate Pairs. Here are some of the conjugate pairs of acids
and bases that we will be focusing on in this course. By convention the acid
form is shown on the left and the conjugate base is on the right. In each
case the removal of a hydrogen ion (proton) from the acid yields the base,
and the addition of a hydrogen ion to the base converts it back into the acid
form.
There are three dissociations of phosphoric acid which occur sequentially.
Each of the dissociations has its own distinct dissociation constant. It is the
middle of the three dissociations which is of physiological relevance in
biological systems.
The dissociation of amino groups and carboxyl groups are very common in
biological systems—most commonly they are found in protein molecules.
The existence of protein amino groups in the positively charged acid form
and carboxyl groups in the negatively charged base form is what leads to the
“zwitterionic” character (carrying both positive and negative charges
simultaneously) of proteins.
Slide 20. Amphoteric compounds. There are certain molecules which can
act as both Bronsted acids and bases. The various dissociated forms of
phosphoric acid are a prime example. In this series of dissociations the
forms H2PO4-1 and HPO4-2 can act as both acids and bases—these forms are
which can donate and accept protons are designated as being amphoteric.
The slide shows a second compound which is of prime importance in
biology. That is carbonic acid which can dissociate to bicarbonate. In a
second dissociation bicarbonate can form the dianion, carbonate. Carbonic
acid is formed in living cells by the reaction of carbon dioxide with water.
As the name suggests, carbonic acid is a weak acid which is capable of
acidifying a solution (or an organism). One of the primary tasks of
respiration in humans is to get rid of excess carbon dioxide, so that we do
not acidify ourselves to death. On a more benign note, the high acidity of
carbonated beverages is due to the reaction of carbon dioxide with water.
Slide 21. Henderson-Hasselbalch equation. A consideration of the
Bronsted theory leads us naturally into a discussion of the HendersonHasselbalch equation. The Henderson-Hasselbalch equation can be
derived quite easily from the dissociation equation which was presented on a
previous slide. However, in the form given it has proven to be more useful
in a biological context. It indicates that the pH of a solution is equal to the
pKa plus the log to the base 10 of the concentration of conjugate base over
the concentration of the Bronsted acid. In this equation the conjugate base is
always in the numerator and the conjugate acid in the denominator. The
subscript “a” in the pKa indicates that we are dealing with a Bronsted acid.
In case you were wondering where this equation comes from, I am going to
relieve your mind. If you were not wondering, that too is OK. In that case I
will relieve my own mind of the nagging need to raise your consciousness.
Slide 22. Derivation of the H-H equation, part 1. The H-H equation is
derived directly from the dissociation equation for a weak acid, using a few
simple steps. (Simple, if you are mathematically inclined and inscrutable if
you are not.)
 On line one we start with the standard dissociation of a generic weak acid.
 Line two is the standard form for the dissociation equation of that acid.
 For line 2-a, we take the hydrogen ion concentration away from the rest
of the right hand components. The little dot between the parts indicates
that the hydrogen ion concentration is multiplied times those other
components.
 Now on line 3 we take the logarithm of all the components. Because we
separated out the hydrogen ion concentration that term gets its own log
expression. The resulting equation tells us that the log of the dissociation
constant is equal to the log of the hydrogen ion concentration plus the log
of the concentration of the conjugate base divided by the conjugate acid.
(These are all common and allowable mathematical manipulations.)
Slide 23. Derivation of the H-H equation, part 2. This slide continues the
derivation of the H-H equation.
 The information on step 3 is repeated on this slide.
 We then transpose two terms from one side of the equation to the other.
After doing that, the log of the hydrogen ion concentration is on the left
side of the equation, and the log of the dissociation constant is on the
right side. Please recall from algebra that when we do this the signs of
these terms become negative.
 The negative log of the hydrogen concentration is defined as the pH, and
we now define the negative log of the dissociation constant as the pK.
When these two terms are substituted into the equation the H-H equation
magically emerges.
Slide 24. Definition of a buffer. The H-H equation allows us to exam a
number of pH related processes in biochemistry. One of these is the
phenomenon of buffering. The more general definition of a buffer is
something that resists changes. For example the shock absorbers on your car
resist the effects of pot holes in the road— this keeps your car from
bouncing around like crazy. As we use the term in biochemistry, a buffer
consists of a mixture of a weak acid and its conjugate base. In a
biological system the mixture of a weak acid and its conjugate base can
resist changes in pH when a strong acid or strong base is added to the
system. To state it in a slightly different way, in the presence of a buffer the
change in hydrogen ion concentration is minimized.
