






Study with the several resources on Docsity
Earn points by helping other students or get them with a premium plan
Prepare for your exams
Study with the several resources on Docsity
Earn points to download
Earn points by helping other students or get them with a premium plan
Community
Ask the community for help and clear up your study doubts
Discover the best universities in your country according to Docsity users
Free resources
Download our free guides on studying techniques, anxiety management strategies, and thesis advice from Docsity tutors
Bioorganic Chemistry is a specialized subject that integrates the principles of organic chemistry with the complexities of biological systems. It aims to give biotechnology students a strong foundation in how organic molecules behave in biological contextis course typically covers: ✅ Structure & Reactivity: Understanding organic molecules important in biology, including carbohydrates, amino acids, nucleic acids, and lipids. ✅ Mechanisms: How organic reaction mechanisms explain enzyme-catalyzed reactions, biosynthetic pathways, and drug interactions. ✅ Stereochemistry: Concepts like chirality and optical activity, critical for biological activity of molecules. ✅ Reaction Types: Such as substitution (SN1, SN2), elimination, addition, oxidation-reduction in biological molecules. ✅ Catalysis: Chemical catalysis vs enzyme catalysis — coenzymes, metal ions, acid-base and covalent catalysis.
Typology: Study notes
1 / 11
This page cannot be seen from the preview
Don't miss anything!
RATE LAW AND MECHANISM : A rate law describes how the rate of a reaction changes with
the concentration of reactants, while a reaction mechanism explains the step-by-step process of how
reactants transform into products.
Rate Law: A rate law is an equation that expresses the relationship between the rate of a reaction
and the concentrations of reactants. It is determined experimentally, not derived from the balanced
chemical equation of the overall reaction.
For a simple reaction, the rate law might be expressed as:
Rate = k[A]
m
n
Where k is the rate constant, [A] and [B] are the concentrations of reactants, and m and n are the
reaction orders with respect to each reactant.
Reaction Mechanism: A reaction mechanism is the sequence of elementary reactions (individual
steps) that lead to the overall transformation of reactants into products. These elementary steps
are often described as single molecule collisions or a series of such collisions. The slowest step in
the mechanism is called the rate-determining step, which dictates the overall reaction rate and
influences the rate law.
Relationship between Rate Law and Mechanism: A proposed mechanism should be consistent
with the experimentally determined rate law. If a mechanism predicts a certain rate law and that
rate law is experimentally observed, it provides evidence for the validity of the
mechanism. However, a mechanism can only be inferred from the rate law, not the other way
around. Multiple mechanisms could potentially be consistent with the same rate law, so
experimental evidence beyond kinetics is often needed to confirm a mechanism.
Dr. N. Devika
TRANSITION STATES AND INTERMEDIATES : In chemical reactions, both transition states
and intermediates represent points along the reaction pathway, but they differ significantly in their
stability and lifetime. Transition states are the fleeting, unstable configurations that exist at the peak
of the energy barrier between reactants and products, whereas intermediates are more stable,
discrete species that can exist for a measurable period.
Transition States: The highest energy point in a reaction pathway, representing the moment where
reactants are transitioning to products.
Lifetime: Extremely short, lasting only a few seconds.
Stability: Highly unstable.
Energy: Highest energy point in the reaction.
Role: The point of maximum energy that must be overcome for a reaction to proceed.
Intermediates: Stable, discrete species that are formed during a reaction and can exist for a longer
period before being consumed.
Lifetime: Can range from nanoseconds to even days, depending on their stability.
Stability: More stable than transition states.
Energy: Exist at a local energy minimum on the reaction pathway.
Examples: Carbocations, radicals, and biradicals.
Differences:
Feature Transition State Intermediate
Stability Highly unstable Relatively stable
Lifetime Femto seconds Nanoseconds to days
Energy Highest energy point Local minimum on energy profile
Isolatable No Potentially yes, under certain conditions
MICROSCOPIC REVERSIBILITY : It is a principle stating that in a system at equilibrium, any
molecular process and its reverse occur at the same rate on average. This principle, also known as
detailed balance, is fundamental in understanding rate and equilibrium phenomena, particularly in
chemical kinetics.
