CHAPTER 13 - OVERVIEW OF METABOLISM
Introduction
Last semester you learn about the structure of proteins, carbohydrates and membranes.
During this semester we’ll see how these molecules, as components of cells and organisms,
extract energy from the environment and use it to carry out cellular work. The overall process
by which organisms acquire and use free energy comprises metabolism, which is the sum of all
the chemical transformations which occur in an organism, and which is typically divided into
two parts, catabolism (breakdown of nutrients to generate energy - exergonic), and anabolism
(biosynthesis, an energy-requiring process (endergonic) in which biomolecules are synthesized
from simpler precursors).
Different organisms have different nutritional requirements depending on their ecological
niche:
Autotrophs synthesize cellular components from simple molecules such as H2O, CO2,
NH3 and H2S.
Chemolithotrophs extract free energy from oxidation of inorganic chemicals
such as NH3, H2S or Fe+2:
2 NH3 + 4 O2 6 2HNO3 + 2 H2O
H2S + 2 O2 6 H2SO4
4 FeCO3 + O2 + 6 H2O 6 4 Fe(OH)3 + 4 CO2
NOTE: O2 is the oxidizing agent, or electron acceptor in all cases. The transfer of electrons in
cells between nutrient and O2 is never a direct transfer, but typically involves an intermediate
carrier such as the coenzymes NAD+ and FAD.
Phototrophs extract energy from light via photosynthesis:
6 CO2 + 6 H2O + light energy 6 6 O2 + C6(H2O)6
Heterotrophs obtain free energy from the oxidation of organic compounds:
C6(H2O)6 + 6 O2 6 6 CO2 + 6 H2O
NOTE: light energy provides the means to reverse the spontaneous flow of electrons from O2 to
H2O in the reactions of photosynthesis. This implies an important symbiotic relationship
between phototrophs (plants) and heterotrophs (us).
O2
(C H2O)n
Heterotrophs
Phototrophs
CO2
H2 O
NOTE: In the above examples O2 was the electron acceptor. Organisms which utilize O2 in this
way are aerobic (obligate - must use O2, or facultative - can also grow in absence of O2).
Obligate anaerobes, in contrast must not be exposed to O2).
Metabolic Pathways
The reactions of metabolism are organized, somewhat arbitrarily, into pathways, either
linear or cyclic. Pathways tend to occur in specific cells or cellular compartments, depending on
their type, so the designation is not entirely arbitrary. For example O2-utilizing reactions occur
in mitochondria (CAC, ET-OxPhos, etc). Chapter 21 will deal with the interdependence of
metabolic functions of various organs.
Catabolic pathways break down larger metabolites into smaller molecules. They often
converge to a single metabolite, such as Acetyl CoA. (See Figure 13-2)
The reactions depicted in Figure 13-2 are all catalyzed by enzymes. These proteins, by virtue of
their substrate specificity, prevent the inefficient formation of byproducts, greatly enhance the
rate of metabolic reactions, and also are capable of rendering a nonspontaneous, endergonic
reaction (delta G > 0) spontaneous by coupling it with an exergonic reaction such as the
hydrolysis of ATP.
All metabolic reactions fall into four major types and are catalyzed by the six classes of
enzymes (see Table 11- 2, page 315): Redox reactions (catalyzed by oxidoreductases), grouptransfer reactions (catalyzed by transferases and hydrolases), eliminations, isomerizations and
rearrangements (catalyzed by isomerases and mutases), and reactoins that make or break C-C
bonds (catalyzed by hydrolases, lyases and ligases).
NH2
N
N
N
N
O H CH 3
O
HS
CH 2 CH 2 NH C
CH 2 CH 2 NH C C C
CH 2 O P O
OH CH 3
Business End
O
O
O
O
CH2
P O
O
H
Pantothenate
(Vitamin)
O
H
H
O
OH
P
H
OH
OH
Nucleotide
HS CoA
O
H3 C C
Ther
S CoA = Acetyl CoA
mody
namics - A major objective of catabolism is to generate energy necessary for biosynthesis, active
transport, mechanical work, etc.; we need to review our brief introduction to thermodynamics
last semester in order to understand the energetics of metabolism.
Recall that the second law of provides us with a thermodynamic criterion for reaction
spontaneity, namely the change in entropy, delta S of the universe. The Gibbs free energy
change, delta G, is a more convenient measure, since it involves only the system of interest.