Slide 25. Acetic acid functions as a buffer. Next we look at a picture of
what happens when a source of hydroxide ion is added to a buffer. What
you see is a titration curve for acetic acid. Understanding what is happening
during such a titration is critical to comprehending two things in
biochemistry. One concept is how buffering really works. The other is how
the charge on a biological molecule is altered as the pH changes. In this
figure the ordinate shows the pH value. Remember that this is a logarithmic
function, so for each unit change the hydrogen ion concentration is changing
by a factor of ten. Along the abscissa we have plotted the number of moles
of hydroxide ion added per mole of acetic acid. Acetic acid has a pKa of
4.76. The undissociated acid has the formula CH3COOH and the conjugate
base is CH3COO-. The curve spans the whole range of a titration from the
point on the left where no hydroxide ion was added to the far right where
one mole of hydroxide ion was added for each mole of acetic acid. At the
right end of the curve the conversion of acetic acid to acetate is essentially
complete. What is plotted in between these two extremes is the pH in the
solution as we go from no hydroxide ion added to one equivalent added.
The entire titration curve (that is, the relationship between the amount of
hydroxide ion added and the pH) is determined through the use of the H-H
equation. One relatively simple point to plot is at the center of the titration,
where the acid and conjugate base concentrations are equal. If A- and HA
are equal then the term ( A-)/(HA) is equal to one. The log of one is zero.
Thus at the midpoint in the titration, the term log ( A-)/(HA) falls out of the
equation, and the pH is equal to the pKa.
This relationship holds true for any weak acid—at the center point of the
titration where (HA) and (A-) are equal, the pH always equals the pKa of that
particular weak acid. The rest of the blue line for the titration can be plotted
using other ratios of (HA) and (A-). Regardless of the weak acid being
titrated the curve will always have this same sigmoidal shape, and the center
of the curve will always occur at the point where the pH is equal to the pK of
the weak acid being titrated. In order to understand how the titration curve is
derived it would be a good learning experience for you to plot a titration
curve yourself using the H-H equation to determine the pH at various points
in the titration.
Keep in mind, that as hydroxide ion is added, the sum of (HA) plus (A-)
is constant.. As (HA) goes down the (A-) rises so that the total of (HA) and
(A-) is a constant. The lower part of the figure shows this diagrammatically.
Now why does this mixture of (HA) and (A-) serve as a buffer? Please focus
your attention on the area of the titration within the stippled trapezoid. This
region extends from one pH unit below to one pH unit above the pKa. It
encompasses the area from 10% titration to 90% titration of the weak acid.
You will see that the curve is relatively flat in that region. That is, for a lot
of hydroxide ion added there is very little change in pH. This region is
generally considered to be the optimal buffering zone for any weak acid.
The region will always occur from one pH unit above to one pH unit below
the pKa of the weak acid being titrated. That is, all weak acids will show
this same pattern of buffering. That pattern will move up or down
depending on the pKa of the particular weak acid, but it will always be
centered between one pH unit below to one pH unit above the pKa.
Slide 26. Titration curves for three weak acids. The relationship between
the titration curves for various weak acids is illustrated on this slide. The
lower curve shows the titration for acetic acid which we have just looked at.
Above that is the curve for imidazole (a functional group in the amino acid,
histidine). Notice the curve has the same shape, but it is now centered
around pH 6.99, equal to the pKa for imidazole. The top curve is for
ammonium ion which has a pKa of 9.25. Again you see a curve with the
same shape, now centered around pH 9.25.
Slide 27. Complete titration of phosphoric acid. Phosphoric acid carries
three dissociable protons, and it thus exhibits three fused titration curves.
Each dissociable proton has a progressively higher pKa. The first proton is
relatively easy to remove. The second proton is harder to remove, and
removing the third proton is even more difficult Notice that all three
titrations have the same shape, but are displaced to higher and higher pH
values as the titration proceeds.
The three pKa’s for this system are 2.1, 7.2 and 12.7. It is the middle
dissociation with a pKa of 7.2 that is of physiological relevance, because
most living cells operate around neutrality (pH 7.0) and thus this central
dissociation is the point at which phosphoric acid can serve as a buffer in
living cells.
Slide 28. The pK values for amino acids. This slide summarizes the pK
values for all of the ionizable groups in the common amino acids. The amino groups in these amino acids vary a bit in their pK a values—a good
mean to use in titrations would be 9.25. Similarly, for the -carboxyl groups
you could use a mean value of 2.25. There are seven side chain R groups
that undergo ionization in the normal pH range. The pK values of these "R"
groups vary from 3.9 to 12.5. These pK values reflect the relative ability of
the groups to donate hydrogen ions, with the best proton donors having low
pK values.
