Ionic Equilibrium: Acids, Bases & Salt Hydrolysis Why ionic equilibrium of acids, bases, and salts matters In water, many substances split into ions and then set up a balance (equilibrium) between unionized and ionized forms. That balance controls pH, which in turn controls reaction rates, solubility, taste (sour/bitter), corrosion, and even your blood chemistry. NEET loves asking: which theory defines the acid/base here, what is the conjugate pair, how does Ka or Kb decide strength, and will a salt solution turn acidic or basic? pH scale (0–14) with familiar examples. At pH 7 and 25 C , [H⁺] = [OH⁻]. Away from 25 C , the neutral pH shifts because K w changes with temperature. Three acid–base theories: scope and use-cases Arrhenius (H⁺ or OH⁻ in water), Brønsted–Lowry (proton donor/acceptor), and Lewis (electron-pair acceptor/donor) — three lenses to classify acids and bases. Pick the broadest one that fits the situation. Arrhenius Acid: produces H⁺ in water; Base: produces OH⁻ in water Aqueous only HCl (hydrogen chloride, hydrochloric acid) NaOH (sodium hydroxide) Simple and quick for aqueous strong/weak electrolytes Brønsted–Lowry Acid: proton donor; Base: proton acceptor Wider: any solvent where H⁺ transfer occurs CH3COOH (ethanoic/acetic acid) NH3 (ammonia) — accepts H⁺ to give NH4⁺ Lewis Acid: electron-pair acceptor; Base: electron-pair donor Broadest: includes non-proton acids/bases BF3 (boron trifluoride) — accepts e⁻ pair F⁻ (fluoride), NH3 — donate e⁻ pair Covers reactions with no proton transfer (e.g., BF3·NH3 adduct) Model Comparing acid–base theories Theory Definition Scope Example acid Example base What it covers that others don’t Quick pick rule: If you see H⁺ moving, use Brønsted–Lowry. If you see an empty orbital accepting a lone pair (like BF3), use Lewis. If it’s simple acid/base in water giving H⁺ or OH⁻, Arrhenius works. tip Water uses its lone pair to accept a proton from acetic acid, forming hydronium and acetate. The forward and reverse happen until equilibrium. Aqueous, 298 K Curved arrow from water’s lone pair to H of –COOH Lone pair on water attacks the acidic proton of acetic acid; O–H bond forms on water. Curved arrow from O–H bond to O O–H bond of –COOH breaks; electrons go to the carbonyl oxygen, giving acetate. Reversible Dynamic equilibrium between CH3COOH + H2O and H3O⁺ + CH3COO⁻ establishes. Brønsted–Lowry proton transfer (example with acetic acid) ethanoic acid acetic acid Brønsted–Lowry acid (proton donor) oxidane (water) Brønsted–Lowry base (proton acceptor) water conjugate acid of water hydronium oxidanium ethanoate conjugate base of acetic acid acetate gpt-image-2 Mechanism diagram of Brønsted–Lowry proton transfer: HA + H2O -> H3O+ + A-. Show curved arrows from water lone pair to acidic proton and from O–H bond back to A. Label conjugate pairs (HA/A− and H2O/H3O+). Clean 2D vector, red arrows, black atoms, white background. 2026-05-26T17:04:32.728Z Proton-transfer diagram with conjugate-pair labels: HA + H2O H3O⁺ + A⁻; show HA/A⁻ and H2O/H3O⁺ as conjugate pairs. Conjugate acid–base pairs When an acid donates H⁺, what remains is its conjugate base. When a base accepts H⁺, what forms is its conjugate acid. They differ by one proton and are on opposite sides of the equilibrium — like partners in a seesaw. HCl Cl⁻ H2O H3O⁺ CH3COOH (ethanoic acid) CH3COO⁻ (ethanoate) H2O H3O⁺ H2CO3 (carbonic acid) HCO3⁻ (hydrogen carbonate) HCO3⁻ H2CO3 NH4⁺ (ammonium) NH3 (ammonia) NH3 (ammonia) NH4⁺ (ammonium) HSO4⁻ (hydrogen sulfate) SO 4 2- (sulfate) SO 4 2- HSO4⁻ Pair Acid Conjugate base Base Conjugate acid Typical conjugate acid–base pairs (25 examples in water) tip Spot the pair fast: two species that differ by exactly one H⁺ are conjugates. The one with more H is the acid; the one with less H is the base. Ionization of water and the pH scale Even pure water ionizes slightly. This self-ionization sets the reference for pH. At 25 C , [H⁺] = [OH⁻] = 1.0 10 -7 M in pure water, so pH = 7 is neutral at this temperature. Believing that pH 7 is always neutral. pH 7 is neutral only at 25 C . As temperature increases, K w increases, [H⁺] and [OH⁻] both increase, and the neutral pH drops slightly below 7. K w changes if we dilute or concentrate the solution. K w depends only on temperature (for a given T), not on concentrations added. Dilution changes [H⁺] and [OH⁻] but their product remains K w at that T. Weak acids and bases: Ka, Kb, and degree of ionization Strong acids/bases dissociate almost completely in water (e.g., HCl, NaOH). Weak ones establish an equilibrium with water and only partially ionize. Their strength is quantified by equilibrium constants: K a for acids and K b for bases. Example: ethanoic acid (acetic acid, CH3COOH; SMILES: CC(=O)O) has K a = 1.8 10 -5 ( pK a 4.74). Ammonia (NH3; SMILES: N) as a weak base reacts with water: NH3 + H2O NH4⁺ + OH⁻ with K b = 1.8 10 -5 ( pK b 4.74). Strong acid vs weak acid in water: strong acid shows nearly complete ionization; weak acid shows many undissociated molecules with a few ions. Ostwald dilution law for a weak monobasic acid at low degree of ionization ( 1). For a weak monobasic acid (if C [H⁺]). Analogous forms apply to weak bases with K b . 2026-05-26T17:04:34.541Z Ostwald dilution: as concentration (C) decreases, degree of ionization (α) rises following α = K a /C for a weak acid. gpt-image-2 Graph of alpha (0 to ~0.2) vs concentration ( 10 -1 to 10 -5 M, log scale) for a weak acid with Ka fixed. Smooth curve rising as C decreases. Clean axes, neutral textbook colors, vector style. Strength vs concentration: Strength is an intrinsic property ( K a / K b ). Concentration is how much solute you put. A 0.001 M HCl is still a strong acid — just dilute. remember Common weak acids — pK a (25 ) and approximate K a Acid pK a1 pK a2 pK a3 K a1 (approx.) K a2 (approx.) K a3 (approx.) Acid HF (hydrofluoric acid) 3.17 6.8 10 -4 HCOOH (methanoic/formic acid) 3.75 1.8 10 -4 CH3COOH (ethanoic/acetic acid) 4.74 1.8 10 -5 HCN (hydrocyanic acid) 9.21 6.2 10 -10 H2S (hydrogen sulfide) 7.00 12.92 1.0 10 -7 1.2 10 -13 H2CO3 (carbonic acid) 6.35 10.33 4.5 10 -7 4.7 10 -11 H3PO4 (phosphoric acid) 2.15 7.20 12.38 7.1 10 -3 6.3 10 -8 4.2 10 -13 Salt hydrolysis: will the solution be acidic, basic, or neutral? A salt is made from an acid and a base. In water, its ions may react with water (hydrolyze). The conjugates decide the pH direction: - Strong acid + strong base salt (e.g., NaCl): both ions are too weak to hydrolyze → neutral. - Weak acid + strong base salt (e.g., CH3COONa): anion (CH3COO⁻) is a base → basic. - Strong acid + weak base salt (e.g., NH4Cl): cation (NH4⁺) is an acid → acidic. - Weak acid + weak base salt (e.g., CH3COONH4): depends on K b of anion vs K a of cation. If K b (anion) > K a (cation) → basic; if K b < K a → acidic; if comparable → near neutral. Flowchart: Start with 'Salt AB from acid HA and base BOH'. Branch by strong/weak of parent acid/base. Leaves show pH: neutral/basic/acidic with examples (NaCl, CH3COONa, NH4Cl, NH4CN). Minimalist vector, green=basic, red=acidic, gray=neutral. gpt-image-2 Decision tree for salt hydrolysis: identify parent acid/base strengths → check which ion hydrolyzes → predict pH direction. 