Base (chemistry) Study Guide

📖 Core Concepts Arrhenius base – dissociates in water to give OH⁻ (e.g., NaOH, Ca(OH)₂). Brønsted–Lowry base – accepts a proton (H⁺); OH⁻ and NH₃ are classic examples. Lewis base – donates an electron pair to an electron‑acceptor; does not need H⁺ (e.g., CO₃²⁻, NH₃). $Kb$ – equilibrium constant for the reaction B + H₂O ⇌ BH⁺ + OH⁻; larger $Kb$ → stronger base. Strong base – fully protonated by water (quantitative conversion to BH⁺). Weak base – only partially protonated; $Kb$ ≈ $1.8 × 10^{-5}$ for NH₃ at 25 °C. Superbase – stronger than OH⁻; cannot exist in aqueous solution. Monoprotic vs. polyprotic bases – number of OH⁻ released per formula unit (e.g., NaOH = 1, Ca(OH)₂ = 2). Leveling effect – in water, all bases stronger than OH⁻ appear as OH⁻ (i.e., they are “leveled” to a strong base). --- 📌 Must Remember Arrhenius: base → OH⁻ in water → pH > 7. Brønsted: base = proton acceptor; conjugate acid = BH⁺. Lewis: base = electron‑pair donor (no proton needed). Strong bases: alkali‑metal hydroxides, alkaline‑earth hydroxides (Ca(OH)₂, Ba(OH)₂). Weak base example: NH₃, $Kb = 1.8 × 10^{-5}$ (25 °C). Indicators: litmus blue, phenolphthalein pink, bromothymol blue stays blue, methyl orange turns yellow. pH of a basic solution: > 7. Superbases: cannot be in water; they deprotonate water completely. Neutralization: acid + base → H₂O + salt; strong acid + strong base → complete neutralization. --- 🔄 Key Processes Base + Water (Bronsted–Lowry) \[ \text{B} + \text{H}2\text{O} \;\rightleftharpoons\; \text{BH}^+ + \text{OH}^- \quad (Kb) \] Neutralization Acid dissociates: $\text{HA} \rightarrow \text{H}3\text{O}^+ + \text{A}^-$. Base dissociates: $\text{BOH} \rightarrow \text{B}^+ + \text{OH}^-$. $\text{H}3\text{O}^+ + \text{OH}^- \rightarrow \text{H}2\text{O}$; remaining ions form the salt. Non‑hydroxide base neutralization (e.g., $\text{NH}3$, $\text{CO}3^{2-}$) Accepts a proton from the acid → forms conjugate acid (NH₄⁺, HCO₃⁻) → proceeds to water formation. Lewis base coordination Electron‑pair donor (L) → metal center (Lewis acid) → MLₙ complex formation. --- 🔍 Key Comparisons Arrhenius vs. Brønsted vs. Lewis Arrhenius: must generate OH⁻ in water. Brønsted: must accept H⁺ (includes OH⁻, NH₃, CO₃²⁻). Lewis: must donate an electron pair (includes all Brønsted bases and species like CO₃²⁻ that never donate H⁺). Strong vs. Weak vs. Superbase Strong: complete ionization, $Kb$ large (practically infinite). Weak: partial ionization, measurable $Kb$ (e.g., $10^{-5}$). Superbase: $Kb$ > that of OH⁻; cannot survive in water. Monoprotic vs. Diprotic (Polyprotic) Bases Monoprotic: 1 OH⁻ per formula (NaOH). Diprotic: 2 OH⁻ per formula (Ca(OH)₂, Ba(OH)₂). --- ⚠️ Common Misunderstandings All bases contain OH⁻ – false; NH₃ and CO₃²⁻ are bases without OH⁻ in the solid. A larger $Kb$ always means a higher pH – $Kb$ reflects strength in dilute solution; concentration also matters. Superbases can be used in aqueous titrations – impossible; they react with water to give OH⁻. Strong base = always “very high” pH (> 12) – in very dilute solutions, pH may be only slightly > 7. Leveling effect means all bases are the same – only in water; in non‑aqueous solvents stronger bases retain their identity. --- 🧠 Mental Models / Intuition “Proton sponge” – think of a base as a molecule with an empty seat (lone pair) eager to grab a proton. Electron‑pair donor – picture a Lewis base as a hand offering its electrons to a metal “grabber.” Leveling effect analogy – water is a “flattening board” that only shows the strongest base (OH⁻); anything stronger is flattened down to OH⁻. --- 🚩 Exceptions & Edge Cases Superbases: cannot exist in water; they deprotonate water completely (e.g., NaH, organolithiums). Carbonate (CO₃²⁻): neutralizes acid despite lacking OH⁻; acts via proton acceptance. Polyprotic bases: each OH⁻ released may have a different $Kb$ (stepwise dissociation). Lewis acids can be neutral molecules (BF₃) or high‑oxidation‑state metals (Fe³⁺). --- 📍 When to Use Which Identify base type → use the appropriate definition: OH⁻ present → Arrhenius is sufficient. Proton acceptance without OH⁻ → Brønsted–Lowry. Electron‑pair donation, no proton involvement → Lewis. Choose $Kb$ vs. $Ka$ → for base‑related equilibria use $Kb$; for conjugate acid equilibria use $Ka$ (remember $Ka · Kb = Kw$). Select indicator → based on expected pH range: phenolphthalein (pH ≈ 8.2–10), bromothymol blue (pH ≈ 6.0–7.6). Use strong base → when you need complete neutralization or to generate a known concentration of OH⁻. Use weak base → for buffer preparation or when only partial neutralization is desired. --- 👀 Patterns to Recognize Formula contains OH⁻ → likely a strong base (unless polyprotic with low solubility). Presence of lone‑pair‑rich atoms (N, O, S) without H⁺ → potential Lewis base. Color change of litmus to blue → basic environment. pH > 7 and indicator pink → phenolphthalein confirms basicity. Reaction of a solid base with water producing a clear solution → strong, fully dissociated base. --- 🗂️ Exam Traps “All bases turn red litmus blue” – true, but methyl orange turns yellow, not pink; mixing indicator choices can mislead. Choosing Na₂CO₃ as a “non‑base” – it is a base (accepts protons) despite lacking OH⁻. Assuming a large $Kb$ always gives pH > 12 – concentration matters; a dilute strong base may give pH ≈ 8. Confusing $Kb$ with $Ka$ – remember they are reciprocals multiplied by $Kw$ ( $Ka · Kb = 1.0 × 10^{-14}$ at 25 °C). Identifying a superbasis as usable in aqueous titration – impossible; they react with water before the intended reaction. ---