Acids, Bases, and Electron Flow in Organic Chemistry

Curved Arrows, Arrow Pushing, and Chemical Reactions

Curved Arrows as Electron Bookkeeping

Curved arrows are a fundamental tool in organic chemistry for tracking the movement of electrons during chemical reactions. They help visualize how bonds are broken and formed, and how charges are redistributed.

• Full-headed arrow: Represents the movement of a pair of electrons. The arrow starts at the electron source (such as a lone pair or a bond) and points to the electron sink (such as an atom or bond where electrons are accepted).

• Bond dissociation: When a molecule AB dissociates, a curved arrow shows the movement of electrons from the bond to one of the atoms, resulting in changes in formal charges.

• Multiple arrows: In more complex reactions, multiple curved arrows may be used to show simultaneous electron movements, such as in proton transfer or nucleophilic substitution.

Example: In the dissociation of HCl, a curved arrow from the H–Cl bond to Cl shows the formation of H+ and Cl−.

Multiple Bonds and Curved Arrows

• When a curved arrow begins at a multiple bond (such as a π bond), it indicates the movement of only two electrons from the π bond, while the σ bond remains intact.

• π bonds can act as electron sources (nucleophiles) or shift to allow a carbon atom to accept electrons (electrophilic addition).

Example: In electrophilic addition to an alkene, a curved arrow from the π bond to an electrophile shows the formation of a new σ bond.

Acids and Bases: The Brønsted-Lowry View

Brønsted-Lowry Acids and Bases

The Brønsted-Lowry theory defines acids and bases based on proton transfer.

• Acid: A species that donates a proton (H+) to a base.

• Base: A species that accepts a proton from an acid.

• Conjugate pairs: Species related by the gain or loss of a single proton. The acid forms its conjugate base after donating a proton, and the base forms its conjugate acid after accepting a proton.

Example: In the reaction of acetic acid with water:

$\mathrm{CH_3CO_2H} + \mathrm{H_2O} \rightleftharpoons \mathrm{CH_3CO_2^-} + \mathrm{H_3O^+}$

Acetic acid is the acid, water is the base, acetate is the conjugate base, and hydronium is the conjugate acid.

The Acid Dissociation Constant (Ka)

The strength of an acid is measured by its acid dissociation constant, Ka, which quantifies the equilibrium between the acid and its conjugate base in water.

$\mathrm{HA} + \mathrm{H_2O} \rightleftharpoons \mathrm{A^-} + \mathrm{H_3O^+}$

$K_a = \frac{[\mathrm{A^-}][\mathrm{H_3O^+}]}{[\mathrm{HA}]}$

• A larger Ka value indicates a stronger acid.

• Ka for an acid is related to Kb for its conjugate base.

pKa: A Brain-friendly Acidity Constant

Because Ka values can span many orders of magnitude, chemists use the negative logarithm, pKa, for convenience:

$pK_a = -\log_{10} K_a$

• A smaller pKa value means a stronger acid.

How Structure Affects Acid Strength

Structural Factors Influencing Acidity

Several structural factors influence the acidity of a proton in organic molecules. Understanding these helps predict and explain acid strength trends.

• Bond strength (H–A bond): As the size of atom A increases, the H–A bond becomes weaker, making it easier to lose H+ and increasing acidity.

• Electronegativity of A: Higher electronegativity of atom A increases the polarization of the H–A bond, making the proton more acidic.

• Inductive effects: Electronegative atoms near the acidic proton pull electron density away, increasing acidity. This effect is called induction.

• Resonance delocalization: If the conjugate base can delocalize the negative charge through resonance, it is stabilized, making the acid stronger.

Factor 1: H–A Bond Strength

• As the atom bonded to hydrogen increases in size (down a group in the periodic table), the bond becomes weaker and easier to break, increasing acidity.

• Hydrohalic acids (e.g., HF, HCl, HBr, HI) illustrate this trend: acidity increases from HF to HI.

Factor 2: Electronegativity of A in H–A

• As the electronegativity of A increases (across a period), the H–A bond becomes more polarized, and the proton is more easily lost.

• Covalent hydrides of second-row nonmetals (e.g., CH4, NH3, H2O, HF) show increasing acidity from left to right.

Factor 3: Inductive Effects

• Electronegative atoms near the acidic proton increase acidity by pulling electron density away (inductive effect).

• Example: Trifluoroethanol (CF3CH2OH) is much more acidic than ethanol (CH3CH2OH) due to the strong inductive effect of fluorine atoms.

Resonance Delocalization

• When the conjugate base can delocalize its negative charge via resonance, it is stabilized, making the parent acid stronger.

• Example: Acetic acid (CH3CO2H, pKa 4.75) is much more acidic than ethanol (CH3CH2OH, pKa 16) because the acetate ion is resonance stabilized, while the ethoxide ion is not.

Acid-Base Equilibria

Acid-Base Equilibria Favor the Weak

The position of acid-base equilibria can be predicted using pKa values. The equilibrium favors the side with the weaker acid (higher pKa).

• The equilibrium constant for a generic acid-base reaction is calculated from the Ka values of the acids on each side.

• If K > 1, products are favored; if K < 1, reactants are favored.

• Acid-base equilibria favor the side with the weaker acid (higher pKa).

Example: For the reaction:

$\mathrm{HA} + \mathrm{B} \rightleftharpoons \mathrm{A^-} + \mathrm{HB^+}$

The equilibrium favors the side with the acid of higher pKa.

Acids and Bases: The Lewis View

Lewis Theory of Acids and Bases

The Lewis theory broadens the definition of acids and bases to include electron pair transfer, not just proton transfer.

• Lewis acid: A species that accepts an electron pair (electrophile).

• Lewis base: A species that donates an electron pair (nucleophile).

• All Brønsted acids and bases are also Lewis acids and bases, but not all Lewis acids/bases are Brønsted acids/bases.

Example: BF3 + OEt2

• Diethyl ether (OEt2) donates a pair of electrons from oxygen, acting as a Lewis base (nucleophile).

• Boron trifluoride (BF3) accepts the electron pair at boron, acting as a Lewis acid (electrophile).

• A new bond forms between oxygen and boron, with both electrons coming from the Lewis base.

Importance of Lewis Theory

• Lewis acids (electrophiles) and bases (nucleophiles) are central to organic reaction mechanisms.

• Partially or fully negative atoms often act as nucleophiles; partially or fully positive atoms act as electrophiles.

• Every polar reaction can be viewed as an exchange of electrons between nucleophiles and electrophiles.

• Structural factors affecting Brønsted acid/base strength also influence Lewis acid/base strength.

Summary Table: Factors Affecting Acid Strength

Factor	Effect on Acidity	Example
Bond Strength (H–A)	Weaker bond increases acidity	HI > HBr > HCl > HF
Electronegativity of A	Higher electronegativity increases acidity	HF > H2O > NH3 > CH4
Inductive Effects	More electronegative groups increase acidity	CF3CH2OH > CH3CH2OH
Resonance Delocalization	Resonance stabilization increases acidity	Acetic acid vs. ethanol

Additional info: The above notes expand on the original content by providing definitions, examples, and equations for clarity and completeness, as would be expected in a modern organic chemistry study guide.