Chapter 6 Notes
6.1: The Flow of Energy in Living Systems
I.
Two states of energy
A. Kinetic – energy of motion
B. Potential – stored energy
NOTE: Much of the work that living organisms carry out involves transforming potential energy
into kinetic energy
II.
Energy can be measured in many ways
A. Heat most convenient way of measuring energy because all other forms can be converted
into heat. Thermodynamics means “heat changes”
Most common unit of heat in biology is kilocalorie (kcal); 1 kcal = 1000 calories
NOTE: 1 calorie of heat is required to raise temperature of 1 gram of water 1 degree C.
III.
What happens to all the energy from the Sun?
A. Approximately 40 million billion calories per second (13 x 10 23 calories per year)
B. Plants, algae, and certain bacteria capture only a fraction of this energy through
photosynthesis
C. Photosynthesis Preview: Energy from the Sun is used to combine small molecules, CO2 and
H2O, into more complex molecules, C6H12O6. Converts Carbon from inorganic to organic
form. Energy from Sunlight is stored as potential energy in the covalent bonds between
atoms in the sugar molecules. Breaking bonds between atoms requires energy of C-H bonds.
Example: Fat molecules have many C-H bonds, breaking those bonds provides lots of energy.
IV.
TEST: Explain the relationship between electrons and energy.
A. During oxidation an electron is lost; atom that gains electron during reduction has a higher
energy level than the oxidized form.
B. The transfer of electrons releases energy that can be used to do the work of the cell. In
addition, electrons are needed to produce ATP and ATP can be thought of as energy
storage.
6.2: Laws of Thermodynamics and Free Energy
I.
TEST: Laws of Thermodynamics govern all energy changes in the universe.
A. First Law of Thermodynamics:
 Deals with the quantity of energy in the universe; this quantity is limited.
 States: Energy can neither be created nor destroyed; but can be transformed from one
state to another (one form to another). Ex: potential energy to kinetic energy.
 Chemical potential energy stored in some molecules can be shifted to other molecules
and then stored in different chemical bonds.
 During each energy conversion, some energy is lost as heat.
 Energy stored in bonds of macromolecules can be liberated through metabolism.
B. Second Law of Thermodynamics:
 States: The disorder in the universe, more formally known as entropy, is continuously
increasing. Tendency toward disorder/chaos.
 Simplified: Disorder is more likely than order OR “entropy increases.”
 EXAMPLE: It is more likely for a column of bricks to fall over than for a pile of bricks to
spontaneously form a column.

II.
At the formation of the universe, it held all the potential energy it will ever have. It has
become increasingly disordered with every energy exchange increasing the amount of
entropy.
Free Energy:
A. Energy available to do work in any system.
B. For molecules within a cell, where pressure and volume do not change, free energy is
denoted G (Gibb’s Free Energy). G is equal to the energy within the molecule’s chemical
bonds, called enthalpy (H) together with energy term (TS) related to degree of disorder in a
system, where S is the symbol for entropy and T is absolute temperature in K (K = C°+ 273).
The EQUATION for calculating free energy is therefore is G = H – TS.
C. In chemical reactions we break some bonds in the reactants and form new ones in the
products.
D. These chemical reactions produce changes in the free energy (G).
E. When reactions in living systems occur under conditions of constant temperature, pressure,
and volume, the change in free energy is symbolized by  (delta) …..so change in free energy
is G = H - TS
F. Can use G to predict whether chemical reactions are spontaneous or not.
G. TEST: If G is positive (G>0), then products of reaction contain more free energy than the
reactants.
 These reactions are endergonic because they require an “input of energy.”
 G is positive because high energy products can be used for work.
 Not spontaneous
H. TEST: If G is negative (G<0) then reactants of a reaction have more free energy than the
products. Energy is released.
 The reactions are exergonic because excess free energy has been released as heat
(“outward energy”).
 These reactions are spontaneous. (NOTE: not instantaneous)
I. Chemical reactions are reversible. A reaction that is exergonic in the forward direction will
be endergonic in the reverse direction
(input)
EXAMPLE:
6CO2 + 6H2O + energy (sunlight)
C6H12O6 + 6O2 (Photosynthesis)
(released)
C6H12O6 + 6O2
6CO2 + 6H2O + energy (Cellular Respiration)
Which can also be written…
6CO2 + 6H2O + energy (sunlight)
C6H12O6 + 6O2 + energy

III.