2026-05-26T17:04:34.591Z Salt Salt From acid From base Hydrolysis pH at 25 Salt hydrolysis examples NaCl HCl (strong) NaOH (strong) Negligible Neutral (≈ 7) CH3COONa (sodium ethanoate) CH3COOH (weak) NaOH (strong) Anion hydrolyzes (basic) > 7 NH4Cl (ammonium chloride) HCl (strong) NH3/NH4OH (weak) Cation hydrolyzes (acidic) < 7 (≈ 5) CH3COONH4 (ammonium ethanoate) CH3COOH (weak) NH3/NH4OH (weak) Both ions weak; compare K b vs K a ≈ neutral NH4CN (ammonium cyanide) HCN (weak) NH3/NH4OH (weak) Both ions; K b (CN⁻) > K a (NH4⁺) Basic (> 7) Strong acid + strong base → salt + water (no hydrolysis; neutral solution). Classic neutralization sits outside hydrolysis because conjugates are very weak and do not affect pH. NEET trap: A salt from a strong acid and weak base is NOT neutral. NH4Cl solution is acidic because NH4⁺ donates H⁺ to water. neet-alert Strong acid + weak base salt is neutral. False. The cation (conjugate acid of the weak base) hydrolyzes, releasing H⁺, making the solution acidic. Example: NH4Cl gives acidic solution. Only strong acids (HCl, HBr, HI, HClO4, HNO3, H2SO4 (first step)) are nearly fully ionized in water. Weak acids (e.g., CH3COOH, HF) are only partially ionized at equilibrium. All acids ionize completely. Strength equals concentration. A strong acid is strong even if very dilute because it dissociates fully. A weak acid remains weak even if concentrated because its K a is small. Real-life link: Blood pH stays in a tight 7.35–7.45 range via the H2CO3/HCO3⁻ buffer. Disturbances cause acidosis/alkalosis. Hospitals use Henderson–Hasselbalch to quantify buffers (you will study buffer pH in the next concept). remember Henderson–Hasselbalch equation for buffer solutions. Detailed buffer pH calculations are covered in NTCH06/04. Key terms at a glance Acid produces H⁺ in water; base produces OH⁻ in water. Classical aqueous acid–base Arrhenius acid/base Acid donates a proton; base accepts a proton. Protonic acid–base Brønsted–Lowry acid/base Acid accepts an electron pair; base donates an electron pair. Lewis acid/base Two species differing by exactly one H⁺ (acid has the extra H). Conjugate acid–base pair Acid dissociation constant: K a = [H⁺][A⁻]/[HA]. K a Base dissociation constant: K b = [BH⁺][OH⁻]/[B]. K b K w = [H⁺][OH⁻]; equals 1.0 10 -14 at 25 C . K w (ionic product of water) Self-ionization of a solvent (e.g., 2 H2O H3O⁺ + OH⁻). Autoionization Fraction of molecules that ionize; for weak acids at low α, α = K a /C . Degree of ionization (α) Relates α to concentration C for weak electrolytes; α increases on dilution. Ostwald dilution law Reaction of salt ions with water affecting pH (acidic/basic/neutral). Salt hydrolysis Strong vs weak electrolytes (fast checklist) Strong acids: HCl, HBr, HI, HClO4, HNO3, H2SO4 (first step) → strong electrolytes. Strong bases: Group 1 hydroxides (NaOH, KOH), Ba(OH)2, Sr(OH)2 → strong electrolytes. Weak acids: CH3COOH, HF, HCN, H2CO3 → weak electrolytes. Weak bases: NH3, amines → weak electrolytes. Strong acid + strong base salts (e.g., NaCl) → neutral; otherwise consider hydrolysis. clinical Healthcare connects here: arterial blood pH is 7.35–7.45. The H2CO3/HCO3⁻ system resists pH change when metabolic acids/bases load the blood. This is ionic equilibrium at work.