Notes on Reaction Mechanisms, Acids/Bases, and Equilibria
Reaction mechanisms and arrow pushing
Purpose: Reaction mechanisms explain what actually happens in a reaction, including the step-by-step movement of electrons and the sequence of events.
Mechanism types: Sometimes reactions are single-step; other times they are multistep.
Example of single-step: a + b → c + d (classic from general chemistry, where a and b are reactants, c is intermediate or product depending on context, and d is a product).
Multistep example: a + b → c, followed by c → d. In this case, the product of the first step (c) becomes the reactant for the next step, i.e., an intermediate is formed.
Intermediates: The species on the left-hand side of the arrow that feed into the next step (such as c in the example) are called intermediates.
Arrow-pushing rules (moving electrons, not atoms):
Curved arrows always move electrons from regions of higher electron density to lower electron density (electrons flow down their gradient).
Atoms themselves do not move via curved arrows; only electrons move.
Only certain electrons are “accessible” for movement: electrons in π bonds (double/triple bonds) and lone-pair electrons (nonbonding electrons).
When electrons are moved toward a carbon, consider the rest of the bonding: if that carbon currently holds the appropriate valence, you often need to push the adjacent π bond or lone pair to another atom, often onto a more electronegative atom.
If the push would violate valence or place electrons where they cannot be accommodated, that arrow movement is not allowed.
Two-electron moves are shown with a full-headed curved arrow; one-electron moves (radical-type steps) are shown with a half-headed arrow.
Simple mechanism sketching approach: drawing nonbonding electrons and sometimes visualizing the bonds helps see which bonds break and how electrons reorganize during the reaction.
Example: HCl reacting with water (acid-base mechanism)
H–Cl is a polar covalent bond with a dipole: partial positive on H, partial negative on Cl.
Water has lone-pair electrons on oxygen to donate.
Mechanistic step: the lone pair on water attacks the acidic proton of HCl, forming H3O+; the H–Cl bond breaks with electrons moving onto Cl, forming Cl−.
Resulting species: hydronium ion ext{H}_3 ext{O}^+ and chloride ion ext{Cl}^-
Visualization: this illustrates movement of electrons (not atoms) and bond-breaking forming a clear two-step sequence in a proton-transfer process.
Lewis acids, Lewis bases, Bronsted–Lowry acids/bases
Lewis acid: an electron-pair accepter; Lewis base: an electron-pair donor.
Example: ext{BF}_3 is a Lewis acid (it accepts electrons from a donor).
Example: water (H₂O) can donate electrons to a Lewis acid, acting as a Lewis base in many cases.
Relationship to Bronsted–Lowry (BL):
BL acids/bases are hydrogen donors/acceptors; every BL acid/base is also a Lewis acid/base, respectively, because there is an electron acceptor/donor involved in the reaction.
However, not all Lewis acids/bases are BL acids/bases, since BL definitions require hydrogen transfer.
Example: ext{BF}_3 is a Lewis acid but not a BL acid/base (no hydrogen involvement).
Carbocations and carbanions as Lewis acids/bases:
Carbocation (C⁺, sp²) is electron-deficient and can act as a Lewis acid (electrophile).
Carbanion (C⁻, sp³) has lone electrons and can act as a Lewis base (nucleophile).
Electrophiles and nucleophiles (built on charge/positioning):
Electrophiles = electron-poor centers (often Lewis acids) that seek electrons.
Nucleophiles = electron-rich species (often Lewis bases) that donate electrons.
Carbonyl carbons are classic electrophilic centers with a partial positive charge, attracting nucleophiles.
Note: nucleophiles are attracted to positive or electron-deficient centers; electrophiles do not need a full positive charge to act as reactive centers.
Carbons, heterolysis vs homolysis; carbocations and carbanions
Heterolysis: a bond breaks and both electrons go to one atom (usually the more electronegative one). This can generate a cation (carbocation) on the other atom.
Homolysis: a bond breaks evenly, each atom taking one electron (radical pair).
Carbocation basics:
Formed via heterolysis that leaves carbon positively charged (C⁺).
Typical coordination around the cation: three substituents (three regions) → trigonal planar geometry → sp² hybridization.
Carbanion basics:
Formed when bond electrons stay with carbon, generating C⁻ (carbanion).
Four substituents around carbon → tetrahedral geometry → sp³ hybridization.
Consequences for reactivity:
Carbocations are Lewis acids (electron pair acceptors).
Carbanions are Lewis bases (electron pair donors).
Electrophiles vs nucleophiles (reframing from Lewis acid/base)
In many organic reactions, viewing species as electrophiles or nucleophiles (rather than strictly as Lewis acids or bases) helps predict the course of the reaction.
Carbonyl groups as electrophiles:
The carbonyl carbon bears a partial positive charge and is an electrophilic center.
Nucleophiles attack the carbonyl carbon, moving electron density toward the more electronegative oxygen.
Core idea: reactions are guided by electron density and charge distributions; toward regions of low electron density (electrophiles) are sites of nucleophilic attack; toward regions of high electron density are attacked by electrophiles.