Key aspects of microscopic reversibility:
Dynamic Description of Equilibrium: It provides a dynamic view of equilibrium, where individual
molecular processes are constantly occurring in both forward and reverse directions.
Rate Equivalence: The principle states that for every forward reaction, there is a corresponding
reverse reaction that proceeds at the same rate under equilibrium conditions.
Equilibrium Constant and Rate Constants: The principle helps relate the ratio of forward and
reverse rate constants to the equilibrium constant.
Mechanism of Reaction: It implies that the forward and reverse reaction pathways follow the same
sequence of molecular structures at the microscopic level, meaning the same transition state is
involved.
Applications:
Chemical Kinetics: The principle is used to understand reaction rates, equilibrium constants, and
reaction mechanisms.
Thermodynamics: It helps in understanding the dynamic nature of equilibrium.
Molecular Machines: It is a key principle in understanding the operation of molecular machines,
like motors in biological systems.
Optics: The principle of reversibility in optics, where light paths are reversible, is a related
concept.
In chemical reactions, kinetic control refers to the product that forms fastest, while
thermodynamic control refers to the most stable product, under reversible conditions. Kinetic
control is dominant at lower temperatures where reactions are generally irreversible, and
thermodynamic control is dominant at higher temperatures where reactions can be reversible.
Kinetic Control: The product that forms at the fastest rate dominates the reaction. Typically occurs
at lower temperatures where the reaction is irreversible. The reaction is not reversible, (i.e) once the
kinetic product is formed, it doesn't readily revert back to the reactants or another product.
Thermodynamic Control: The most stable product dominates the reaction under reversible
conditions. Typically occurs at higher temperatures where the reaction is reversible, allowing for
equilibrium between products. Reactions are reversible, meaning the reaction can proceed in both
directions, and the equilibrium favors the more stable product.
Reversibility and Thermodynamic Control: For a reaction to be under thermodynamic control, it
must be reversible, meaning the reaction can proceed in both directions. The relative stability of
the products determines the equilibrium constant, with the more stable product being favored.
Under thermodynamic control, the reaction will proceed until it reaches the lowest energy state,
which corresponds to the most stable product.
PRIMARY AND SECONDARY ISOTOPES : Primary isotopes are those involved in the bond-
breaking or bond-forming step of a reaction, while secondary isotopes are those that are not directly
involved but still influence the reaction rate.
Primary Kinetic Isotope Effect: When a bond to an isotopically labeled atom is broken or formed
during the rate-determining step of a reaction, the effect is considered a primary kinetic isotope
effect.
Mechanism: These effects arise from changes in zero-point vibrational energies between reactants
and transition states.
Magnitude: Primary effects are generally larger than secondary effects.
Examples: The isotope effect observed when a hydrogen atom is replaced with deuterium in a
reaction involving the breaking of a C-H bond is a classic example of a primary isotope effect.
Secondary Kinetic Isotope Effect: When the isotopically labeled atom is not directly involved in
the bond-breaking or bond-forming step of the reaction, but its presence still affects the reaction
rate, it's a secondary kinetic isotope effect.
Mechanism: Secondary effects can arise from changes in hybridization, hyper conjugation, or
other factors that affect the electronic structure of the molecule.
Types: Secondary effects can be further categorized into alpha, beta, and gamma effects,
depending on their location relative to the bond being broken.
Magnitude: Secondary effects are generally smaller than primary effects.
Example: A secondary isotope effect could be observed when a hydrogen atom is replaced with
deuterium on a carbon atom adjacent to the bond being broken in a reaction.
In essence, primary isotope effects are directly related to the bond being broken or formed in
the rate-determining step, while secondary effects are influenced by the isotopic substitution at a
location further from the reaction center.
THE ARRHENIUS EQUATION: The Arrhenius equation is an expression that provides a
relationship between the rate constant (of a chemical reaction), the absolute temperature, and the A
factor (also known as the pre- exponential factor). The expression of the Arrhenius equation is:
Where,
K = rate constant of the reaction,
A = pre-exponential factor which, in terms of the collision theory, is the frequency of
correctly oriented collisions between the reacting species
e = base of the natural logarithm (Euler’s number)
a =
activation energy of the chemical reaction (in terms of energy per mole)
R = universal gas constant
T = absolute temperature associated with the reaction (in Kelvin)
If the activation energy is expressed in terms of energy per reactant molecule, the universal gas
constant must be replaced with the Boltzmann constant (k B
) in the Arrhenius equation.
