Unit 3A – Acid/Base By: Nikola Popovic, Rory Fencl, Jonathan Hooyman

Strong Acids and Bases Strong Acids –HCl, HBr, HI, H2SO 4, HNO 3, HClO 4 Strong Bases –First 2 columns of periodic table + OH -

Example of a typical reaction “molecular” HCl + NaOH  H 2 O + NaCl “total ionic” H + + Cl - + Na + + OH -  Na + + Cl - + H 2 O “net ionic” H + + OH -  H 2 O

General pH rules For a strong acid reacting with a weak base, the pH will be lower then a weak acid with weak base, or strong acid with strong base For a weak acid reacting with a strong base, the pH will be higher then a weak acid with weak base, or strong acid with strong base

Acid/Base Definitions Bronsted – Lowery –Acid: proton (H + ) donor –Base: proton acceptor According to Lewis Model –Acid: electron acceptor –Base: electron donor Lewis acids and bases do not always involve H+ –Ex. When BF 3 and NH 3 react, the electron rich NH 3 donates an electron pair to BF 3 to form a stable bonding interaction. BF 3 is the Lewis acid, and NH 3 is the Lewis base

Example Bronsted-Lowery H 2 O + H 2 O  OH - + H 3 O + acid base conjugate conjugate base acid

General pH rules part II pH scale –1-6 = acid –7 = neutral –8-14 = base pH = -log[H + ] pOH = -log[OH - ] 14 – pH = pOH

Formulas Equilibrium Equation  N H * V A * C A = N OH * V B * C B