Introduction to Organic chemistry: Chapter 2 NOTES

Acids and Bases

A great many organic reactions are acid–base reactions.
Many organic reactions are catalyzed by proton-donating acids, such as H3O+ and CH3OH2+. Others are catalyzed by Lewis acids, such as AlCl3 .

2.1 What Are Arrhenius Acids and Bases?
According to Arrhenius definitions, an acid is a substance that dissolves in water to produce H+ ions, and a base is a substance that dissolves in water to produce OH_ ions.

2.2 What Are Brønsted–Lowry Acids and Bases?
According to Brønsted–Lowry definitions, an acid is a proton donor, a base is a proton acceptor, and an acid–base reaction is a proton-transfer reaction.
According to the Brønsted–Lowry definitions, any pair of molecules or ions that can be interconverted by the transfer of a proton is called a conjugate acid–base pair.
When an acid transfers a proton to a base, the acid is converted to its conjugate base.
When a base accepts a proton, the base is converted to its conjugate acid.
By studying the examples of conjugate acid–base pairs in the table, note the following points:
1. An acid can be positively charged, neutral, or negatively charged. Examples of these charge types are H3O+, H2CO3 , and H2PO4+
2. A base can be negatively charged or neutral. Examples of these charge types are Cl_ and NH3.
3. Acids are classified as monoprotic (HCl, HNO3 and CH3COOH), diprotic (H2SO4 and H2CO3), or triprotic (H3PO4).
4. Several molecules and ions appear in both the acid and conjugate base columns; that is, each can function as either an acid or a base. For example the bicarbonate ion, HCO3_.
5. There is an inverse relationship between the strength of an acid and the strength of its conjugate base. The stronger the acid, the weaker is its conjugate base.

2.3 How Do We Measure the Strength of an Acid or Base?
A strong acid or strong base is one that ionizes completely in aqueous solution.
When HCl is dissolved in water, a proton is transferred completely from HCl to H2O to form Cl_ and H3O+. We conclude that HCl is the stronger acid and H3O+ is the weaker acid. Similarly, H2O is the stronger base and Cl_ is the weaker base.
Strong acids in aqueous solution: HCl, HBr, HI, HNO3 , HClO4 , and H2SO4.
Strong bases in aqueous solution: LiOH, NaOH, KOH, Ca(OH)2 , and Ba(OH)2 .
A weak acid or weak base is one that only partially ionizes in aqueous solution.
Most organic acids and bases are weak. For example: carboxylic acids.
The equation for the ionization of a weak acid, HA, in water and the acid ionization constant Ka for this equilibrium are, respectively.
The acids with ionization constants Ka considerably less than 1.00 are weak acids.

2.4 How Do We Determine the Position of Equilibrium in an Acid–Base Reaction?
In an acid–base reaction, the position of equilibrium always favors reaction of the stronger acid and stronger base to form the weaker acid and weaker base.
At equilibrium, the major species present are the weaker acid and weaker base.
For example, In the reaction between acetic acid and ammonia, therefore, the equilibrium lies to the right, and the major species present are acetate ion and ammonium ion.

2.5 What Are the Relationships between Acidity and Molecular Structure?
The most important factor in determining the relative acidities of organic acids is the relative stability of the anion, A⁻, formed when the acid, HA, transfers a proton to a base.
We can understand the relationship involved by considering (A) the electronegativity of the atom bonded to H, (B) resonance, (C) the inductive effect, and (D) the size and delocalization of charge on A⁻.

A. Electronegativity: Acidity of HA within a Period of the Periodic Table
The more electronegative an atom is, the greater its ability to sustain electron density around itself.
The greater the electronegativity of A, the greater the stability of the anion A⁻, and the stronger the acid HA.

B. Resonance Effect: Delocalization of the Charge in A⁻
Carboxylic acids are weak acids (pKa = 4 to 5).
Alcohols are very weak acids (pKa = 15 to 18).
Carboxylic acid is a stronger acid than an alcohol because there is no resonance stabilization in an alkoxide anion, but in carboxylate ion there is a delocalization of the charge.

C. The Inductive Effect: Withdrawal of Electron Density from the HA Bond
The inductive effect is the polarization of electron density transmitted through covalent bonds by a nearby atom of higher electronegativity.
For example trifluoroacetic acid (pKa 0.23) is stronger than acetic acid (pKa 4.76) because Fluorine is more electronegative than carbon and polarizes the electrons of the C―F bond, creating a partial positive charge on the carbon of the _CF3 group, which withdraws electron density from the negatively charged _CO2 group.
The withdrawals of electrons delocalizes the negative charge and makes the conjugate base of trifluoroacetic acid more stable than the conjugate base of acetic acid.

D. Size and the Delocalization of Charge in A
The more stable the anion, the greater the acidity of the acid.
The larger the volume over which the charge of an anion (or cation) is delocalized, the greater the stability of the anion.
Recall from general chemistry that atomic size is a periodic property:
1. For main group elements, atomic radii increase going down a group in the Periodic Table. Thus, for the halogens (Group 8A elements), iodine has the largest atomic radius and fluorine has the smallest (I > Br > Cl > F).
2. Anions are always larger than the atoms from which they are derived. For anions, the nuclear charge is unchanged, but the added electron(s) introduce new repulsions and electron clouds شسwell. Among the halide ions, I⁻ has the largest atomic radius and F⁻ has the smallest (I⁻ > Br⁻ > Cl⁻ > F⁻).
Hence, HI is the strongest acid and that HF is the weakest.
The negative charge on iodide ion is delocalized over a larger area than is the negative charge on bromide ion.

2.6 What Are Lewis Acids and Bases?
According to the Lewis definition, an acid is a species that can form a new covalent bond by accepting a pair of electrons; a base is a species that can form a new covalent bond by donating a pair of electrons.
“donating” means that the electron pair is shared with another atom to form a covalent bond.
Isolated protons do not exist in solution; rather, a proton attaches itself to the strongest available Lewis base. For example, when HCl is dissolved in methanol, the strongest available Lewis base is a CH3OH molecule.
An oxonium ion is an ion that contains an oxygen atom with three bonds and bears a positive charge.