Arrhenius Plot: When logarithms are taken on both sides of the equation, the Arrhenius equation
can be written as follows:
ln k = ln(Ae
Solving the equation further:
lnk=ln(A)+ln(e
lnk=ln(A)+(-E a
/RT)=ln(A)– (E a
Since ln(A) is a constant, the equation corresponds to that of a straight line (y = mx + c)
whose
slope(m) is - E a
/R. When the logarithm of the rate constant (lnK) is plotted on the Y-axis
and the inverse of the absolute temperature (1/T) is plotted on the X-axis, the resulting graph is
called an Arrhenius plot.
Fig.1 Arrhenius Plot
The Arrhenius plot for the decomposition of nitrogen dioxide is illustrated above.
Temperature Dependence: The equation explicitly includes the temperature, allowing for the
calculation of reaction rates at different temperatures.
Activation Energy and Entropy: It incorporates both the activation energy (enthalpy) and the
activation entropy, providing a more complete picture of the reaction mechanism.
Rate Constant Calculation: The Eyring equation allows for the calculation of the rate constant
(k) based on the transition state properties.
Applications: The Eyring equation finds applications in various fields, including chemical
kinetics, reaction dynamics, and even transport processes like viscosity and diffusion.
The Eyring equation can be expressed as:
Where,
k : rate constant
kB: Boltzmann constant Δ
T: absolute temperature
h: Planck's constant
K_‡: equilibrium constant for the transition state
ΔG‡: Gibbs free energy of activation
R: ideal gas constant
The Eyring equation states that the rate constant (k) is proportional to the temperature and
the equilibrium constant for the transition state (K_‡), and is inversely proportional to the free
energy of activation (ΔG‡).
For example, if you know the Gibbs free energy of activation (ΔG‡) for a reaction, you
can use the Eyring equation to calculate the rate constant (k) at a given temperature
(T). Conversely, if you measure the rate constant (k) at a specific temperature (T), you can use
the Eyring equation to estimate the Gibbs free energy of activation (ΔG‡).
In thermodynamics, ΔG (change in Gibbs free energy), ΔS (change in entropy), and ΔH
(change in enthalpy) are interconnected and determine the spontaneity of a reaction. Coupled
reactions utilize the energy released by a spontaneous reaction to drive a non-spontaneous
reaction.
k = (kBT/h) * K_‡ * exp(-ΔG‡/RT)
Interrelationship of ΔG, ΔS, and ΔH:
The Gibbs free energy change (ΔG) is calculated from the enthalpy change (ΔH) and the
entropy change (ΔS) at a given temperature (T).
Spontaneity:
If ΔG < 0, the reaction is spontaneous(proceeds in the forward direction).
If ΔG > 0, the reaction is non-spontaneous(proceeds in the reverse direction).
If ΔG = 0, the reaction is at equilibrium.
Factors influencing spontaneity:
ΔH < 0 and ΔS > 0: The reaction is always spontaneous (ΔG < 0) at all temperatures.
ΔH > 0 and ΔS > 0: The reaction can be spontaneous at high temperatures (ΔG < 0)
where the TΔS term dominates.
ΔH < 0 and ΔS < 0: The reaction can be spontaneous at low temperatures (ΔG < 0)
where ΔH dominates.
ΔH > 0 and ΔS < 0: The reaction is always non-spontaneous (ΔG > 0) at all
temperatures.
Coupled Reactions: Coupled reactions involve linking a thermodynamically favorable reaction
(exergonic, ΔG < 0) to a thermodynamically unfavorable reaction (endergonic, ΔG > 0) to
drive the latter.
Mechanism: The favorable reaction releases energy, which is then used to overcome the
energy barrier of the unfavorable reaction.
Shared intermediates: Coupled reactions often involve shared intermediates, where the
products of one reaction act as reactants for the other.
Examples:
ATP hydrolysis: The hydrolysis of ATP (a favorable reaction) provides energy for other
processes.
Muscle contraction: The coupling of ATP hydrolysis to muscle protein interactions
drives muscle contraction.
Active transport: Coupled reactions can drive the movement of molecules across cell
membranes against their concentration gradient.