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Chapter 8.1:
Metabolism, Energy, and the Laws of
Thermodynamics
By Richie Kovac
What is Metabolism?
• Metabolism – the sum of an organism’s chemical reactions. It
is an emergent property resulting from interactions between
molecules within the environment of the cell.
• Metabolic pathway – a series of chemical reactions. It begins
with a molecule, is changed through a series of reactions
(which are catalyzed by enzymes), and is made into a product.
Metabolic pathways in an organism intersect like roads on a
map intersect.
• Metabolism is like a large map of all the metabolic pathways in
an organism.
Metabolism:
sum of all the
metabolic
pathways in an
organism
One street is like a
metabolic pathway. The
pathways intersect.
Example of a metabolic pathway.
Molecule A is changed into Molecule F
What Does the Metabolism Do?
• Manages material and energy resources in a
cell
• Mechanisms control enzymes that are part of
metabolic pathways to either:
– a) catalyze a metabolic pathway to produce a
certain product if it is in deficiency or
– b) to prevent enzymes from catalyzing a reaction,
preventing the production of a product (if the
product is in excess).
Too much of a
product inhibits the
enzyme involved in
making it. As a result,
accumulation of the
product stops
Types of Metabolic Pathways
• Catabolic pathways – metabolic pathways that
break down complex molecules into simpler
Amino
Acids
compounds,
which releases energy
form a
polypeptide
–
– Example
an example of
an anabolic
pathway.
– cellular respiration
Cellular
Respirationan example
of a
catabolic
pathway.
• Anabolic pathways – pathways that consume
energy to build complicated molecules from
simpler ones
– Example – proteins are built from amino acids.
Breakdown
Proteins to Amino Acids, Starch to Glucose
Synthesis
Amino Acids to Proteins, Glucose to Starch
Chapter 5
Chapter 5
Chapter 5
Quick Questions
•
•
•
•
What is Metabolism?
What does metabolism do?
What are 2 types of metabolic pathways?
Which metabolic pathway consumes energy to
build complicated molecules from simpler
ones?
• Using a map as a symbol, what would a road
on a map symbolize?
Energy
• Energy – the capacity to cause change/the
ability to rearrange a collection of matter.
• Bioenergetics – the study of how energy flows
through living systems.
• Can be used to perform work – pushing
matter against opposing forces.
Types of Energy
• Kinetic energy – the relative motion of objects
– The motion of large or small objects
• Heat (thermal energy) – kinetic energy associated with the random
movement of atoms or molecules (in a body of matter, the amount of heat
is a measure of the matter’s total kinetic energy due to the motion of its
molecules).
– Has to do with the motion of particles like atoms and molecules
• Light
• Potential energy – energy that is not kinetic. It is energy that matter
possesses because of its location or structure.
– Example – water held in a dam above sea level had potential energy. A diver on a diving
board has potential energy.
• Chemical energy – a form of potential energy, this is energy available in
molecules for release in a chemical reaction
Potential Energy
Quick Questions 2
•
•
•
•
•
What is energy?
What is work?
Chemical energy is a type of _____ energy.
How are thermal and kinetic energy similar?
A cheetah running displays which type of
energy in use?
Thermodynamics
• Thermodynamics – the study of energy
transformations that occur in a collection of
matter.
• “System” – name given to a piece of matter
under study. Everything outside the system is
known as its “surroundings”
• Isolated system – cannot exchange matter or
energy with its surroundings.
• Open system – can exchange matter or energy
with its surroundings.
1st Law of Thermodynamics
• Says energy of the universe is constant. Energy
can be transferred and transformed, but it
cannot be destroyed or created.
Matter is conserved – it’s
not created or destroyed
2nd Law of Thermodynamics
• Says every energy transfer/transformation
results in some energy becoming unusable
energy. The creation of unusable energy
increases the entropy of the universe.
• Entropy – a measure of unusable energy
within a closed or isolated system. It is also
defined as a measure of disorder or
randomness. Increased unusable energy in a
system, such as the universe, increases chaos
in the universe.
• To understand the concept of entropy, think of
an energy transfer as a class of 20 kids walking
from one end of a hallway to another. The kids
represent useable energy and their teachers
represent systems that use the energy. By the
time the kids reach the other end of the
hallway, 2 of the 20 kids have turned crazy and
chaotic and can no longer be taught (“used”)
by the teachers. As the kids move back and
forth, eventually all of them will go crazy.
2nd Law of Thermodynamics
Continued
• This law also states that for a process to occur spontaneously,
it must increase the entropy of the universe.
• Spontaneous process – one that occurs without the input of
energy. Example, a boulder rolling down a hill. The boulder
rolls down the hill, releases heat and so increases entropy,
and the boulder becomes more stable.
• This law permits that the entropy of a particular system may
decrease as long as the total entropy of the universe
increases.
Spontaneous
processes, like an
uncared office falling
apart after a while,
increases the entropy
of the universe.
Final Quick Questions
• Can an isolated system exchange energy with
its surroundings?
• What does the 1st law of thermodynamics
say?
• What is entropy?
• How is entropy related to the 2nd law of
thermodynamics?
• You’re talking to your friend. He tells you “I
performed a chemical reaction today. I ended
up with less energy than I started with.” Why
is your friend wrong?
• Metabolism Animation
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Free energy change- portion of a system’s energy that can
perform work when temperature and pressure are uniform
throughout the system
The free energy change of a reaction tells us whether or not
the reaction occurs spontaneously
Symbolized by the letter G
• For the professor J. Willard Gibbs, who defined a very useful function
called the Gibbs free energy of a system (without considering its
surroundings ) in 1878


