Chapter 2: Acid-Base Concepts and Electron Flow
Electrons move; atoms do not
Every bond-making and bond-breaking event can be reduced to the movement of an electron pair from a site that can share electrons (a Lewis base) to a site that can accept them (a Lewis acid). Lewis bases include lone-pair donors such as water, ammonia, and halide ions, as well as alkenes and other pi systems. Lewis acids include protons, carbocations, metal ions, and atoms with incomplete valence shells. A practical recognition sequence is to mark lone pairs and pi bonds first, then check for positive charge, an empty orbital, or an incomplete octet. Thus nitrogen in trimethylamine and oxygen…
Formal charge: electron bookkeeping
A Lewis structure is complete only when every atom's share of the electrons has been counted. Formal charge is that count: the number of valence electrons an isolated atom brings, minus the nonbonding electrons it keeps, minus half of the electrons it shares. The formal charges in a structure always sum to the net charge of the species, which makes the calculation a running error check on the drawing. A few patterns cover most of organic chemistry. Oxygen with two bonds and two lone pairs is neutral; with three bonds it is +1, as in the hydronium ion (H₃O⁺); with one bond it is −1, as in…
Two rules for drawing curved arrows
Curved arrows are a bookkeeping tool. The tail of the arrow marks where the electrons are coming from, and the head shows where they are going. Two fundamental moves exist. In an association, a lone pair or pi bond donates to an electron-deficient acceptor, and one curved arrow suffices. In a displacement, an incoming donor pushes out a leaving group: one arrow forms the new bond between donor and central atom while a second arrow moves the old bonding electrons onto the departing group. Both arrows must be drawn; using only one is the most common error and produces a structure that violates…
Resonance: delocalized electrons and the hybrid
Acetate presents a problem that one Lewis structure cannot solve. The drawing puts a double bond to one oxygen and a negative charge on the other, yet both carbon–oxygen bonds have the same measured length and both oxygens carry the same charge. Two drawings are needed, differing only in where the electrons are placed. Such drawings are resonance contributors, written with a double-headed arrow between them, and the curved arrows of the previous section interconvert them on paper: a lone pair on the negatively charged oxygen becomes a π bond, and the original π bond becomes a lone pair on the…
Proton transfer and conjugate pairs
Brønsted-Lowry theory focuses on proton transfer. When an acid donates H+, the electrons that held the proton to the molecule remain — forming the conjugate base. When a base accepts H+, it uses a lone pair to bind the proton, generating the conjugate acid. Acid and conjugate base are related by one proton; base and conjugate acid are related by one proton. A molecule that can act as either an acid or a base depending on what it encounters is called amphoteric. Water is the classic example: it accepts H+ from a stronger acid (behaving as a base) or donates H+ to a stronger base (behaving as…
The pKa scale
The equilibrium expression for proton transfer to water gives Ka = [H₃O⁺][A⁻]/[HA]. Ka values for organic acids range from roughly 10⁻⁵ for carboxylic acids down to 10⁻⁵⁰ for sp3 C–H bonds — a range too wide to compare directly. The pKa scale compresses these into manageable numbers: HCl has pKa ≈ –7, acetic acid ≈ 4.7, ethanol ≈ 16, and acetylene ≈ 25. Under the same solvent, temperature, and analytical concentration, a stronger acid is more dissociated at equilibrium and generally produces a lower pH; pH itself still depends on the composition and concentration of the particular solution. A…
Typical pKa ranges for organic functional groups
The position of a proton on the pKa scale governs its reactivity toward bases. Carboxylic acids (pKa 4–5) are deprotonated by amines, carbonates, and hydroxide under mild conditions. Phenol has pKa ≈ 10 because its phenoxide conjugate base delocalizes negative charge into the aromatic ring. Protonated alcohols and ethers (oxonium ions, R–OH₂⁺) have pKa values near −2 and occur under strongly acidic conditions. Protonated amines are ammonium ions (R–NH₃⁺), commonly with pKa values near 9–11; this number describes the ammonium conjugate acid, not the neutral amine. Neutral amine N–H bonds have…
Predicting which side of equilibrium is favored
Proton transfer is reversible, and predicting the equilibrium position requires comparing the two possible acids: the starting acid and the conjugate acid produced in the forward reaction. Equilibrium lies on the side bearing the weaker acid (higher pKa) because that side represents the more stable distribution of charges and bonds. If a neutral amine reacts with a carboxylic acid (pKa ≈ 5), the product-side acid is the ammonium ion, the amine's conjugate acid (pKa ≈ 10). Equilibrium therefore favors the ammonium carboxylate salt. The equilibrium constant Keq ≈ 10^(pKa(product acid) –…
Five factors that control acidity
Acidity can always be analyzed through conjugate-base stability: the more stable the conjugate base, the stronger the acid. Electronegativity: within a row of the periodic table, placing the negative charge on a more electronegative atom (O vs N vs C) increases stability. Size: down a group, larger, more polarizable atoms stabilize negative charge better than small atoms, even though they are less electronegative — this is why H2S is more acidic than H2O despite oxygen being more electronegative. Hybridization: an sp carbon has more s character than sp2 or sp3, holds electrons closer to the…
Noncovalent interactions: the forces between molecules
The covalent bonds inside a molecule do not explain why water boils at 100 °C while methane, a molecule of similar size, boils at −162 °C. The difference lies in noncovalent interactions — attractions between molecules that are far weaker than covalent bonds but act in enormous numbers. Four kinds matter here, and they form a hierarchy of typical strength: ion–ion attractions between fully charged species, hydrogen bonds, dipole–dipole attractions, and London dispersion forces. A hydrogen bond forms when a hydrogen bonded to oxygen or nitrogen is shared with a lone pair on a neighboring…
ΔG° and the equilibrium constant
Every acid-base equilibrium has an associated standard free-energy change ΔG° = –RT ln Keq where R = 8.314 J mol⁻¹ K⁻¹ and T is in Kelvin. A reaction with ΔG° = 0 kJ/mol has Keq = 1 (equal amounts of reactants and products at standard conditions). Because Keq is exponential in ΔG°/RT, even modest energetic differences produce large equilibrium shifts: –5.7 kJ/mol (at 25 °C) corresponds to a tenfold preference for products. This amplification means that reactions do not need enormous thermodynamic driving forces to go essentially to completion — a Keq of 10^5 (ΔG° ≈ –28 kJ/mol) already…
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