One last generalization about the ionizable amino acid side chain "R"
groups. These functional groups basically come in two flavors--those
that are neutral at lower pH and become negative as the pH is raised (that
would include the side chains of glutamate, aspartate, cysteine and tyrosine),
and those that are positive at lower pH and become neutral as the pH is
raised. (That includes histidine, argininine, and lysine.) There is no single
"R" group that can carry both negative and positive charges as a function of
pH.
Slide 29. The pH profiles for enzymes. Enzyme proteins, such as trypsin
and alkaline phosphatase, which operate at neutral pH values generally have
a pH optima in the neutral or slightly alkaline range. In contrast, pepsin
which functions in the acidic environment of the stomach has a pH optimum
around 2.0.
Slide 30. The pH values in human physiology. The pH of foodstuffs can
vary over a wide range—from pH of for carbonated beverages of 2.0 to pH
12.0 or above for certain alkaline foods.
Within the human body most tissues operate at near neutral pH levels, but
there are exceptions such as the stomach. Raw gastric juice has a pH range
of 1.0 to 2.0.
Slide 31. pH and dental caries. We have here a summary of the classic
relationship between dental plaque (my dentist used love the term plaque,
and he would bemoan the use of the term “tarter”), pH, buffering and dental
carries. The bacteria in dental plaque erode the enamel by secreting acid.
These bacteria secrete acid when they are provided with a good source of
fermentable sugar such as glucose or sucrose. The buffers in the saliva, such
as bicarbonate, tend to moderate the acidifying effects of the plaque
organisms. In this context it appears that a critical pH for enamel dissolution
is about 5.5. Below that pH, nasty things tend to happen, which I will not
describe here.
Slide 32. The pH of dental plaque after exposure to carbohydrates. The
pH of the dental plaque was measured in three patients after exposure to
galactose and glucose solutions. There was a significant drop in pH after
exposure to the galactose solution, but the lowest pH was above the critical
level at which enamel begins to dissolve. In contrast, treatment with a
glucose solution resulted in a greater drop in pH. In two of the three subjects
the pH level was at or below the point where the enamel dissolves.
Slide 33. The carbon dioxide/bicarbonate buffer system. The primary
product of carbon metabolism in humans is carbon dioxide. The carbon
dioxide is in equilibrium with carbonic acid which can dissociate to produce
bicarbonate and a hydrogen ion. The expeditious removal of carbon dioxide
by respiration is necessary to prevent acidosis.
The carbonic acid (and its conjugate base, bicarbonate), also serves as a
major buffer system in humans, helping to maintain neutral pH in the blood
and tissues.
Slide 34. The carbon dioxide/bicarbonate system buffers blood plasma
at pH 7.4. The general rule is that a buffer functions well only within one
pH unit of its pKa. The carbonic acid/bicarbonate system has a pKa of 6.1
which, on the surface, would make it a poor buffer for maintaining the
plasma pH (normally set at about pH 7.4). However, because the buffer
system is in equilibrium with carbon dioxide, which can be regulated by
respiration, the system works very well at pH 7.4.
Slide 35. Equilibration of carbon dioxide with carbonic acid and
bicarbonate ion. Gaseous carbon dioxide is in equilibrium with dissolved
carbon dioxide which in turn equilibrates with carbonic acid and bicarbonate.
-When excess hydrogen ions are produced by metabolism, they can react
with bicarbonate to form carbonic acid which is then converted to carbon
dioxide. The carbon dioxide can be removed by respiration. This is a
temporary buffering mechanism because it lowers serum bicarbonate levels.
The only way to permanently restore these bicarbonate levels is to remove
the excess hydrogen ions by secreting them in the urine.
-When excess hydroxyl ions are added, they can react with carbonic acid to
form more bicarbonate. The pool of carbonic acid can be increased by a
reduction in the rate of respiration which in turn decreases the rate of carbon
dioxide release. The retained carbon dioxide can be converted to carbonic
acid. This is also a temporary buffering mechanism because if continued
indefinitely it would lead to the buildup of bicarbonate to unacceptable
levels. Permanent compensation would require that the bicarbonate be
secreted in the urine.
Slide 36. The respiratory system removes or retains carbon dioxide.
Respiration provides a mechanism for elimination excess carbon dioxide
produced by metabolism. It can also provide a temporary mechanism for
buffering other acids (or bases) produced by metabolism. However, when
acids or bases are produced from metabolic sources other than carbon
dioxide, they must ultimately be eliminated in the urine. The body cannot
remove an excess of metabolically produced acid or base by means of
respiration.