Ch 06- Chemical Reactivity.pdf

(1)

Chapter 6

(2)

6.1 Enthalpy

• Enthalpy (ΔH or q) is the heat energy exchange between the reaction and its surroundings at constant pressure

• Breaking a bond requires the system to absorb energy. WHY?

(3)

6.1 Enthalpy

• Bonds can break homolytically or heterolytically

• Bond dissociation energy (BDE) or ΔH for bond breaking generally represents the energy associated with

(4)

(5)

6.1 BDEs

• Explain how heat energy is exchanged between the reaction (system) and the solution (surroundings) for each scenario below

1. H• and F• free radicals come together to form bonds

2. A C–Br bond is broken

3. A strong bond is broken and a weak bond is formed

(6)

6.1 Enthalpy Δ

H

• Match the reaction coordinate diagrams below with the statements below

– The reaction causes the surrounding temp. to DECREASE

(7)

6.2 Entropy Δ

S

• Although most reactions are EXOthermic, there are many ENDOthermic reactions that occur

• Enthalphy and entropy must BOTH be considered when predicting whether a reaction will occur

• ENTROPY (ΔS) can be though of as molecular disorder, randomness, or freedom

(8)

6.2 Entropy Δ

S

• If the energy of molecules can be distributed in a higher number of vibrational, rotational, and translational

states, the sample will have a greater entropy.

• Molecules exhibit vibrational, rotational, and translational motion. Explain HOW

(9)

6.2 Entropy Δ

S

• Consider the ENTROPY change for the following process

(10)

6.2 Entropy Δ

S

• The total entropy change will determine whether a process is spontaneous (favors the forward direction)

• If ΔS_tot is positive, the process is spontaneous. What if ΔS_tot is negative?

• When the volume of a gas expands to fill a container, what should ΔS_surr be?

(11)

6.2 Entropy Δ

S

• For each of the reactions below, predict the sign for ΔS_sys

• Consider how a change in a molecule’s structure affects the number of possible translational, rotational, and/or vibrational distributions for the molecules?

(12)

6.2 Entropy Δ

S

• How would the following conditions affect spontaneity (the degree to which the reaction is product favored)?

1. The reactions are exothermic

(13)

6.3 Gibbs Free Energy Δ

G

• We know that the spontaneity of a process depends only on ΔS_tot

• ΔS_syscan be measured or estimated

• ΔS_surrdepends on ΔH_sys

(14)

G

• Multiply both sides by Temperature

» or

(15)

6.3 Gibbs Free Energy Δ

G

• Consider the example reaction

• Predict the sign (+ or -) for ΔS_sys

• In the reaction, two pi bonds are converted into two sigma bonds. Predict the sign (+ or -) for ΔH_sys

• Predict the sign (+ or -) for ΔS_surr

• Predict the sign (+ or -) for ΔG

(16)

G

• If a process at a given temperature is calculated to have a (-) ΔG, the process is exergonic

•• It will be spontaneous and favor the

products

• Note that G is plotted rather than H

(17)

6.3 Gibbs Free Energy Δ

G

• If a process at a given temperature is calculated to have a (+) ΔG, the process is endergonic

• It will be

NONspontaneous and favors the reactants

(18)

6.4 Equilibria

• Consider an exergonic process with a (-) ΔG. Will every molecule of A and B be converted into products?

– No, an equilibrium will eventually be reached

– A spontaneous process will simply favor the

products meaning there will be more products than reactants

– The greater the

(19)

6.4 Equilibria

• Why doesn’t an exergonic process react 100% to give products? Why will some reactants remain?

– The diagram shows one unit of A react with one unit of B

– In reality, moles of reactants are present

– How will

(20)

6.4 Equilibria

• In any reaction, collisions are necessary

• As [A] and [B] decrease collisions between A and B will occur less often

• As [C] and [D] increase, collisions between C and D will occur more often

– The more often C and D collide, the more often collisions will occur with enough free energy for the reverse reaction to take place

(21)

6.4 Equilibria

• Equilibrium is also the state with the lowest free energy overall

• Why does the equilibrium mixture have the least G?

(22)

6.4 Equilibria

• An equilibrium constant (K_eq) is used to show the degree to which a reaction is product or reactant favored

(23)

6.4 Equilibria

(24)

6.5 Kinetics

• Recall that a (-) sign for ΔG tells us a process is product favored (spontaneous)

• That does NOT tell us anything about the RATE or

kinetics for the process

• Some spontaneous processes are fast such as explosions. Can you think of other examples?

