Polar Covalent Bonds; Acids and Bases – Resonance, Acid-Base Theories, and Organic Examples

Chapter 2: Polar Covalent Bonds; Acids and Bases

2.6 Drawing Resonance Forms

Resonance is a key concept in organic chemistry, describing the delocalization of electrons within molecules that have conjugated bonds or lone pairs adjacent to multiple bonds. Resonance structures are different Lewis structures for the same molecule, showing possible distributions of electrons.

• Resonance Forms: Different resonance forms of a substance do not have to be equivalent. For example, in acetone anion, one resonance form has the negative charge on carbon, another on oxygen.

• Curved Arrows: A curved arrow always indicates the movement of electrons, not atoms.

• Valency Rules: Resonance forms must obey normal rules of valency. Structures violating the octet rule (e.g., 10 electrons on carbon) are not valid resonance forms.

• Resonance Hybrid: The actual structure (resonance hybrid) is more stable than any individual resonance form due to electron delocalization.

Example: 2,4-Pentanedione has multiple resonance structures, with electrons delocalized between oxygen atoms and the central carbon. The carbonate ion also has three equivalent resonance structures, each with a double bond to a different oxygen.

2.7 Acids and Bases: The Brønsted–Lowry Definition

The Brønsted–Lowry theory defines acids and bases based on their ability to donate or accept protons (H+ ions).

• Acid: A substance that donates a hydrogen ion (proton).

• Base: A substance that accepts a hydrogen ion.

• Conjugate Acid-Base Pairs: When an acid donates a proton, it forms its conjugate base; when a base accepts a proton, it forms its conjugate acid.

General Equation:

$\ce{HA + B <=> A^- + HB^+}$

Example: Acetic acid ($\ce{CH3COOH}$) donates a proton to water, forming acetate ion ($\ce{CH3COO^-}$) and hydronium ion ($\ce{H3O^+}$).

2.8 Acid and Base Strength

The strength of an acid in water is quantified by the acidity constant ($K_a$) and its logarithmic counterpart, $pK_a$.

• Acid Dissociation Equilibrium: $\ce{HA + H2O <=> A^- + H3O^+}$

• Acidity Constant: $K_a = \frac{[\ce{H3O^+}][\ce{A^-}]}{[\ce{HA}]}$

• pKa Definition: $pK_a = -\log K_a$

• Water's pKa: The $pK_a$ value for water is 15.74.

Table: Relative Strengths of Some Common Acids and Their Conjugate Bases

Acid	Name	pKa	Conjugate Base	Name
CH3CH2OH	Ethanol	16.0	CH3CH2O-	Ethoxide ion
H2O	Water	15.74	HO-	Hydroxide ion
HCN	Hydrocyanic acid	9.31	CN-	Cyanide ion
H2PO4-	Dihydrogen phosphate ion	7.21	HPO42-	Hydrogen phosphate ion
CH3CO2H	Acetic acid	4.76	CH3CO2-	Acetate ion
H3PO4	Phosphoric acid	2.16	H2PO4-	Dihydrogen phosphate ion
HNO3	Nitric acid	-1.3	NO3-	Nitrate ion
HCl	Hydrochloric acid	-7.0	Cl-	Chloride ion

Additional info: Table shows the trend that as acid strength increases (lower pKa), the conjugate base becomes weaker.

2.9 Predicting Acid–Base Reactions from pKa Values

The direction of acid-base reactions can be predicted using pKa values. The proton (H+) will transfer from the stronger acid to the stronger base, forming a weaker acid and a weaker base.

• Reaction Direction: Reaction proceeds from the side with the stronger acid and base to the side with the weaker acid and base.

Example:

$\ce{CH3COOH + OH^- -> H2O + CH3COO^-}$

Acetic acid (stronger acid, $pK_a = 4.76$) reacts with hydroxide ion (stronger base) to form water (weaker acid, $pK_a = 15.74$) and acetate ion (weaker base).

2.10 Organic Acids and Organic Bases

Organic acids and bases are common in biological and chemical systems. Their strength and reactivity depend on molecular structure and the stability of their conjugate ions.

Some Organic Acids

• Methanol: $pK_a = 15.54$ (weak acid)

• Acetic Acid: $pK_a = 4.76$ (strong acid)

• Acetone: $pK_a = 19.3$ (very weak acid)

Anion Stabilization: The stability of the conjugate base (anion) increases if the negative charge is on a highly electronegative atom or is delocalized by resonance.

• Acetone anion is stabilized by resonance between oxygen atoms.

• Carboxylic acids (–CO2H group) are abundant in living organisms and are involved in metabolic pathways (e.g., acetic acid, pyruvic acid, citric acid).

Some Organic Bases

• Organic bases contain an atom with a lone pair of electrons that can bond to H+ (e.g., methylamine, methanol, acetone).

• Amino acids like alanine can exist in uncharged or zwitterion forms, depending on pH.

2.11 Acids and Bases: The Lewis Definition

The Lewis definition broadens the concept of acids and bases to include electron pair transfer.

• Lewis Acid: Electron pair acceptor (e.g., Mg2+, BF3).

• Lewis Base: Electron pair donor (e.g., H2O, NH3).

• Lewis Acid-Base Complex: Formed when a Lewis base donates an electron pair to a Lewis acid.

Example: Boron trifluoride (Lewis acid) reacts with dimethyl ether (Lewis base) to form an acid-base complex.

Some Lewis Acids

• Neutral proton donors: H2O, HCl, HBr, HNO3, H2SO4

• Carboxylic acids, phenols, alcohols

• Cations: Li+, Mg2+

• Metal compounds: AlCl3, TiCl4, FeCl3, ZnCl2

Some Lewis Bases

Compound	Type
H2O	Water
CH3CH2OH	An alcohol
CH3OCH3	An ether
CH3CHO	An aldehyde
CH3COCH3	A ketone
CH3COCl	An acid chloride
CH3COOH	A carboxylic acid
CH3COOCH3	An ester
CH3CONH2	An amide
CH3NH2	An amine
CH3SCH3	A sulfide
Organotriphosphate ion	An organotriphosphate ion

Additional info: Lewis bases are characterized by the presence of lone pairs that can be donated to electron-deficient species (Lewis acids).