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Acids & bases Basic Chemistry

Definition of the pH in H2O Water molecules dissociate to give hydronium (H3O+) hydroxide (OH-) and ions: (autoprotolysis or self-ionization) 2H2O H3O+ + OH- In neutral water at 25 °C, [H3O+] = [OH-] = 10-7 M. The P operator indicates -log10: pH = -log10([H3O+]) = 7 in neutral water Addition of acid to water leads to pH < 7, ACIDIC Addition of base to water leads to pH > 7, BASIC/ ALKALINE

Importance and applications Some examples: pH of human blood ≈7.35-7.40; pathology indicator Acids and bases are very important in industrial processes: As reactants: Ammonium nitrate production As catalysts: Formation of carbocations in oil refining Control of pH and understanding of acid-base chemistry is important in virtually all scientific research structure design & maintenance, safety…..

Lewis acids & bases Do not necessarily comprise a different group! ….A different feature is emphasized: Lewis base: electron-pair donor; Lewis acid: electron-pair acceptor Example:

Typical acids & bases Acids: Hydrogen halides (HF, HCl, HBr, HI) Oxyacids (H2CO3, H3PO4, HNO3, HNO2…) (Lewis acids) High-charge (transition) metals (Al3+ , Cr3+, Fe3+) Bases Alkali and alkaline earth hydroxides (NaOH, KOH, Ca(OH)2…) Alkali and alkaline earth oxides (K2O, CaO, MgO…) Metal hydrides (LiH, NaH, KH….)

Acid dissociation constant Ka; pKa = -log(Ka) The strength of acids and bases is described by their equilibrium constants General acid HA: HA+ H2O H3O+ + A- Equilibrium constant K: 𝐾= 𝑎 H3O+ 𝑎[A−] 𝑎 HA 𝑎[H2O] ≈ H3O+ A− HA H2O …But the concentration of water is almost constant: 𝐾 𝑎 =K H2O = H3O+ [A−] HA Acid dissociation constant Ka; pKa = -log(Ka)

General base B: B + H2O BH+ + OH- 𝐾 𝑏 = BH+ [OH−] B The strength of acids and bases is described by their equilibrium constants General base B: B + H2O BH+ + OH- 𝐾 𝑏 = BH+ [OH−] B Base dissociation constant Kb; pKb = -log(Kb)

Acid-base conjugate pairs HA/A- and B/BH+ are both acid/base conjugates: HA + B A- + BH+ Acid Base Conjugate base Conjugate acid Reactions always proceed in the direction of the weaker acid and base! pKa + pKb = pKw ; The conjugate base of a strong acid is a very weak base (non-reactive); The conjugate base of a weak acid is a weak base

Acid-base conjugate pairs

While-> HF < HCl < HBr < HI? Acid strength: Why? NH3 < H2O < HF ? Electronegativity: N (3.04) < O (3.44) < F (3.98) While-> HF < HCl < HBr < HI? Bond dissociation energies (kJ/mol): HI (287) < HBr (354) < HCl (419) < HF (543)

pH calculation in aqueous solutions Strong acids (pKa < 0) dissociate completely: [H3O+] = CHA, pH = -log(CHA) Weak acids dissociate negligibly, so [HA] ≈ CHA 𝐾 𝑎 = H3O+ [A−] HA [H3O+] = [A-] = X X2 = Ka[Ha] ≈ KaCHA X = [H3O+] = 𝐾 𝑎 𝐶 𝐻𝐴 pH = −log 𝐾 𝑎 𝐶 𝐻𝐴 = 1 2 (𝐾 𝑎 −𝑙𝑜𝑔 (𝐶 𝐻𝐴 ))

Summary Reactions proceed to yield the weaker acid and base Acids are proton donors or electron-pair acceptors Bases are proton acceptors or electron-pair donors pKa of acids and pKb of bases describes their strength; the lower pKa (pKb) the stronger the acid (base) The strength of a conjugate base is easily deduced from the strength of the acid: pKa + pKb = pKw Reactions proceed to yield the weaker acid and base The pH of a strong acid in aqueous solution: pH ≈ -log(CHA) The pH of a weak acid in aqueous solution: pH ≈ 1 2 (𝐾 𝑎 −𝑙𝑜𝑔 (𝐶 𝐻𝐴 ))

Let’s look at Henry’s law… The concentration of a gas dissolved in any liquid is proportional to it pressure in the atmosphere/gas in contact with the liquid: C = KP, with K the constant, C the concentration in the liquid, and P the partial pressure Henry’s constants for water at 25 °C Gas K (mol L-1 atm-1) CO2­­­­­­­­­­­­ 3.4×10-2 N2 7×10-4 He 3.7×10-4 But note: 𝐻 2 𝐶 𝑂 3 (𝑎𝑞) 𝐶 𝑂 2(𝑎𝑞) =1.8 ×10 −3 Finally: for H2CO3, Ka1 = 2.5×10-4; Ka2 =4.7×10-11