13 | Bioenergetics and
Biochemical Reaction
Types
© 2017 W. H. Freeman and Company
Life Needs Energy
• Recall that living organisms
are built of complex
structures.
• Building complex structures
that are low in entropy is
only possible when energy is
spent in the process.
• The ultimate source of this
energy on Earth is sunlight
Laws of Thermodynamics
Apply to Living Organisms
• Living organisms cannot create energy from nothing.
• Living organisms cannot destroy energy into nothing.
• Living organism may transform energy from one form to
another.
• In the process of transforming energy, living organisms
must increase the entropy of the universe.
• In order to maintain organization within themselves,
living systems must be able to extract useable energy
from their surroundings and release useless energy
(heat) back to their surroundings.
Energy transfer in metabolism
• In order to understand the flow of energy in
biological systems (bioenergetics) we need to
determine the amount of energy associated
with ATP
• The most useful thermodynamic term for this
purpose is G, the free- energy change for a
chemical process
4
Free-energy change (G)
• Free energy (G = term defined by J. Gibbs) is a
measure of the energy available to do work at
constant temperature and pressure, which describes
the living state
• In molecular terms G is the energy stored within the
structure of the molecule and is a measure of the
amount of energy that will be released or is required
in the course of the reaction
• The change in free energy (G) is the energy change
occurring when a reaction proceeds from start to
equilibrium
• For biochemical purposes G is used to predict
whether energy is released or required during the
course of the reaction 5
Reaction is spontaneous and
exergonic:
For the reaction:
 G  0 (negative; GA  GB)
A  B
Reaction is non-spontaneous
and endergonic:
 G  0 (positive; GB  GA)
(G = GB – GA)
Reaction is at equilibrium:
 G = 0 (GA = GB)
6
Exergonic reaction (energy released)
Free energy
Reactants
Amount of
energy
released
(∆G < 0)
Energy
Progress of the reaction
7
Products
Endergonic reaction (energy required)
Free energy
Products
Amount of
energy
required
(∆G > 0)
Reactants
Energy
Products
Progress of the reaction
8
• Spontaneous reactions:
Glc + 6O2  6CO2 + 6H2O (G  0)
ATP  ADP + Pi (G  0)
• Non- spontaneous reactions:
6CO2 + 6H2O  Glc + 6O2 (G  0)
ADP + Pi  ATP (G  0)
9
• G does not say anything about the velocity of the
reaction
– The oxidation of glucose
•  sec (bomb calorimeter)
•  min (cell)
• The velocity is dependent on the activation energy
of a reaction
• Catalysts/Enzymes accelerate reactions by 
activation energy
10
Go (Standard free-energy change)
• We can define standard conditions for any
process and then use these standard
conditions as the basis for comparing
reactions
• For solutes the standard state is taken as a
1 molar (M) concentration or for gases, at
partial pressures of 101.3 kPa or 1 atm.
11
pH = - log
pH = - log [H+]
[H+]
= - log (1)
= - log (1 x 10-7)
=0
= - (-7)
=7
12
For any general reaction:
aA + bB  cC + dD
We can write an equation that relates the free-energy change (G)
for the reaction under any conditions to the free-energy change under
standard conditions (Go )
G = Go + RT ln [C]c [D]d
[A]a [B]b
= Go + RT ln Q
(Q= mass-action ratio)
13
• Go is the free-energy change under standard
conditions:
– Initial concentrations of all reactants and products = 1M or
for gases 1 atm; T = 273 + 25 oC = 298 K;
• Go is constant for a reaction
• R
= Gas constant
= 1.99 x 10-3 kcal/mol/K
= 8.31 x 10-3 kJ/mol/K
(1kcal = 4.184 kJ)
14
• T = absolute temperature
• G = free-energy change under non-standard/ any
conditions and is dependent on Go (nature of the
reactants/products) and the actual intracellular
concentrations of reactants/products (given by the
second term of the equation)
• [ ] = molar concentrations of the products/reactants
(A, B, C, D)
15
Modified standard state for biochemical
reactions
• Since biochemical reactions do not occur
naturally at a hydrogen-ion concentration of 1
