Unit 3A – Acid/Base

By: Nikola Popovic, Rory Fencl, Jonathan Hooyman

Strong Acids and Bases

• Strong Acids– HCl, HBr, HI, H2SO4, HNO3, HClO4

• Strong Bases– First 2 columns of periodic table + OH-

Example of a typical reaction

• “molecular” HCl + NaOH H2O + NaCl

• “total ionic” H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O

• “net ionic” H+ + OH- H2O

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 BF3 and NH3 react, the electron rich NH3

donates an electron pair to BF3 to form a stable bonding interaction. BF3 is the Lewis acid, and NH3 is the Lewis base

Example Bronsted-Lowery

• H2O + H2O OH- + H3O+

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 NH * VA * CA = NOH * VB * CB