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Bioorganic chemistry, Study notes of Organic Chemistry

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.

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BIO-ORGANIC CHEMISTRY
UNIT-III, KINETICS AND MECHANISM
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[B]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
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BIO-ORGANIC CHEMISTRY

UNIT-III, KINETICS AND MECHANISM

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

[B]

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.

KINETIC AND THERMODYNAMIC REVERSIBILITY:

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)

E

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

  • Ea/RT

Solving the equation further:

lnk=ln(A)+ln(e

  • Ea/RT

lnk=ln(A)+(-E a

/RT)=ln(A)– (E a

/R)(1/T)

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‡).

ΔG , ΔS, ΔH, THERMODYNAMICS OF COUPLED REACTIONS:

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:

ΔG = ΔH - TΔS

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.