Delta G = Delta H - TDeltaS (H = enthalpy (heat) and S = entropy). The criteria we need are
summarized below:
- Delta G < 0, reaction is spontaneous (thus irreversible) and proceeds left to right
- Delta G > 0, reaction is not spontaneous, thus proceeds right to left
- Delta G = 0, reaction is at equilibrium
- Recall that there are both enthalpic (H) and entropic components of )G. Processes that
are neither enthalpically or entropically favored will never proceed spontaneously; processes that
are both enthalpically and entropically favored will always proceed spontaneously, processes
that are not favored enthalpically but are favored entropically will proceed only if temperature is
raised to a point such that the entropic component overrules the enthalpic component (example heating up a solution to dissolve the solute)
For a chemical reaction A + B WC + D, the equation for )G becomes
Delta G = DeltaG0' + RT lnQ, Q = [C][[D]/[A][B]
(1)
where Q has the appearance of an equilibrium constant, but is not numerically equal to the
equilibrium constant unless concentrations are equilibrium concentrations. At equilibrium,
DeltaG = 0 and
DeltaG0' = -RT lnQeq = -RT ln Keq
(2)
Consider:
If initial concentrations are not equilibrium values but rather are such that DeltaG < 0,
then Q is not numerically equal to K, and the reaction will proceed spontaneously, left to right as
dictated by the second law of thermodynamics. As the reaction proceeds left to right,
concentrations of reactants will decrease and concentrations of products increase, increasing Q
and also ln Q, until the RT lnQ term in equation (1) offsets the DeltaG0' term such that DeltaG =
0. At this point the reaction is at equilibrium. If initial conditions are such that DeltaG > 0, the
reaction will proceed right to left, Q and RT lnQ will get smaller, etc.
NOTE: An important feature of G is that it is a function of state; that is, )G is pathindependent. This implies that )G values are additive, which implies that coupling of
unfavorable reactions with favorable reactions is possible. Consider, for example, the following:
A + B 6 C + D,
DeltaG1 < 0
D + E 6 F + G,
DeltaG2 < 0
Reaction 1 is exergonic, reaction 2 endergonic. If an appropriate mechanism exists (typically, a
common intermediate between reactions 1 and 2, then it is possible to couple the reactions:
A + B + E 6 C + F + G,
)G = )G 1 + )G2
A common strategy in metabolism is to couple an otherwise unfavorable (endergonic) reaction
with an exergonic reaction (i.e., ATP hydrolysis) such that the overall DeltaG is < 0.
Now recall the difference between thermodynamics and kinetics. Kinetics has to do with
rates of reactions, whereas thermodynamics can tell us nothing about rates, but only the position
of ultimate equilibrium.
Many of the individual reactions in a typical metabolic pathway are near equilibrium
(i.e., DeltaG -0 and Q - K). The net flux, J, (J = vf - vr, where vf and vr are forward and reverse
rates, respectively ) of metabolites through such a near-equilibrium reaction is -0, hence the
concentrations of reactants and products don’t change appreciably.
Enzymes catalyzing near-equilibrium reactions are typically very efficient. Enzymes
lower the activation barrier between product and reactant for all reactions such that both vf and vr
are increased. This means that both vf and vr are essentially limited only by the amounts of
reactants and products, and the net flux of such a reaction is said to be under thermodynamic
control. If for some reason [reactants] increases relative to [products] J will increase
correspondingly, and vice versa. It would be clearly impossible for the organism to exert direct
regulation, or control, over a near-equilibrium reaction.
Other metabolic reactions are highly exergonic and operate far from equilibrium. It is the
inevitable existence of such reactions that render metabolic pathways irreversible (because the
reverse reaction is endergonic and thus will not occur spontaneously). It is via exergonic
reactions that biological regulation, or control, is typically exerted, and not through nearequilibrium reactions. Thermodynamics favor formation of products for exergonic reactions, but
enzymes catalyzing these reactions are typically inefficient, resulting in large activation barriers,
and reactants therefore tend to accumulate such that such exergonic reactions typically constitute
the rate-determining step of a pathway. Such a reaction typically occurs early in the pathway,
and is also known as the committed step, because it “commits” its product to continue down the
pathway. Committed steps frequently occur at branch points in metabolic pathways. Changes in
concentrations of reactants and products have little effect on the flux through these reactions; the
net flux, J, is controlled by regulating the activity of the enzyme catalyzing this reaction. Such
enzymes are typically allosteric and subject to modulation via allosteric modulators, which
serve as regulators. Note that in contrast to near-equilibrium reactions, which are typically under
thermodynamic control, highly exergonic reactions are said to be under kinetic control. Nature
does not let metabolism proceed uncontrolled. When the metabolic goal of a pathway is
accomplished, the final product can serve as a negative modulator and reduce the activity of the
enzyme catalyzing a highly exergonic, perhaps committed, step “upstream.” This constitutes
feedback inhition.