(25)

6.5 Kinetics

• The rate of a reaction tells us how many molecules are reacting in a given period of time

• Give some examples for typical reaction rate units

(26)

6.5 Kinetics

• The reaction rate (the number of collisions that will

result in product production in a given period of time) is affected by multiple factors

1. The concentrations of the reactants 2. The Activation Energy

3. The Temperature

4. Geometry and Sterics

5. The presence of a catalyst

(27)

6.5 Rate Law Equations

• To quantify how much the reactant concentration affects the rate of reaction, the Rate Law equation is used

• The degree to which a change in [reactant] will affect the Rate is known as the order.

(28)

6.5 Rate Law Equations

• Consider a generic reaction that is known to be first

order with respect to A and zero order with respect to B: A + B  C + D

– Write the appropriate Rate Law

– How should he rate change if [A] were doubled?

(29)

6.5 Rate Law Equations

• Consider a generic reaction that is known to be first

order with respect to A and first order with respect to B: A + B  C + D

(30)

6.5 Rate Law Equations

• Consider a generic reaction that is known to be second order with respect to A and first order with respect to B: A + B  C + D

– How should he rate change if [B] were doubled?

(31)

6.5 Factors that Affect Rates

• Locate the Activation Energy in figure 6.13?

• Why must the free energy (G) increase

before the products can be formed?

• Why would E_act be

different for different reactions?

Free energy

(32)

6.5 Factors that Affect Rates

• Temperature is a measure of a system’s average

kinetic energy

• Would you expect there to be a temperature

below which the reaction rate is zero and above

which the reaction rate instantaneous for all

molecules? WHY or WHY

Free energy

(33)

6.5 R Factors that Affect Rates

• Why does a lower E_act result in a greater reaction rate?

Free energy

(G)

Free energy

(34)

(35)

• How might geometry and sterics affect the reaction rate?

• How might the presence of a catalyst affect the reaction rate?

Free energy

(36)

6.6 Energy Diagrams

• Distinguish between kinetics and thermodynamics

Free energy

(G)

Free energy

(37)

• For the energy diagram below, which pathway do you think is favored? WHY?

6.6 Kinetics vs Thermodynamics

• Will a decrease in temperature affect which pathway is favored?

• Will an increase in temperature affect which pathway is favored?

Free energy

(38)

6.6 Kinetics vs Thermodynamics

• For the energy diagram below, which pathway is kinetically favored?

• Which pathway is thermodynamically favored?

• How can temperature be used to control

which set of products predominate?

Free energy

(39)

6.6 Transition States vs Intermediates

Free energy

(40)

6.6 Transition States

• A transition state occurs at an energy maxima

• Transition states exist for a fleeting moment; they cannot be

isolated or directly observed

• Why are transition states so unstable?

Free energy

(41)

6.6 Intermediates

• An intermediate occurs at an energy minima

• Intermediates often exist long enough to observe

because bonds are NOT in the process of breaking or forming

Free energy

(42)

6.6 The Hammond Postulate

• Two points on an energy diagram that are close in energy should be similar in structure

Free energy

(43)

• For each of the diagrams below, will the transition state structure look more like the reactants or the products?

6.6 The Hammond Postulate

Free energy

(G)

Free energy

(44)

6.6 The Hammond Postulate

• Draw a reaction coordinate diagram for the generic exergonic reaction sequence below. Label the axes, reactants, products, intermediates, and transition states

A  B + C C + D  E Net reaction: A + D  B + E

(45)

6.7 Nucleophiles and Electrophiles

• A major focus in this course is on predicting reaction products for ionic reactions and explaining HOW such reactions work

• Ionic or polar reactions result from the force of attraction between opposite charges

• Ionic reactions are also guided by the octet rule

(46)

6.7 Nucleophiles

• When an atom carries a formal or partial negative charge and an available pair of electrons, it is

considered a nucleophile

• It will love to attack a nucleus. WHY?

• Explain how the molecules below are nucleophiles

(47)

6.7 Electrophiles

• When an atom carries a formal or partial positive charge and can accept a pair of electrons, it is

considered a electrophile

• It will love available electrons. WHY?

• Explain how the molecules above are electrophiles

(48)

6.7 Electrophiles

(49)

6.8 Mechanisms and Arrow Pushing

• We use arrows to show how electrons move when bonds break and form

• It will be a huge benefit in this course to master the skill of arrow pushing

• There are four main ways that electrons move in ionic reactions

1. Nucleophilic Attack

2. Loss of a Leaving Group

(50)

6.8 Nucleophilic Attack

• When you identify a nucleophilic site and an

electrophilic site, the arrow shows the nucleophile attacking

• The tail of the arrow starts on the electrons (- charge)

• The head of the arrow ends on a nucleus (+ charge)

(51)

6.8 Nucleophilic Attack

• Nucleophilic attack may appear to occur in two steps

• The alcohol is the nucleophile in this example. It attacks a carbon with a δ+ charge

• The second arrow shows the flow of negative charge. WHY is it necessary?