Photosynthesis is an _________________ reaction, because energy is
______________________ ____________________________.
 Cellular Respiration is a _______________ reaction, because energy is
_____________________ _____________________________.
Spontaneous chemical reactions require activation energy
A. Since exergonic reactions tend to occur spontaneously, why haven’t all such reactions
already occurred? Let’s consider the gas in your car. Your gas tank does not spontaneously
explode. Why? Oxidation!
 (Hold up! What is oxidation? _________________________. And, the
loss, or transfer, of an electron is associated with the release of
_____________, right? Wait…which of the types of reactions is
associated with a release of __________? ____________. AHAH!
Sooooo, oxidation can be said to be an _______________ reaction.
Now, where were we? Ah, yes! The gas in your car.)
 One reason all exergonic reactions have NOT already spontaneously occurred is that
most reactions require an input of energy to get started. (Gas in your car is ignited by
an electrical spark generated in the engine cylinders. This electrical spark produces a
controlled explosion.)
B. On a biological level, before new chemical bonds can be formed, even those bonds with less
energy, existing bonds must first be broken, and that requires energy input.
C. Extra energy needed to destabilize existing chemical bonds and initiate chemical reactions is
called activitation energy. Simplified: energy required to start a reaction.
D. Even Exergonic reactions have activation energy. The larger activation energy reactions are
slower, because fewer molecules succeed in getting over the initial energy hurdle. But, rate
of reactions, the speed at which they proceed, can be increased in two ways:
 Increase energy of reacting molecules (How could this be accomplished? Add heat,
heat up reactants. What does adding heat do to the molecules?
 Lower activation energy (catalyst) →adding stress to a chemical bond makes it easier
to break. Catalysis is the process used to influence chemical bonds in a way that
lowers the activation energy needed to initiate a reaction.
E. A catalyst is a substance used to speed up a reaction by lowering the activation energy
required. EXAMPLE: enzymes. (Refer to Fig. 6.5 on page 111)
J. How Catalyst Work
 Affect the transition state in a reaction.
 Energy needed to reach transition state is activation energy.
 Stabilize transition state, lowering the activation energy.
 O NOT make endergonic reactions spontaneous
 Accelerates both forward and reverse reactions by the same amount
 DO NOT alter proportion of reactant ultimately converted into product.
 Reduce the energy barrier that prevents a reaction from proceeding.
 In living systems, enzymes are the catalyst.
6.3: ATP the Energy of Cells
I.
ATP
A. Nucleotide ( see chemical structure pg 112, Fig. 6.6)
 50carbon sugar, ribose
 Adenine: organic molecule made of 2 carbon-nitrogen rings; each N has an unshared
pair of electrons and weakly attracts H ions; forms weak base
 Chain of 3 phosphates (review Fig. 6.6 on page 112 and identify the location of high
energy bonds in ATP)
B. Function:
 Energy Storage
i. Key to energy storage – triphosphate group; similar to a “coiled spring” with
phosphates straining away from one another.
ii. Phosphate groups are highly negatively charged, thus strongly repel one
another. This electrostatic repulsion makes the covalent bonds unstable.
II.
C. Activation Energy:
 ATP has low activation energy which means its bonds are easily broken by hydrolysis.
 When phosphate groups break, they release considerable amount of energy. So,
hydrolysis of ATP has a –∆ G(exergonic)
 The energy released can be used for work in the cell.
D. Steps of Hydrolysis of ATP:
 Outermost, high-energy phosphate bond hydrolyzed
 Phosphate group on end is cleaved (broken off from ATP)
 ATP then becomes ADP + an inorganic phosphate (Pi)
 Energy released = 7.3 kcal/mol
E. ATP hydrolysis drives endergonic reactions
 Endergonic reactions NOT spontaneous
 ATP cleavage of outer phosphate group releases more energy than other, endergonic
reactions
 Two reactions coupled so energy released by ATP hydrolysis supplies energy needed
by endergonic reaction
 Together -∆G (net release of energy) and therefore exergonic and proceed
spontaneously
 Most reactions in cells require less energy than that released by cleavage of ATP.