The existence of a single, highly exergonic reaction in a pathway will confer
directionality to the pathway (it is irreversible). The net flux through a pathway will typically not
be zero, but a positive, constant value. This corresponds to a steady-state (as opposed to an
equilibrium) condition.
Since a complementary, biosynthetic, pathway will typically exist between the same start
and end points, anabolic pathways must differ at least in some respects from the corresponding
catabolic pathway. At the very least, a highly exergonic reaction in the catabolic pathway must
by bypassed such the overall )G of the anabolic pathway is > 0.
Some other mechanisms of metabolic control (in addition to allosteric control) are control
by covalent modification, substrate cycles, and genetic control. Note that in a substrate cycle
(see page404) the forward and reverse reactions are not catalyzed by the same enzyme. In such
cases a modulator typically exists that increases (or decreases) vf, but decreases (or increases) vr
(i.e., the same modulator has reciprocal effects on the forward and reverse reactions). The net
flux through such a substrate cycle can then be more efficiently controlled than through a near
equilibrium reaction in which both forward and reverse reactions are catalyzed by the same
enzyme.
“High Energy” Compounds
Consider: We have seen that since G is a function of state, Delta G values are additive,
implying that energy coupling is possible, giving rise to an important means of getting
unfavorable reactions to occur. Successful coupling of two reactions requires an efficient
coupling mechanism, such as the occurrence of a common intermediate, as suggested above.
Biological coupling mechanisms typically involve the participation of high energy molecules,
such as ATP. Consider the complete oxidation of glucose:
C6H12O6 + 6 O2 6 6 CO2 + 6 H2O, )G0' = -2850 kJ/mole
In order for this energy to be conserved to perform useful work, it must first must be stored as
chemical energy in high energy compounds such as ATP (figure 3, p.406)
NH2
N
N
OH
OH
OH
N
N
HO P O P O P OCH2 O
O
O
O
H
H
H
H
OH OH
A high energy biomolecule like ATP has a large - Delta G of hydrolysis, which is due to
destabilization of reactant and/or stabilization of product. For ATP hydrolysis,
ATP + H2O WADP + Pi
ATP is destabilized by electrostatic repulsion and resonance destabilization, whereas Pi, one of
the hydrolysis products is resonance stabilized:
HO
O
O
O
P O
P O
P OR
O
O
O
ATP
O
HO
P O
O
Pi
The oxygen sandwiched between two P’s acquires a positive charge if it forms a double bond.
The above illustration therefore is an unstable resonance form for ATP, whereas any of the four
O’s in Pi can be involved in a double bond (resonance stabilization of product).
Greater solvation of products also plays a role in the large free energy gap for ATP
-Other examples of high energy compounds.
- Phosphoenolpyruvate
O
O
C
C
OPO3
CH2
O
C
H C
O PO3
OH
CH2OPO3
- Phosphocreatine
O
C
O
NH2
O
H2C N C HN P O
CH3
O
NOTE: The absolute value of Delta G0' of hydrolysis values for high energy, phosphorylated
intermediates (see table 13-2) is called the phosphoryl group transfer potential.
Example (#, p.426): Predict whether the creatine kinase reaction will proceed in the direction of
ATP synthesis or phosphocreatine synthesis at 250C when [ATP] = 4 mM, [phosphocreatine] =
2.5 mM, [ADP] = 0.15 mM, [creatine] = 1 mM
phosphocreatine + ADP W ATP + creatine
This is a typical coupling problem where, depending on concentrations, phosphocreatine
transfers a phosphate group to ADP to form ATP and creatine, or vice versa, depending on the
sign of Delta G:
DeltaG = Delta G0' + RT lnQ, Q = [ATP][creatine]/[phosphocreatine][ADP]
From Table 13-2, )G0' = -43.1 + 30.5 kJ/mole = -12.6 kJ/mole
Delta G = Delta G0' + RT ln[4][1]/[0.15][2.5] = -12.6 + 5.9 = -6.7 kJ/mole
Thus the reaction proceeds spontaneously from left to right, in the direction of ATP synthesis.