(52)

6.8 Loss of a Leaving Group

• Loss of a leaving group occurs when a bond breaks and one atom from the bond takes BOTH electrons

• For the molecule below, draw the structure that will result after the leaving group is gone

(53)

6.8 Proton Transfers

• Recall from Chapter 3 that a base is protonated when it uses a pair of electrons to take an H+ _{from the acid.}

• The acid retains its electron pair

• A group can also be deprotonated (sometimes shown by writing –H+ _{over the reaction arrow)}

(54)

6.8 Proton Transfers

• Multiple arrows may be necessary to show the

complete electron flow when a proton is exchanged

(55)

• Carbocations can be stabilized by neighboring groups through slight orbital overlapping called

hyperconjugation

(56)

6.8 Carbocation Rearrangements

• Hyperconjugation and induction can both be invoked to explain the stability trend below. HOW?

• If a carbocation can INTRAmolecularly rearrange to

(57)

6.8 Carbocation Rearrangements

• Two types of carbocation rearrangement are common

– Hydride shift

– Methyl shift

• Shifts can only occur from an adjacent carbon. WHY?

(58)

6.9 Combining Arrow Pushing Patterns

(59)

6.9 Combining Arrow Pushing Patterns

• Many times a single step in a mechanism will include more than one arrow pushing pattern

• Identify the patterns below

(60)

6.10 Arrow Pushing Rules

• To draw reasonable mechanisms, a few key rules should be followed

1. The arrow starts ON A PAIR OF ELECTRONS (a bonded pair or a lone pair)

(61)

6.10 Arrow Pushing Rules

• A few key rules should be followed

(62)

(63)

4. Draw arrows that follow the 4 key patterns we outlined

(64)

(65)

6.10 Arrow Pushing Rules

(66)

• When you encounter a carbocation, you must consider all possible rearrangements (Hydride and methyl shifts)

1. Identify all adjacent carbons

2. Identify all –H and –CH₃ groups on the adjacent

(67)

• In this case, a hydride shift will result in a more stable tertiary carbocation

6.11 Carbocation Rearrangements

• When you encounter a carbocation, you must consider all possible

rearrangements (Hydride and methyl shifts)

(68)

• Complete the same analysis for the molecule below

1. Identify all adjacent carbons

2. Identify all –H and –CH₃ groups capable of shifting

(69)

6.12 Reversible and Irreversible

Reaction Arrows

• Why are some reactions are drawn as equilibria and others are essentially irreversible?

(70)

6.12 Reversible and Irreversible

Reaction Arrows

• Consider nucleophilic attack

• Draw a mechanism for the reverse reaction

(71)

Reaction Arrows

• If the attacking nucleophile is also a good leaving group, it will be a reversible attack

– The reverse reaction will have a relatively low transition state energy (kinetically favored)

– The reactants and products of the reaction will be similar in energy allowing significant quantities of both to exist at

(72)

6.12 Reversible and Irreversible

Reaction Arrows

• If the attacking nucleophile is a poor leaving group, it will essentially be an irreversible attack

– The reverse reaction will have a relatively HIGH transition state energy (kinetically disfavored)

(73)

Reaction Arrows

• Consider loss or a leaving group

• Draw a mechanism for the reverse reaction

(74)

6.12 Reversible and Irreversible

Reaction Arrows

• Consider proton transfer

(75)

Reaction Arrows

• Consider proton transfer

(76)

Reaction Arrows

(77)

Reaction Arrows

• When considering thermodynamic equilibrium, in

(78)

Additional Practice Problems

• Consider electrocyclic ring opening

• Predict the sign (+ or -) for ΔS_sys

• In the reaction, 1 pi bond is converted into 1 sigma bond. Predict the sign (+ or -) for ΔH_sys

• Predict the sign (+ or -) for ΔS_surr

• Predict the sign (+ or -) for ΔG

(79)

Additional Practice Problems

(80)

• Reactants A and B can react by two different pathways. Pathway 1 is thermodynamically favored, and pathway 2 is kinetically favored. Draw a reaction coordinate

(81)

• Propose a mechanism for the reaction below that explains its Rate Law.

(82)

Additional Practice Problems

(83)

• Draw a mechanism for a generic nucleophilic attack followed by a proton transfer.

(84)