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Mechanisms . One of the most practical aspects of organic chemistry is the study and application of chemical reactions. Due to the large number of reactants that can be used, it is virtually impossible to memorize all possible reactions.

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Mechanisms

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Mechanisms

  • One of the most practical aspects of organic chemistry is the study and application of chemical reactions.

  • Due to the large number of reactants that can be used, it is virtually impossible to memorize all possible reactions.

  • Organic reactions are often organized into groups based on their mechanisms and the intermediates that are involved.


Mechanisms

  • Mechanism:

    • a step-by-step pathway from reactants to products that shows which bonds break, which bonds form and the order in which they happen

    • includes structures of all reactants, intermediates and products and curved arrows showing the movement of electrons


Mechanisms

  • Your success in this class depends in large part on learning the mechanisms of key reactions and applying these mechanisms to predict the products formed from starting materials you have not used before.


Mechanisms

  • Halogenation of alkanes:

    alkane + halogen alkyl halide(s) + HX

    CH4 (g) + Cl2 (g) CH3Cl+ CH2Cl2 +

    CHCl3 + CCl4 + HCl

D or hu

D or hu


Mechanisms

  • The halogenation of alkanes is a substitution reaction that occurs via a chain reactionmechanism.

  • Substitution reaction:

    • a reaction in which one atom substitutes for or replaces another atom

      • In the chlorination of methane, a chlorine atom replaces a hydrogen atom.


Mechanisms

  • Three types of steps occur in all chain reactions:

    • Initiation:

      • generates a reactive intermediate

      • Reactive intermediate:

        • a short-lived species that reacts as quickly as it is formed

          • never present in high concentration


Mechanisms

  • Three types of steps occur in all chain reactions:

    • propagation

      • reactive intermediate reacts with a stable molecule to form a new reactive intermediate and a new stable molecule

        • continues until reactants are exhausted or reactive intermediate is destroyed

    • termination

      • side reaction that destroy the reactive intermediate

        • slows or stops reaction


Mechanisms

  • Chlorination of Methane (Mechanism)

    • Initiation:

      • Chlorine absorbs hu generating twofree radicals.

Use half arrows to show movement of one electron.


Mechanisms

  • Free radical

    • areactive intermediate with one or more unpaired electrons

      • also called a radical

      • electron deficient (doesn’t have octet)


Mechanisms

  • Chlorination of Methane (Mechanism) cont.

    • Propagation:

      • First propagation step:

        • Chlorine radical collides with a methane molecule and abstracts a hydrogen atom


Mechanisms

  • Chlorination of Methane (Mechanism) cont.

    • Second Propagation step:

      • methyl radical reacts with chlorine molecule, generating product and another reactive species

      • The new chlorine radical continues the chain by abstracting another hydrogen atom from methane, etc.


Mechanisms

  • Chlorination of Methane (Mechanism) con’t.

    • Termination:

      • Any reaction that produces fewer reactive intermediates than it uses will slow or stop the reaction:


Mechanisms

  • More possible termination reactions:


Thermodynamics

  • Information about chemical reactions is obtained using thermodynamics and kinetics.

  • Thermodynamics: used to study the stability of reactants and products

    • predicts which compounds are favored by the equilibrium


Thermodynamics

  • For an equilibrium reaction:

    a A + b B c C + d D

    DGo = -RTlnKc

  • Spontaneous reaction (favors products):

    • Kc > 1

    • DG = neg

  • Nonspontaneous reaction (favors reactants):

    • Kc < 1

    • DG = pos


Thermodynamics

  • Two thermodynamic quantities contribute to DG:

    DG = DH - TDS

  • DH = enthalpy change (amount of heat gained or lost)

    • DH = positive (endothermic: heat gained)

    • DH = negative (exothermic:heat lost)

  • DS = entropy change

    • change in the randomness or disorder


Thermodynamics

  • For many organic reactions, TDS is small relative to DH

    • DG ~ DH

  • Therefore, most exothermic organic reactions tend to favor the formation of products.

  • The DHrxn can be estimated using the bond dissociation energies of the bonds broken and formed during the reaction.


Thermodynamics

  • Bond dissociation energy:

    • the amount of energy required to break a bondhomolytically

      • equally

      • each atom in the bond being broken gets one electron

        • forms free radicals

  • As BDE increases, more energy is needed to break the bond:

    • stronger bond


Thermodynamics

Example: Which of the following bonds is the strongest? The weakest?

F - F, Cl - Cl, CH3 - F, CH3 - Cl, H - F, or H - Cl

F - F

Cl - Cl

CH3 - F

CH3 - Cl

H - F

H - Cl

38 kcal/mol

58

109

84

136

103


Thermodynamics

Example: Rank the following C-H bonds in order from the easiest to the hardest to break homolytically.

1o H

98 kcal/mol

Methyl H

104 kcal/mol

3o H

91 kcal/mol

2o H

95 kcal/mol


Thermodynamics

  • As BDE increases, it is harder to break the bond:

    • Ease of homolytic cleavage:

      3o > 2o > 1o > methyl

  • The stability of methyl, 1o, 2o, and 3o free radicals follows the same trend:

    3o > 2o > 1o > methyl

    • i.e. 3o free radicals are the most stable and methyl radicals are the least stable

(easiest)

(hardest)


Kinetics

  • Many reactions that have favorable energy changes (DG = neg or DH = neg) occur so slowly that the reaction is imperceptible.

    • Very slow reaction rate

  • For the general reaction:

    a A + b B c C + d D

    Rate = k [A]m[B]n

    where k = rate constant

    m = reaction order with respect to A

    n = reaction order with respect to B


Kinetics

  • The reaction rate depends on:

    • collision frequency

    • a probability or orientation factor

    • activation energy (Ea)

  • The reaction rate increases as the number of collisions between reacting species increase.

    • Concentration

    • temperature


Cl .

Cl .

Cl .

Br

Br

Br

H

H

H

Kinetics

  • Collisions must occur in a particular orientation for reactions to occur.

  • For the reaction:

    Cl.+ H - Br H - Cl + Br.

No HCl formed

No HCl formed

HCl can form


Kinetics

  • Collisions must occur with a specific minimum amount of energy in order for a reaction to take place.

    • Activation energy (Ea)

      • the minimum energy the reactants must have for a reaction to occur

      • the energy difference between the reactants and thetransition state


Kinetics

  • Transition state:

    • a particular arrangement of atoms of the reacting species in which bonds are partially broken and partially formed

    • the state of highest energy between reactants and products

    • a relative maximum on the reaction-energy diagram


Reaction Energy Diagram

  • Reaction energy diagram:

    • a plot of potential energy changes that occur as reactants are converted to products


Hammond Postulate

  • What does the transition state look like?

    • The appearance of the transition state depends on whether the reaction is endothermic or exothermic.

      • governed by the Hammond Postulate

  • Hammond Postulate:

    • Related species that are closer in energy are also closer in structure.

      • The structure of the transition state resembles the structure of the closest stable species.


Hammond Postulate

  • For an endothermic reaction, the transition state more closely resembles the products.

  • For an exothermic reaction, the transition state more closely resembles the reactants.

exothermic

endothermic


Rate Determining Step

  • Chlorination of methane has two propagation steps.

    • The first propagation step controls the rate of the overall reaction and is called the rate-determining step.

  • Rate-determining step (rate-limiting step):

    • the slowest step in a multi-step process

    • the step with the highest energy transition state


Rate Determining Step

intermediate


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