M, a biochemical standard state (Go’ = delta
G zero prime) is used, where the [H+] is set at
1 x 10-7 M
G’ = Go’ + RT ln [C]c [D]d
[A]a [B]b
16
Go’ (Standard free-energy change) of a
reaction and its relation to the equilibrium
constant
• At equilibrium G’ = 0
0 = Go’ + RT ln [C]c [D]d
[A]a [B]b
Go’ = - RT ln [C]c [D]d
[A]a [B]b
17
• The concentrations are now equilibrium
concentrations and the eq. can be written as:
Go’ = - RT ln K’eq
• If the concentrations of reactants and
products at equilibrium can be determined,
the Go’ can be calculated
18
Free Energy, or the Equilibrium Constant,
Determines the Spontaneity of Processes
TABLE 13-2
Relationship between Equilibrium Constants and Standard FreeEnergy Changes of Chemical Reactions
DG'°
K'eq
(kJ/mol)
(kcal/mol)a
103
–17.1
–4.1
102
–11.4
–2.7
101
–5.7
–1.4
1
0.0
0.0
10–1
5.7
1.4
10–2
11.4
2.7
10–3
17.1
4.1
10–4
22.8
5.5
10–5
28.5
6.8
10–6
34.2
8.2
a
Although joules and kilojoules are the standard units of energy and are used throughout
this text, biochemists and nutritionists sometimes express DG'˚ values in kilocalories
per mole. We have therefore included values in both kilojoules and kilocalories in this
table and in Tables 13-4 and 13-6. To convert kilojoules to kilocalories, divide the number
of kilojoules by 4.184.
Free Energy, or the Equilibrium Constant,
Determines the Spontaneity of Processes
TABLE 13-3
Relationships among K'eq, DG'˚, and the Direction of
Chemical Reactions
DG'˚ is . . .
Starting with all components at
1 M, the reaction . . .
>1.0
negative
proceeds forward
1.0
zero
is at equilibrium
<1.0
positive
proceeds in reverse
When K'eq is . . .
Experimental determination of Go’
Glucose-6-phosphate  Fructose-6-phosphate
• To begin experiment, 1M concentrations of each reagent are mixed under
standard conditions with the enzyme glucose-6-phosphate isomerase. At
equilibrium the concentrations of [F6P] and [G6P] are 0.67 M and 1.33 M,
respectively
Keq = [F6P] = 0.67 = 0.50
[G6P] 1.33
Go = - RT ln K’eq
= - (8.31x 10-3) (298) ln 0.5
= - (- 1.72)
= + 1.72 kJ/mol
21
• Needs energy to proceed from left to right
under standard conditions
• Endergonic
• Non-spontaneous
• Net flow was from F6P to G6P
• Reaction actually proceeded in the reverse
direction
22
Energetics Within the Cell are Not
Standard
• The actual free-energy change of a reaction in the
cell depends on:
– the standard change in free energy
– actual concentrations of products and reactants
– For the reaction aA + bB
cC + dD:
[C ]c [ D]d
DG  DG '  RT ln
[ A]a [ B]b
• Thus for the reaction under non-standard conditions:
Glucose-6-phosphate  Fructose-6-phosphate
Calculate the G for the above reaction when [F6P] = 0.05
mM and the [G6P] = 1 mM at 37 deg C. Go = 1.72 kJ/molsee slide 21 for calculation. NB! Concentrations need to be
converted from mM to M (x10-3).
G = Go + RT ln [F6P]
[G6P]
1.72 + (8.31x10-3)(310) ln 0.05 x10-3
1x10-3
= - 5.99 kJ/mol
=
Under these conditions the reaction will proceed from left to
24
right to equilibrium
Metabolism Is the Sum of All Chemical
Reactions in the Cell
Metabolic pathway
• Metabolic pathway = a series of enzymecatalyzed reactions
• Pathway = linear, branched, cyclic
• Anabolism (biosynthesis of larger molecules
from smaller ones) vs catabolism (breakdown
of larger molecules to smaller ones)
26
Metabolic reactions may be divided into 2
types:
• Near-equilibrium reactions (reversible
reactions): Go  0
– Add reactants/remove products  reaction shifts
to the right
– Add products/remove reactants  reaction shifts
to the left
– Reactions in the middle of Pathways
– The same E for forward/reverse reaction
27
• Metabolically irreversible reactions:
G o  0
–
–
–
–
Start/end of pathways
Control points
Different enzymes catalyze forward/reverse reactions
These enzymes are usually allosteric enzymes (See later
BCM251)
•
•
Allosteric enzyme: multisubunit protein; a conformational
change in one subunit induces a change in another subunit
Allosteric effector: a substance (activator/inhibitor) that binds
to an allosteric enzyme and affects its activity
28
A
E7

