1. Organic Chemistry, 6th ed.
Marc Loudon and Jim Parise
Chapter 3 Acids and Bases. The Curved-Arrow Notation Electrophiles and Nucleophiles

2. Chapter 3 Overview
3.1 – Lewis Acid-Base Association Reactions 3.2 – Electron-Pair Displacement Reactions 3.3 – Using the Curved-Arrow Notation to Derive Resonance Structures 3.4 – BrØnsted-Lowry Acids and Bases 3.5 – Free Energy and Chemical Equilibrium 3.6 – The Relationship of Structure to Acidity

3. Electron-Deficient Compounds
Electron deficient compounds have atoms that are short of an octet of electrons.
In BF3, boron has only 6 e- and is electron deficient.
Electron deficient compounds act as Lewis acids in order to fulfill their valence-shell octet.
Electrophile Nucleophile Adduct
3.1 Lewis Acid-Base Association Reactions

4. Curved-Arrow Notation
A tool for keeping track of electron pairs in a chemical reaction.
The formation of a bond is described by a “flow” of electrons from the electron donor (Lewis base) to the electron acceptor (Lewis acid).
The electron flow is indicated by a curved arrow drawn from the electron source to the electron acceptor.
3.1 Lewis Acid-Base Association Reactions

5. Curved-Arrow Notation
Application to Lewis acid-base dissociation reaction:
Here the B-F bond that breaks is the source of the electron pair that is transferred to F-.
Note that in going from the reactants to the products that the total charge is conserved.
3.1 Lewis Acid-Base Association Reactions

6. What happens when electrons are directed at atoms that are not electron-deficient?
Not all electrophilic sites are actually electron-deficient.
An electron pair must depart from the atom receiving an electron pair. Bromomethane is not, strictly speaking, an electrophile.
The octet rule is preserved.
3.2 Electron-Pair Displacement Reactions

7. What happens when electrons are directed at atoms that are not electron-deficient?
Curved-Arrow Notation for Displacement
Displacement reactions require two arrows.
Here, the donated electron pairs originate from a bond.
3.2 Electron-Pair Displacement Reactions

8. Nucleophile, Electrophile, and Leaving Group
Nucleophile – a species that donates an electron pair to form a new bond. Also called a Lewis Base.
Electrophile – a species that accepts an electron pair from the nucleophile. Sometimes, as below, applied to compounds that have an electrophilic center.
Leaving Group – the group that accepts electrons from the breaking bond.
3.2 Electron-Pair Displacement Reactions

9. Nucleophile, Electrophile, and Leaving Group
In the reverse reaction, the roles are reversed.
Strictly speaking methylammonium ion is neither an electrophile nor a Lewis acid, but it does have an “electrophilic” site. It is an adduct of H3C+ and :NH3. The bromide ion, however, is an actual nucleophile.
3.2 Electron-Pair Displacement Reactions

10. Nucleophile, Electrophilic Center, and Leaving Group
The nucleophilic electron pair can originate from a bond in addition to an unshared electron pair.
3.2 Electron-Pair Displacement Reactions

11. Curved-Arrow Notation for Resonance
Resonance structures always differ only by movement of electrons; nuclei do not move.
Curved-arrow notation is ideal to help derive resonance structures.
3.3 Using the Curved-Arrow Notation to Derive Resonance Structures

12. Proton Transfers = Curved Arrows the Wrong Way
Curved-arrows show the movement of electron pairs not nuclei
Electrons are responsible for chemistry!
3.2 Electron-Pair Displacement Reactions

13. Brønsted Acids and Bases
BrØnsted Acid: A species that donates a proton (H+)
BrØnsted Bases: A species that accepts a proton.
A BrØnsted acid-base reaction is an electron-pair displacement reaction where the hydrogen is the electrophilic center (the “proton itself” is the electrophile).
3.4 Brønsted-Lowry Acids and Bases

14. Conjugate Acids and Bases
When a BrØnsted acid loses a proton, its conjugate base is formed.
When a BrØnsted base gains a proton, its conjugate acid is formed
conj. base
base
conj. acid
acid
3.4 Brønsted-Lowry Acids and Bases

15. Really Important
The most important single thing is that you can’t say a molecule is an acid (or Lewis acid or Electrophile) or a base (or Lewis base) by looking only at its structure. Acid/base labels define behavior, not structure. For example “a Brønsted acid is a proton donor in the specific transformation being evaluated”. However, potential nucleophilic and electrophilic sites can be identified by examining a Lewis structure.
Actual Lewis acid (electrophile) behavior is the rarest event. It requires accepting an electron pair (thus making a new bond) without any other bonds being broken. The other partner in a nucleophilic displacement reaction is not an electrophile even though it has an electrophilic center. Any compound in a reaction that can be classed as a Brønsted acid is NOT an electrophile: the proton is the electrophile.
No electrophile in : H2O H-Cl goes to H3O Cl-

16. Amphoteric Compounds
Compounds that can act as either an acid or a base are called amphoteric.
A species can only be defined as an acid or base based on a specific reaction.
Water is amphoteric. It can be either an acid or a base. In the example below, water is acting as a base.
3.4 Brønsted-Lowry Acids and Bases