 ATP provides most energy the cell needs
ATP Cycle
A. Continuous cycle
B. Not good as a long-term energy storage molecule due to instability of phosphate bonds
C. Fats and carbohydrates better as long-term energy storage molecules
D. Cells don’t maintain more than few seconds’ worth of ATP at any given time
E. Cells continually produce ADP and Pi (see fig. 6.7)
 Exergonic reactions synthesize ATP from ADP and Pi
 Hydrolyze ATP to provide energy needed for endergonic reactions (EX: muscle
contractions)
6.4: Enzymes – Biological Catalysts
I.
Enzymes Affect Activation Energy
A. Enzymes stabilize temporary association between substrates. (Substrates are molecules that
will undergo a chemical reaction.)
B. Decreases activation energy required for new bonds to form by bringing substrates together
in correct orientation OR stressing particular chemical bonds.
II.
Characteristics of Enzymes
A. Enzymes are not changed or consumed by chemical reactions
B. Enzymes are needed in small amounts
C. Enzymes are used over and over
D. Thousands of differed enzymes
E. Each catalyzes one or a few specific chemical reactions
F. Facilitating particular reactions, they determine the course of metabolism, the total of all
chemical reactions, in that cell
G. Different cells need different enzymes. The difference contributes to structural and
functional variations among cells
IV.
V.
VI.
VII.
VIII.
Enzyme Active Sites
A. Active sites conform to fit shape of substrates
B. Active site is defined as one or more pockets or clefts on the surface of an enzyme where
substrates can bind to the enzyme. Catalysis can only occur if substrate fits precisely into
active site.
C. Enzymes, which are globular proteins, are induced by substrates to adjust shape slightly,
leading to a better induced fit.
Enzyme Cycle
D. Substrates bind with enzymes to form enzyme-substrate complexes. ( Fig 6.10 on pg 114)
E. Amino acid side groups of enzymes are very close to substrates at the active sites, and
chemically interact with substrate, destabilizing a particular bond and lowering activation
energy.
F. Substrates are broken apart; new bonds formed
G. Substrates converted to products
H. Substrates dissociate from enzyme
I. Enzyme ready to bind next substrate
J. Enzyme cycle begins again.
Enzyme Locations
A. Many are found suspended in the cytoplasm
B. Some are found within the cellular membrane
C. Some are found within organelles (Lysosome, Peroxisome)
D. Form multi-enzyme complexes to carry out reaction sequences
E. Some may consist of RNA rather than being only protein
TEST: Ribozymes (Review Fig. 6.12, pg 116)
A. RNA molecules that act as catalyst
B. Accelerate rate of particular biochemical reactions
C. Extraordinary substrate specificity
D. Two kinds
a. Intramolecular – folded structures catalyze reactions on themselves
b. Intermolecular – act on other molecules without being changed
E. EXAMPLE: Ribosomes are Ribozymes
F. Evolutionary: Discovery of the ribozyme points to possibility RNA evolved before proteins and
may have catalyzed the formation of the first proteins.
TEST: Factors Affecting Enzyme Function
A. Concentration of substrate and enzyme
B. Temperature (Fig. 6.12a)
1. Uncatalyzed reactions increase when heat is added due to increased molecular motion.
2. Rate of catalyzed reaction increases with addition of heat, but only to optimum
temperature.
a. Below optimum temperature - enzyme’s shape not flexible for induced fit that is
optimum for catalysis.
b. Above optimum temperature - bond forces too weak to maintain shape against
random movement of atoms in enzyme.
c. At higher temperatures, enzymes denature (unfolding of protein tertiary structure)
3. Most human enzymes’ optimum temperature is normal body temperature (35 to 40°C
range)
4. Optimum temperature = normal temperature of organism
IX.
X.
C. pH (Fig 6.12b)
1. Optimum pH for enzymes is between 6 to 8
2. Involves ionic interactions between oppositely charged amino acid residues such as
glutamine acid (-) and lysine (+)
3. Interactions sensitive to hydrogen ion (H+) concentration of fluid in which enzyme is
dissolved.
