Reactions can be organized in 2 ways:
        1. What kinds of reactions occur.
        2. How reactions occur

Consider #1 first, kinds of reactions
1) Addition reaction         A + B ------> C

2) Elimination reaction         A -----> B + C
        (reverse of addition)

3) Substitution reaction         A + B ------> C + D

4) Rearrangement reaction         A -------> B

Consider how reactions occur - MECHANISM
        detailed step-by-step description of a reaction
        describes all bond-breaking and bond making processes.

Bond breaking and making can occur in two ways:
        1). Homolytic bond cleavage (radical)
                one electron to each fragment

        2). Heterolytic bond cleavage (polar)
                two electrons to one fragment

Let's consider a Radical Substitution Reaction

Chlorination of methane is a good example - what is the mechanism

There are three steps
1. Initiation

2. Propagation

repeat over and over-- overall process is a chain reaction

3. Termination

Polar Reactions:
        In polar reaction the electrons move in pairs. Why?
        We have said that certain bonds can be polar
        i.e.

Equilibria & Rates or Thermodynamics & Kinetics

In any reaction, an energy change must occur since the overall energy required to break all necessary bonds rarely equals energy given off by forming new bonds.

We call this Gibbs free energy DG°:         DG° = Free Energy Change

It may be expressed as
                DG° = -RT ln K_eq
                R = gas const. (1.987) cal / deg.kelvin · mole
                T = absolute temp (degrees kelvin)
                e = 2.718

Discuss signs
        K = 1 DG = 0
        K > 1 DG = negative -- exothermic
        K < 1 DG = positive -- endothermic

Table 4-1 gives conversion to products as a function of DG°

DG° dependent on 2 factors
        DG° = DH° - TDS°
                DS° = entropy term is a measure of disorder.
                DH° = standard heat of reaction (larger term)

DH° may be related to the energies required to break and form bonds
        DH° = DH° (bonds broken) - DH° (bonds formed)
                negative = exothermic
                positive = endothermic

If we know Bond Dissociation Energies
        we can calculate DH°

Table 4-2 (p. 142) gives some values of Bond Dissociation Energies

        Calculate DH° for reaction:
                CH₄ + Cl₂ --------> CH₃Cl + HCl

        DH°= 162 - 187 = -25 kcal/ mole (exothermic)

Can do same for
        first propagation step       DH° = +1 kcal/mole
        second propagation step  DH° = -26 kcal/mole
        Note initiation step is not included

Consider Energy Profile

Rate of Reaction

Rate = Collision frequency x Energy factor x Probability Factor

Collision frequency depends on:
        1) concentration
        2) size of particles
        3) temperature (how fast moving)

Probability Factor depends on:
        1) kind of reaction
        2) geometry of particles

Energy Factor (most important) depends on
        1) temperature
        2) energy of activation (characteristic of reaction)

Consider the energy distribution of molecules at some constant temperature

Boltzman, Gaussian (Bell curve) distribution
        When E_act = E1 a greater fraction of molecules possess needed
        energy than when E_act = E2

mathematics tells us that we can integrate a given area under the curve to define

        e^-E_act/RT = fraction of collisions which have energy > E_act

        where
                e = 2.718 (base of natural log)
                R = 1.986 (gas const.)
                T = absolute temp.
        then let
                P = probability factor
                Z = collision frequency

then Rate = P Z e^-E_act/RT (= P Z/e^E_actRT)

observe E_act has a large effect on rate

For example at 275°

Eact =	15 kcal/mole	1/1,000,000	collisions effective	(rate = 1)
	10	100/1,000,000		100
	5	10,000/1,000,000		10,000

What happens if we change temp.?
        T increases ===> rate increases
                (look at rate expression)

Reaction Profiles
        Consider changes in the structure of a system as a reaction proceeds.
                bond breaking ----- require energy
                bonds forming ----- release energy
                        overall change = DH° for reaction

                DH° negative=====> exothermic rxn
                DH° positive=====> endothermic rxn

Look at Reaction Coordinates - Potential Energy Diagram

Notice that there is an energy barrier (Activation Energy) even for exothermic processes.

i.e. CH₄ (natural gas) + O₂ --------> CO₂ + H₂O + energy
                        (Not spontaneous!)

The top of the energy barrier is the Transition State(‡)
Exactly equal probability of returning to reactants or proceeding on to products.

The transition state (T.S.) is a transient, high energy state; hence, we cannot deduce its exact structure. The structure of the TS is important because it tells us about how a reaction occurs.

Some reactions have more than one step
        between each step an intermediate is involved.
        represented by a minimum on the rxn cordinate
        stability of an intermediate is a function of the depth of this well.
i. e.

The lower the energy barrier (E_act) the larger the rate constant, k

The slow step (highest energy TS) is called the rate determining step.

Isotope Effects

Deuterium is the isotope of H with mass 2. The C-D bond is slightly stronger than the C-H bond and hence harder to break. Reactions which involve C-D cleavage go slower than C-H.

Evidence implies that a CH bond is broken in the rate determining step

A problem - polyhalogenation can lead to product mixtures

Can stop at mono-chloro stage using excess CH₄ (not practical)

Halogenation of Higher Alkanes

Most alkanes give mixture of products (book gives propane & gets slightly differant selectivities)

By taking into account: actual product ratios and number of each type of H, we can calculate relative reactivities of the 3 types (1°, 2°, 3°) of hydrogens towards chlorination.

Relative reactity toward chlorination:
                        R₃CH > R₂CH₂ > RCH₃
                          5.0         3.5         1.0

Reflects the relative stablity of the intermediate radicals:
                        R₃C· > R₂CH· > RCH₂·
                          3°           2°         1°

In terms of an energy diagram:

Consider the chlorination of propane vs. the bromination of propane:
        propane ----> 60% isopropyl chloride
        propane ----> 97% isopropyl bromide

Endothermic process: TS resembles more the products
Exothermic process: TS resembles more the reactants
Look at PE diagrams for these two processes:

Hammond Postulate Related species that are similar in energy are also similar in structure. The structure of a transition state resembles the structure of the closest stable species.

Reactive Intermediates short lived species that are never present in high conc because they react as quickly as they are formed.

Carbocations trivalent carbon bearing a positive charge

Electron donating groups can stabilize positive charge. Alkyl groups can be e-donating by hyperconjugation.

Resonance can also stabilize a carbocation

Free Radicals are also sp² hybridized and planar

since they are also electron deficient, they follow the same stability order as carbocations and can also be stabilized by resonance

Carbanions trivalent carbon with a negative charge. Now hybridization changes to sp³.

Because alkyl groups are electron donating, they destabilize a negative charge but resonance can be stabilizing

Carbenes uncharged divalent carbon

Carbenes are quite reactive and can tend to dimerize