E1
Irreversible reaction
Go’ < 0
B

E2
C

E3
Reversible reactions
G0’ ~ 0
D

E4
E
E6

E5
F
29
Irreversible reaction
Go’ < 0
Chemistry of metabolism
• Thousands of cellular reactions
• Reactions can be sorted in 6 main
types/categories
30
TYPE OF REACTION
Oxidation-reduction
(electron transfer)
ENZYME CLASS
Oxidoreductases
(eg. Dehydrogenases,
reductases)
Group transfer
Transferases
(transfer of functional groups) (eg. Kinases transfer
phosphoryl groups)
Hydrolases
Hydrolytic cleavage
(cleavage of bond by addition (eg. Glycosidases,
of water = hydrolysis)
peptidases)
Non-hydrolytic cleavage
(cleavage of bond; water is
not involved)
Lyases
(F1,6BP  DHAP + GAP)
31
Isomerization and rearrangement
Isomerases
*Epimerases (UDP-Glc  UDPGal)
*Isomerases (Aldose  Ketose)
* Mutases (G1P  G6P)
Bond formation using
energy
Ligases
*Synthetase (ATP provides
energy)
* Synthase (ATP does not
provide energy directly)
32
Phosphoryl Transfer from ATP
ATP is frequently the donor of the phosphate in the
biosynthesis of phosphate esters.
(- H2O)
6. Acid + Acid
R’-COOH
O
II
- O - P- OH
I
O-

HOOC- R”
O
O
R’- C- O- C – R”
ACID ANHYDRIDE
O
O
II
II
- O - P- O - P- OI
I
O- O-
O
II
HO - P- OI
O-
PHOSPHORIC ANHYDRIDE
34
Hydrolysis of ATP Is Highly Favorable
Under Standard Conditions
• Better charge separation
in products
• Better solvation of
products
• More favorable resonance
stabilization of products
DG of ATP Hydrolysis Is
Mg++ Dependent
Actual DG of ATP Hydrolysis
Differs from DG’
• The actual free-energy change in a process depends on:
– the standard free energy
– the actual concentrations of reactants and products
• The free-energy change is more favorable if the reactant’s
concentration exceeds its equilibrium concentration.
• The true reactant and the product are Mg-ATP and Mg-ADP,
respectively.
– DG also
Mg++
dependent
[MgADP  ]  [Pi ]
DG  DG ' RT ln
[MgATP 2 ]
TABLE 13-5 Total Concentrations of Adenine Nucleotides, Inorganic
Phosphate, and Phosphocreatine in Some Cells
Concentration (MM)a
ATP
ADPb
AMP
Pi
PCr
Rat hepatocyte
3.38
1.32
0.29
4.8
0
Rat myocyte
8.05
0.93
0.04
8.05
28
Rat neuron
2.59
0.73
0.06
2.72
4.7
Human erythrocyte
2.25
0.25
0.02
1.65
0
E. coli cell
7.90
1.04
0.82
7.9
0
a
For erythrocytes the concentrations are those of the cytosol (human erythrocytes lack a nucleus and
mitochondria). In the other types of cells the data are for the entire cell contents, although the cytosol and
the mitochondria have very different concentrations of ADP. PCr is phosphocreatine, discussed on p. 516.
b This value reflects total concentration; the true value for free ADP may be much lower (p. 509).
Cellular ATP concentration is usually far above the equilibrium
concentration, making ATP a very potent source of chemical energy.
Several Phosphorylated Compounds Have
Large DG’ for Hydrolysis
• Again, electrostatic repulsion within the reactant
molecule is relieved.
• The products are stabilized via resonance, or by more
favorable solvation.
• The product undergoes further tautomerization.
Phosphates: Ranking by the Standard
Free Energy of Hydrolysis
Phosphate can be transferred from compounds with higher ΔG’
to those with lower ΔG’.
Reactions such as:
PEP + ADP =>
Pyruvate + ATP
are favorable and
can be used to
synthesize ATP.
Hydrolysis of Thioesters
• Hydrolysis of thioesters is
strongly favorable.
– such as acetyl-CoA
• Acetyl-CoA is an important
donor of acyl groups.
– feeding two-carbon units
into metabolic pathways
– synthesis of fatty acids
• In acyl transfers, molecules
other than water accept
the acyl group.
Acid + Thiol
(-H2O)