T
his formula is used to determine the free energy change that occurs
when a system changes
• Example: during a chemical reaction

∆G is the free energy of the system

Formula uses only properties of the system (the reaction) itself.

∆H symbolizes the change in system’s enthalpy (in biological systems,
this is equivalent to total energy)

∆S is the change in the system’s entropy

T is the absolute temperature in Kelvin (K) units
∆G= G
final state
–G
initial state
 Gfinal state = free energy of the final state of system
 Ginitial state = free energy of the initial state of the system

Spontaneous reactions have a negative ∆G
• ∆G<0
• Spontaneous reactions must either give up enthalpy (H must decrease), give
up order (TS must increase), or both
• This means that every spontaneous process decreases the system’s free
energy


Processes that have a positive or zero ∆G are never spontaneous
Spontaneous does not mean that a reaction will be instantaneous
or rapid
More free energy (higher G)
Less stable
Greater work capacity
In a spontaneous change:
The free energy of the system decreases
(∆G <0)
The system becomes more stable
The released energy can be harnessed to
do work
Less free energy (lower G)
More stable
Less work capacity





Chemical reactions which proceed downhill, are said to be
spontaneous.
Spontaneous process may also require energy to overcome an
activation barrier.
• This means that some processes may be spontaneous, but do
not occur at a noticeable rate.
Spontaneous reactions occur without outside intervention.
They may occur quickly
Ex’s:
• Combustion of hydrogen
• Graphite turning to diamond
• Expansion of a gas to fill the volume available to it.
• Cooling of a hot object to the temperature of the surroundings

According to the second law of thermodynamics, when a
spontaneous process occurs, there must be an increase in total
entropy.
A
reaction which cannot occur without
the input of work from an external source.
ΔG > 0 for nonspontaneous reactions at T
and P.



Free energy is a measure of a system’s instability
Unstable systems (higher G) tend to change in such a way that they
become stable (lower G)
As a reaction reaches equilibrium, the free energy of the mixture
decreases
• Conversely, the free energy of the mixture increases when a reaction is somehow
pushed away from equilibrium

Any change from equilibrium state will have a positive ∆G and will
not be spontaneous
• For this reason systems never spontaneously move away from equilibrium
• Because a system at equilibrium can’t spontaneously change, it can do no work
• A process is spontaneous and can perform work only when moving toward
equilibrium

For a system at equilibrium, G is at its lowest possible value in the
system

Exergonic reaction- proceeds with a
net release of free energy; energy
released
• Negative ∆G
• Occur spontaneously
• Example: cellular respiration
 C6H12O6 + 6O2 → 6CO2 + 6H2O
 ∆G= -686 kcal/mol (-2870 kJ/mol)

Endergonic reaction- absorbs free
energy
from its surroundings; energy required
• Stores free energy in molecules- ∆G is
positive
• Non-spontaneous
• Example: photosynthesis
 6CO2 + 6H2O + light energy → C6H12O6 +
6O2
 ∆G= 686 kcal/mol
•Exergonic- Release energy
•Usually occur spontaneously since they do not require energy to
occur.
•Some exergonic reactions do not occur spontaneously. Instead, they
require a small input of energy from some outside source to cause the
reaction. This outside energy is called the activation energy of the
reaction.
•Endergonic - Asorb energy to form bonds
•Do not occur spontaneously.
•Occur only if energy is available to be used in the reaction.
•Endergonic reactions also have an activation energy which is
considerably higher than the activation energy for most exergonic
reactions.
•Many reactions in cells are endergonic, so cells require a method of
storing energy until it is needed in a chemical reaction.
Is the reaction endothermic or exothermic?
This is an exothermic reaction (heat is a product), so heat
is released to the surroundings
Ellie Quinby




Chemical work: The pushing of endergonic reactions that
would not occur spontaneously
◦ Ex: synthesis of polymers and monomers
Transport work: The pumping of substances across
membranes against the direction of spontaneous
movement
◦ Ex: proteins through membranes
Mechanical work: Movement of the cell itself
◦ Ex: Beating of cilia, contraction of muscle cells
Energy coupling: use of an exergonic process to drive
and endergonic one (manages cell’s energy use)


