Acid-Base Properties of Salt Solutions: Understanding the Role of Ions, Exercises of Chemistry

Download Acid-Base Properties of Salt Solutions: Understanding the Role of Ions and more Exercises Chemistry in PDF only on Docsity! 1 18.7 Acid-Base Properties of Salt Solutions – The acidity (basicity) of salt solutions depends on the acid-base properties of their ions • Acidic cations – act as weak acids in water – The cations (conjugate acids) of weak bases (NH4 +, CH3NH2 +, …) → act as weak acids – Small, highly charged metal cations (Al3+, Fe3+, Cr3+, Cu2+, …) → act as weak acids • Neutral cations – do not influence the pH – The cations of strong bases (Group I, Ca2+, Sr2+, Ba2+) and metal cations with +1 charge (Ag+, Cu+, …) are extremely weak acids (weaker than H2O) → do not influence the pH • Basic anions – act as weak bases in water – The anions (conjugate bases) of weak acids (F-, CN-, S2-, PO4 3- …) → act as weak bases • Neutral anions – do not influence the pH – The anions (conjugate bases) of strong acids (Cl-, Br-, I-, NO3 -, ClO4 - …) are extremely weak bases (weaker than H2O) → do not influence the pH • Amphoteric anions of polyprotic acids – can act as weak acids or bases in water – Anions with ionizable protons (H2PO4 -, HPO4 2-, HS-, HSO3 -, HSO4 -) → act as either weak acids or weak bases depending on the relative values of their Ka and Kb constants) Salts of neutral cations and neutral anions yield neutral solutions Example: NaCl(s) → Na+ + Cl- (neutral solution) Na+ → neutral cation (cation of a strong base, NaOH) Cl- → neutral anion (anion of a strong acid, HCl) Salts of acidic cations and neutral anions yield acidic solutions Example: NH4Cl(s) → NH4 + + Cl- (acidic solution) NH4 + → acidic cation (cation of a weak base, NH3) Cl- → neutral anion (anion of a strong acid, HCl) NH4 + + H2O ↔ H3O+ + NH3 Example: FeCl3(s) → Fe3+ + 3Cl- (acidic solution) Fe3+ → acidic cation (highly charged, small cation) Cl- → neutral anion (anion of a strong acid, HCl) Fe(H2O)6 3+ + H2O ↔ H3O+ + Fe(H2O)5OH2+ Salts of neutral cations and basic anions yield basic solutions Example: Na2S(s) → 2Na+ + S2- (basic solution) Na+ → neutral cation (cation of a strong base, NaOH) S2- → basic anion (anion of a weak acid, H2S) S2- + H2O ↔ HS- + OH- Example: KF(s) → K+ + F- (basic solution) K+ → neutral cation F- → basic anion 2 Salts of acidic cations and basic anions yield either acidic or basic solutions If Ka of the cation is larger than Kb of the anion, the solution is acidic (cation is a stronger acid) If Ka of the cation is smaller than Kb of the anion, the solution is basic (anion is a stronger base) Example: NH4F(s) → NH4 + + F- NH4 + → acidic cation (cation of a weak base, NH3) F- → basic anion (anion of a weak acid, HF) NH4 + + H2O ↔ H3O+ + NH3 Ka(NH4 +) = 5.7×10-10 F- + H2O ↔ HF + OH- Kb(F-) = 1.5×10-11 Ka(NH4 +) > Kb(F-) ⇒ NH4 + is a stronger acid than F- is a base ⇒ the solution is slightly acidic Salts of neutral cations and amphoteric anions yield either acidic or basic solutions If Ka of the anion is larger than its Kb, the solution is acidic (the anion is a stronger acid) If Ka of the anion is smaller than its Kb, the solution is basic (the anion is a stronger base) Example: Predict whether solutions of KH2PO4 and K2HPO4 are acidic, basic or neutral. KH2PO4 → K+ + H2PO4 - K+ → neutral cation (cation of a strong base, KOH) H2PO4 - → amphoteric anion ??? H2PO4 - + H2O ↔ H3O+ + HPO4 2- Ka(H2PO4 -) H2PO4 - + H2O ↔ H3PO4 + OH- Kb(H2PO4 -) Ka(H2PO4 -) = Ka2(H3PO4) = 6.3×10-8 Kb(H2PO4 -) = Kw/Ka1(H3PO4) = 10-14/7.2×10-3 = 1.4×10-12 Ka(H2PO4 -) >> Kb(H2PO4 -) ⇒ H2PO4 - is a stronger acid than it is a base ⇒ the solution is acidic K2HPO4 → 2K+ + HPO4 2- HPO4 2- → amphoteric anion ??? HPO4 2- + H2O ↔ H3O+ + PO4 3- Ka(HPO4 2-) HPO4 2- + H2O ↔ H2PO4 - + OH- Kb(HPO4 2-) Ka(HPO4 2-) = Ka3(H3PO4) = 4.2×10-13 Kb(HPO4 2-) = Kw/Ka2(H3PO4) = 10-14/6.3×10-8 = 1.6×10-7 Ka(HPO4 2-) << Kb(HPO4 2-) ⇒ HPO4 2- is a stronger base than it is an acid ⇒ the solution is basic 18.8 The Lewis Acid-Base Definition – Acids – electron pair acceptors – Bases – electron pair donors • The Lewis acid-base definition does not require the exchange of a proton (Lewis acids don’t have to have H in their formulas) – Expands the scope of possible acids – H+ itself is a Lewis acid since it accepts an e- pair from a base (H+ + :B ↔ H–B+) ⇒ All B-L acids donate a Lewis acid (H+) • Lewis bases must contain an e- pair to donate • Lewis acids must have a vacant orbital in order to accept the e- pair from the base