Acids & Bases Calculating pH and Naming

Finding the pH of Solutions Self- ionization of water – the simple dissociation of water H 2 O H + + OH - Concentration of each ion in pure water: [H + ] = 1.0 x M + [OH - ] = 1.0 x M Where K w = 1.0 x Ion-product constant for water (K w ), Where K w = 1.0 x K w = [H + ] [OH - ] Acid [H + ] > [OH - ] Base [H + ] < [OH - ] Neutral [H + ] = [OH - ]

pH Scale pH = -log[H + ] pOH = -log[OH - ] pH + pOH = 14

[OH - ]pOHpH[H + ] 1 x x x x x x x x x Increasing acidity41 x x x x x x Neutral71 x x x x x x Increasing basicity101 x x x x x x x x x

Example If the [H + ] in a solution is 1.0 x M, is the solution acidic, basic or neutral? 1.0 x M What is the concentration of the [OH - ]? Use the ion-product constant for water (K w ): K w = [H + ] [OH - ] 1.0 x = [1.0 x ] [OH - ] 1.0 x = [OH - ] 1.0 x x 10 -(14-5) pH 5 = acidic 1.0 x OH -

Fill in the chart. [OH - ]pOHpH[H + ] 8 1x x × X X X

Naming Acids Binary acids –Contains 2 different elements: H and another –Always has “hydro-” prefix –Root of other element’s name –Ending “-ic” –Examples: HI, H 2 S, HBr

Naming Acids Ternary Acids - Oxyacids –Contains 3 different elements: H, O, and another –No prefix –Name of polyatomic ion –Ending “–ic” for “-ate” and “–ous” for “- ite” –Examples: HClO 4, H 3 PO 4, HNO 2

Practice H 2 SO 3 –Sulfurous acid HF –Hydrofluoric acid H 2 Se –Hydroselenic acid Perchloric acid –HClO 4 Carbonic acid –H 2 CO 3 Hydrobromic acid –HBr