4. Some, such as pepsin in the stomach, can function in high acidic environments. (stomach
pH is 2, very acidic)
Inhibitors and Activators
A. Specific substances that can bind to an enzyme and cause changes in shape and activity of
enzyme
B. Cells use inhibitors to regulate which enzymes are active or inactive at any given time
C. Increases cells efficiency and ability to control changes in characteristics during development
D. Inhibitor binds to and decreases active y of enzyme
1. Example: The end product of a biochemical pathway acts as an inhibitor of an early
reaction in the pathway. Feedback inhibition
E. Activator binds to and increases activity of an enzyme
F. Enzyme inhibition occurs in two ways: (Review Fig. 6.13)
1. Competitive Inhibitors compete with substrates for the same active site, occupy the
same active site and prevent the substrate from binding.
2. Noncompetitive Inhibitors bind to the enzyme in a location other than the active site,
changing the shape of the enzyme and making it unable to bind to the substrate.
G. Most enzymes exist as allosteric enzymes, or enzymes that can exist in either an active or
inactive state.
1. Most noncompetitive inhibitors bind to enzymes at a specific portion called the allosteric
site.
2. Allosteric sites are chemical on/off switches.
3. Binding a substance to the allosteric site can switch an enzyme between its activate or
inactive configuration.
4. Substance that binds to allosteric site thereby reducing enzyme activity is called an
allosteric inhibitor.
5. Allosteric activators are substances that activate enzymes by binding to their allosteric
sites, thereby increasing enzyme activity.
Cofactors and Coenzymes
A. Cofactors are chemical components that assist enzyme function.
1. Can be metal ions found at active sites and are involved in catalysis
2. Example: The metallic ion zinc is used by carboxypeptidase to digest proteins.
The metallic zinc ions attract electrons away from the covalent bonds, making
the bonds less stable and easier to break.
B. Coenzymes are cofactors that are nonprotein organic molecules.
1. Examples include the B vitamins (B6 and B12) or modified nucleotides
NOTE: Access the on-line textbook, section 6.4. Find the link to the B vitamin under the
section on Coenzymes. Watch the video and make sure to review the Table at the end on
the B Vitamins and their functions as coenzymes.
2. Function of Coenzymes:
a. Involved in oxidation-reduction reactions catalyzed by enzymes
b. Coenzyme serves as electron acceptor, as electron pairs pass from active site to
enzyme
c. Electrons transferred by coenzyme to different enzyme, where electrons and the
energy they bear are released to substrate in another reaction
d. Often combine with protons (H+) to form hydrogen atoms; shuttle energy in from of
hydrogen atom from one enzyme to another
6.5: Metabolism: The Chemical Description of Cell Function
I. Definition
A. Metabolism is the total of all chemical reactions carried out by an organism.
B. Chemical reactions that expend energy to build up molecules are anabolic reactions
anabolism)
C. Chemical reactions that release (harvest) energy to break down molecules are catabolic
reactions/catabolism.
II. Biochemical pathways organize chemical reactions in cells.
A. A biochemical pathway is a sequence of reactions in which the product of one reaction
becomes the substrate of the next reaction in the sequence. (Refer to Fig. 6.14)
B. Biochemical pathways are the organization units of metabolism.
C. Cells use pathways to harness energy and release it slowly so it can be used for cellular work.
Otherwise it would be like gas and a match, and all the energy would be lost as heat and
gone. This slow acquisition of energy is accomplished by redox reactions.
III. Biochemical pathway
A. Many of these sequential enzyme reactions take place in specific compartments within the
cell. EXAMPLE: Steps of the Krebs Cycle take place in the matrix of the mitochondria.
B. Believed these pathways evolved in a stepwise fashion
C. Hypothetical biochemical pathway evolved slowly through time.
D. Final reactions in pathway evolved first and earlier reactions in pathway evolved later.
E. In otherwords, biochemical pathways within organisms are thought to have evolved one step
at a time, backwayds
IV. Feedback Inhibition Regulation
F. Feedback inhibition regulates some biochemical pathways
1. Why? Unnecessary to synthesize a compound when plenty is present, doing so waste
energy and raw materials that could be used in other reactions
2. It is to the cell’s advantage to temporarily shut down biochemical pathways when their
products are not needed.
3. Regulation depends on elegant feedback mechanism.
a. End-product of pathway binds to allosteric site on enzyme that catalyzes the first
reaction in the pathway
b. Switches off or inhibits the beginning reaction from occurring.
c. This mode of regulation is called feedback inhibition. (See Fig. 6.15)