17. Organic Reactions (not all proton transfers)
The Lewis acid-base concept is central to many reactions in organic chemistry.
For example:
…is analogous to:
3.4 Brønsted-Lowry Acids and Bases

18. Strengths of BrØnsted Acids
The strength of a Brønsted acid is determined by how well it transfers a proton to a Brønsted base.
The standard base traditionally used is water
The equilibrium constant is given by:
3.4 Brønsted-Lowry Acids and Bases

19. The Dissociation Constant
As the solvent, [H2O] effectively remains constant.
Another constant, Ka, the dissociation constant, is defined:
Each acid has its own unique dissociation constant.
A larger Ka indicates more H+’s are transferred
3.4 Brønsted-Lowry Acids and Bases

20. pKa Values
Direct pKa determination in an aqueous environment is limited to acids that are less acidic than H3O+ and more acidic than H2O.
Organic chemistry is typically conducted in a non- aqueous environment where pKa values differ significantly for the same acids determined in water. However, relative, pKa values are often nearly the same – the order of acidity the same.
Ka of water ( ) is different than Kw of water (10-14 M2). Be careful to not confuse the two. The pKa of water is 15.7.
3.4 Brønsted-Lowry Acids and Bases

21. Relative Strengths of Some Acids and Bases
___
weak
____
moderate
acidity
______
strong 
3.4 Brønsted-Lowry Acids and Bases

22. Strengths of BrØnsted Bases
If a base is weak, its conjugate acid is strong.
If a base is strong, its conjugate acid is weak.
H-O- is a stronger base than H2O just as H2O is a weaker acid that H3O+.
3.4 Brønsted-Lowry Acids and Bases

23. Equilibria in Acid-Base Reactions
Determine by comparing pKa of both acids.
The equilibrium in the reaction of an acid and a base always favors the side of the weaker acid and weaker base.
3.4 Brønsted-Lowry Acids and Bases

24. Structure Acidity Relationship
We can use the structure of organic molecules to predict trends in their chemical properties (in this case acidity).
Three main considerations for structural effect on acidity:
Element Effect
Charge Effect
Polar Effect
3.6 The Relationship of Structure to Acidity

25. The Element Electronegativity Effect
Look at H-X: electronegativity of X and bond strength.
3.6 The Relationship of Structure to Acidity

26. The Charge Effect
Positively charged compounds give up a proton more readily than neutral ones.
pKa of H3O+ = vs pKa of H2O = 15.7
pKa of H4N+ = vs pKa of H3N = 35
The conjugate base of H3N ( H2N:- ) is one the strongest bases available, as NaNH2.
3.6 The Relationship of Structure to Acidity

27. Relative Acidity versus Relative Basicity of Conjugate Bases
How to think about it & how to predict relative activities
1) Basicity is proton affinity of species. Relative basicity is how much it needs a proton for stability versus some other base.
2) If a species has very little or no affinity for a proton, it’s conjugate acid will be a strong(er) acid.
3) The attraction of an electron pair for a proton decreases when
a) it’s on a more electronegative atom – this work perfectly when the atoms are in the same period (about the same size).
b) with increased delocalization either by resonance or by being in a larger orbital (bigger/heavier atoms)

28. Resonance Delocalization of Electrons Decreases Basicity
Carboxylic acids illustrate the effect
The conjugate base is resonance-stabilized
Compare the pKa of a carboxylic acid (~5, resonance stabilized) to the pKa of an alcohol (~16, no delocalization of the negative charge ( R-O- , not resonance stabilized).
3.6 The Relationship of Structure to Acidity

29. Inductive Effects (another polar effect)
Increasing the number of nearby electronegative atoms in an acid increases acidity:
3.6 The Relationship of Structure to Acidity

30. The Inductive Effect is Distance Dependent
Varying the distance of the electronegative atom to the acidic hydrogen:
3.6 The Relationship of Structure to Acidity

31. Inductive Effects are Polar Effects
Halogens and other electronegative groups exert an electron-withdrawing polar effect.
This lowers the pKa of carboxylic acids; also applies to alcohols.
Other groups can exert an electron-donating polar effect.
This would raise the pKa of carboxylic acids.
3.6 The Relationship of Structure to Acidity

32. Explaining Polar Effects, an example
Electrostatic interactions can be stabilizing or destabilizing.
Electronegative substituents increase the acidity of carboxylic acids (inductive effect or polar effect).
3.6 The Relationship of Structure to Acidity

33. Relative Basicity of Conjugate Bases versus predicting relative acidity
A priori prediction without access to X – H bond strengths difficult for relative acidity. Relative basicity is mostly proton affinity. Easier to envision and predict.
1) If a species has very little or no affinity for a proton, it is a weak(er) base and it’s conjugate acid will be a strong(er) acid.
2) The attraction of an electron pair for a proton decreases when
a) it’s on a more electronegative atom – this work perfectly when the atoms are in the same period (about the same size). b) with increased delocalization either by resonance or by being in a larger orbital (bigger/heavier atoms)
F - is more basic than Br -