O
R’ – C – S – R”
R’-COOH
HS- R’’
THIOESTER
42
Electrons are transferred in different
ways
•
Directly as electrons
Fe2+ + Cu2+  Fe3+ + Cu+
•
As Hydrogen atoms (H = H+ + e-)
AH2  A + 2H+ + 2e- (ox)
B + 2H+ + 2e-  BH2 (red)
AH2 + B  A + BH2
•
As a hydride ion (:H- = H+ + 2 e-)
This happens with NAD+ -linked dehydrogenases
43
Oxidation-Reduction Reactions
Reduced organic compounds serve as fuels from which
electrons can be stripped off during oxidation.
Reversible Oxidation of
a Secondary Alcohol to a Ketone
• Many biochemical oxidation-reduction reactions
involve transfer of two electrons.
• In order to keep charges in balance, proton transfer
often accompanies electron transfer.
• In many dehydrogenases, the reaction proceeds by a
stepwise transfers of proton (H+) and hydride (:H–).
Coenzymes in redox reactions:
Substrate  H2  Substrate + 2H+ + 2 e(red)
(ox)
Coenzyme + 2H+ + 2 e-  Coenzyme  H2
(ox)
(red)
NAD+
NADP+
FAD
FMN
NADH + H+
NADPH + H+
FADH2
FMNH2
46
NAD+ and NADP+ Are
Common Redox Cofactors
• These are commonly called pyridine nucleotides.
• They can dissociate from the enzyme after the
reaction.
• In a typical biological oxidation reaction, hydride
from an alcohol is transferred to NAD+, giving
NADH.
NAD+ /NADH (Nicotinamide adenine
dinucleotide) & NADP+ /NADPH (Nicotinamide
adenine dinucleotide phosphate)
• Niacin/ Vit B3 (vitamin precursor)
• Adenine
Nicotinamide (Pyridine ring)


Ribose
Ribose


Phosphate  Phosphate
• ADP in structure
48
NAD+ and NADP+ Are
Common Redox Cofactors
Formation of NADH Can Be Monitored
by UV-Spectrophotometry
• Measure the change of absorbance at 340 nm
• Very useful signal when studying the kinetics of
NAD-dependent dehydrogenases
FAD/FADH2 (Flavin adenine dinucleotide)
& FMN/FMNH2 (Flavin mononucleotide)
• Riboflavin/Vit B2 (Vitamin precursor)
• FMN = Flavin

Ribitol

Phosphate
• FAD = Flavin

Ribitol

Phosphate 
Adenine

Ribose

Phosphate
51
• ADP in structure
• FMN/FMNH2 is an electron acceptor/donor in the
electron transport chain:
SH2 + FMN  S + FMNH2
• FAD = is an oxidizing agent in catabolism:
SH2 + FAD  S + FADH2
52
Flavin Cofactors Allow
Single Electron Transfers
• Permits the use of molecular oxygen as an ultimate electron
acceptor
– flavin-dependent oxidases
• Flavin cofactors are tightly bound to proteins.