Quick review: Equilibrium concepts in acid–base chemistry
Example dissociation in water: acetic acid in water forms acetate and hydronium; reaction is HA + H₂O ⇌ A⁻ + H₃O⁺.
Equilibrium constant (acid dissociation constant):
Ka = rac{[A^-][ ext{H}3 ext{O}^+]}{[HA]}
Generalized acid-base notation: HA ⇌ A⁻ + H⁺; in water, H⁺ is represented as H₃O⁺.
pKa relationship: ext{p}Ka = -\log Ka
pH relationship (for aqueous solutions): ext{pH} = -\log [ ext{H}^+]
Interpreting pKa values:
Smaller pKa => stronger acid; larger pKa => weaker acid.
Example set of acids mentioned:
Benzoic acid (pKa roughly 4–5 range in water context)
Acetic acid (pKa ≈ 4.76)
Trifluoroacetic acid (pKa ≈ 0.23)
Hydrochloric acid (HCl) effectively very strong with a negative pKa (often treated as <0 in practice)
Water as an acid: pKa ≈ 14 (in the context of H₂O ⇌ H⁺ + OH⁻ in water)
Relative strength and conjugate bases: stronger acids have weaker conjugate bases, weaker acids have stronger conjugate bases.
Example: Cl⁻ is a very weak base; acetate (A⁻ from acetic acid) is stronger; hydroxide (OH⁻) is stronger than acetate.
Predicting reaction direction in acid–base equilibria:
Equilibrium favors formation of the weaker acid and weaker base.
Example reasoning: when comparing acids with very different pKa values, the equilibrium lies heavily toward the side with the weaker acid and weaker base.
Illustration: if you start with a stronger acid and a stronger base, equilibration will shift toward the side with the weaker acid/base (e.g., formation of water in neutralizations is highly favorable).
Practical application: use acid–base equilibria to predict separations and product formation in synthesis and purification.
Applications and real-world relevance
Salt formation and solubility: converting a carboxylic acid to its salt (e.g., benzoic acid + NaOH → sodium benzoate) increases water solubility, aiding separation and purification.
Separatory funnel technique: use acid–base chemistry to move compounds between aqueous and organic phases to separate mixtures.
Solubility trends for large hydrophobic groups: bulky hydrophobic R groups decrease water solubility; salts improve solubility by introducing charge.
Practical workflow: in synthesis and isolation, you can manipulate acidity/basicity to control solubility and partitioning between layers for purification.
Periodic trends in acidity
Trend: acidity increases down a group in the periodic table for halogen hydrides (HF, HCl, HBr, HI).
Rationale:
Down the group, atoms become larger, bond length increases, bond strength weakens, making proton donation easier (stronger acids).
Examples cited:
HF < HCl < HBr < HI in acidity strength (acid strength increases down the group).
Additional notes and clarifications from the lecture
When discussing equilibrium and pKa, the speaker used approximate numerical anchors to illustrate concepts (e.g., water pKa ≈ 14 as a reference, HCl as very strong with effectively a negative pKa, trifluoroacetic acid stronger than acetic acid).
The “RKa” term was used to refer to the acid dissociation constant Ka; the standard formula and interpretation apply.
The distinction between Lewis and Bronsted–Lowry definitions is key: all BL acids/bases are Lewis acids/bases, but not all Lewis acids/bases are BL acids/bases due to the hydrogen requirement in BL definitions.
The narrative emphasizes visualization of electron flow (arrows) and the importance of considering the subsequent redistribution of bonds after electron movement.
The material connects fundamental concepts (electrophiles, nucleophiles, acids/bases, equilibrium) to practical organic chemistry applications (predicting reaction directions, separations, and solubility changes).
Reaction mechanisms detail step-by-step electron movement using curved arrows, always from high to low electron density, involving π bonds or lone pairs. Reactions can be single-step or multistep, forming intermediates. An example is the acid-base reaction of HCl with water, where water's lone pair attacks the acidic proton, forming ext{H}_3 ext{O}^+ and ext{Cl}^-.
Lewis acids are electron-pair acceptors, and Lewis bases are electron-pair donors. Brønsted–Lowry (BL) acids/bases are proton donors/acceptors; all BL acids/bases are Lewis acids/bases, but not vice versa. Electrophiles are electron-poor centers (Lewis acids) that seek electrons, while nucleophiles are electron-rich species (Lewis bases) that donate electrons, such as carbonyl carbons being electrophilic centers. Carbocations (C⁺, sp², trigonal planar) are electron-deficient Lewis acids/electrophiles, typically formed via heterolysis, where a bond breaks with both electrons going to one atom. Carbanions (C⁻, sp³, tetrahedral) are electron-rich Lewis bases/nucleophiles, formed when bond electrons stay with carbon; homolysis forms radicals.
Acid-base equilibrium is described by the acid dissociation constant ( Ka ) and \text{p}Ka ( -\log Ka ); a smaller \text{p}Ka indicates a stronger acid. Equilibrium favors the formation of the weaker acid and weaker base. These principles apply to practical applications like salt formation to increase solubility and separatory funnel techniques for purification. Acidity also increases down a group in the periodic table (e.g., HF < HCl < HBr < HI) due to increasing bond length and weakening bond strength, making proton donation easier.