ATP (adenosine triphosphate)
◦ Contains ribose with nitrogenous base adenine
and three phosphate groups bonded to it
◦ Used to make RNA
◦ Breaking it down release more energy then most
molecules
Hydrolysis◦ Adding water to ATP breaks the molecule into
adenosine diphosphate, releasing energy



Hydrolysis gets rid of one phosphate and it
gets attached to a different molecule
Phosphorylation- the molecule that gains
the phosphate group is called
phosphorylated
Regenerated through phosphorylation of
ADP

Animation



Enzymes can denature in heat or extreme pH.
◦ Most enzymes optimal temperature is 98.6, body
temperature
Many use coenzymes to catalyze reactions
Some chemicals inhibit the ability of the enzyme to
work
◦ Many are in cells naturally and control metabolism


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

Some reactions require to much energy to begin they are
helped by catalysts
Activation energy (Free energy of activation): amount of
energy needed for a reaction to occur represented by EA
Enzymes that act like catalysts decrease the activation
energy of a reaction
Activation energy barrier- the energy a reaction must
absorb to begin- determines rate of reaction
Transition State- when a reactions has enough energy to
begin, bonds can be broken



Heat cannot be used to lower activation energy or it
would denature cells’ proteins
◦ Therefore a catalyst must be used
Enzymes- catalyze specific reactions in the cells,
allowing them to occur quicker
Catalyze in the direction of lower change in G( free
energy change)



Substrate- the reactant and enzyme acts on
◦ The enzyme fits with a specific substrate, and
that it what it catalyzes
◦ Forms and enzyme/ substrate complex when
together
Active site- the area of the enzyme that actually
bonds to the substrate
Induced fit- when the substrate enter the active
site, the enzyme envelops them, allowing a
reaction to occur
• Enzyme Catalysis
Animation

Enzyme Animation
ENZYMES AND METABOLISM
Factors affecting Enzyme Activity
Prerna Balasundaram
OPTIMAL CONDITIONS
•
•
•
•
•
•
Enzymes work better under some
conditions than others; this is
called optimal conditions.
High temperature, increases rate
of enzyme reaction allowing each
enzyme to have the best activity,
however, without it denaturing
because of extreme high
temperatures.
Each enzyme has its own optimal
temperature.
EXAMPLES
- Human enzymes’ optimal
temperature – @35 degrees Celsius
(close to human temperature).
Just like temperature, each enzyme
has its optimal pH.
EXAMPLES
- Most enzyme’s pH 6-8.
- The human stomach enzyme,
pepsin, works best at pH of 2.
COFACTORS AND COENZYMES

Cofactors – non protein
helpers for catalytic activity
that may either be bound
tightly to enzyme as
permanent residents, or may
bind loosely along with
substrate; usually are
inorganic.

EXAMPLES

- Zinc, iron, copper
Coenzymes – cofactors that
are organic.

EXAMPLES
- Vitamins, raw materials
COMPETITIVE & NONCOMPETITIVE
INHIBITORS



Competitive inhibitors –
inhibitors that resemble
normal substrates and reduce
productivity of enzymes by
blocking real substrates from
entering active site
Noncompetitive inhibitors
– do not directly compete with
substrate to bind to enzyme at
active site; rather they
impede enzyme reaction by
binding with another part of
enzyme
EXAMPLE
- DDT is a pesticide that
inhibits key enzymes in
nervous system
ALLOSTERIC REGULATION



Chemical chaos would result
if all of a cell’s metabolic
pathways were operating at
the same time!
Allosteric regulation – any
case in which a protein’s
function at one site is affected
by the binding of a inhibiting
regulatory molecule to a
separate, regulatory site
(sometimes called an
allosteric site).
Made up by two or more
subunits, each composed of a
polypeptide chain and having
its own active site.
COOPERATIVITY & FEEDBACK INHIBITION


•
•
Feedback inhibition – a
metabolic pathway is switched off
by the inhibitory binding of its
end product to an enzyme that
acts early in the pathway.
EXAMPLE
- ATP binds to several enzymes
allosterically inhibiting their
activity. ADP acts as an
activator so if ATP production
lags behind ADP activates
enzymes and speeds up
production, vice versa if ATP
production is too high.
Cooperativity – one substrate
molecule primes an enzyme to
accept additional substrate
molecules amplifying response of
enzymes to substrates
EXAMPLE
- Oxygen transport of protein
hemoglobin
REVIEW QUESTIONS
What is the optimal temperature and pH for
most human enzymes?
 Give an example of a cofactor and a coenzyme.
 Compare competitive inhibitors and
noncompetitive inhibitors.
 What is allosteric regulation?
 Describe how ATP uses feedback inhibition.


Enzyme Inhibition Animation

Enzyme Specificity Animation