Example: We saw earlier that several catabolic pathways converge to acetyl CoA, which
contains a high-energy thioester linkage between CoA and the acyl group. Under standard
conditions hydrolysis of ATP to form ADP does not provide the necessary driving force for
synthesis of acetyl CoA from CoA and acetate, and ATP is instead hydrolyzed to AMP and Ppi,
which is then hydrolyzed to 2 Pi, effectively providing twice the free energy as would be
obtained by hydrolyzing ATP to ADP:
Acetate + CoA + H+ W Acetyl CoA + H2O, )G0' = +33.1 kJ/mole
ATP + H2O W AMP + PP i , )G0' = -32.2 kJ/mole (Table 13-2)
PP i + H2O W 2 Pi , )G0' = -33.5 kJ/mole (Table 13-2)
Acetate + CoA + ATP + H+ WAcetyl CoA + AMP + 2 P i
)G0' = + 33.1 - 32.2 - 33.5 = -32.6 kJ/mole
Summary: We have seen that the hydrolysis of ATP to ADP constitutes an efficient
thermodynamic driving force for otherwise unfavorable reactions. The )G0' of hydrolysis is
large and negative, qualifying ATP as a “high-energy” biomolecule; however, Table 13-2
indicates that there are several other high-energy molecules with a larger, negative )G0', such as
phosphoenolpyruvate and phosphocreatine. The position of ATP roughly in the middle of this
table implies that ATP can transfer (under standard conditions) a phosphoryl group to glucose,
for example, but also that ADP can accept a phosphoryl group from phosphoenolpyruvate to
form ATP. ATP thus has an intermediate phosphoryl group-transfer potential which enhances
the versatility of ATP.
Note: ATP is not the only high energy molecule used by cells as a driving force. GTP, NTP,
for example, are also used. These are interconverted via the action of nucleoside diphosphate
kinases:
ATP + GDP W ADP + GTP
Consider also the adenylate kinase reaction:
AMP + ATP W ADP + ADP
Redox Reactions Consider the oxidation of Glucose
C6H12O6 + 6 O2 6 6 H2O + 6 CO2
Note oxidation states of C and O.
Features:
Electrons are rarely transferred directly from substrate to O2 (recall toxicity of O2
requires 4 electrons).
Major intermediate electron carriers are NAD+ and FAD
H
H
H
O
H
C NH2
N
H
C NH2
H
H
H O
C NH2
H
O
N
H
R
H
H
N
H
R
R
NAD
NADH
H
O
H 3C
N
C
H 3C
N
N
H
H
N
C
O
H
O
H3 C
N
C
H 3C
N
N
R
H
R
FAD
H
N
C
O
FADH
2
Electrons are transferred spontaneously from substrate (glucose) to NAD+ or FAD to O2. The
electrochemical potential is the measurable parameter which is related to )G, thus reaction
spontaneity through the relationship )G = -nF)E, where n is the number of electrons transferred
and F is the Faraday (96,485 J/(mole Volt). Thus a positive )E is indicative of a spontaneous
transfer of electrons
Half-reactions in Table 13-3 are written as reductions, hence the term “reduction
potentials”
As with the use of Table 13-2 for various phosphorylation reactions, two entries are
selected from this table to form a complete redox reactions, one of which must be turned around
to form an oxidation (in this case change the sign of E0'). Spontaneous reactions are those which
correspond to a positive Delta E0' (standard conditions).
Note that Fe+3 ions of the various cytochromes (similar to hemes) have significantly
different reduction potentials indicating that proteins are capable of modulation the affinity of
these groups for electrons.
Recall that NADH and FADH2 are electron-rich redox coenzymes (strong reducing
agents) which will pass their electrons on through the electron-transport chain in mitochondria to
produce ATP (oxidative phosphorylation).
Skip section 4.
Problems: 1, 2, 3, 4, 6, 9, 10, 13