Lecture 17 - 2/22/10 What is the pH of a solution...

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2/22/10 1 100 NH 3 (aq) + H 2 O(l) ! NH 4 + (aq) + OH - (aq) What is the pH of a solution containing 0.100 M NH 3 and 0.100 M NH 4 Cl? A) 3.25 B) 2.16 C) 6.50 D) 9.26 E) 7.00 HINT: you don’t have to do a calculation HCOOH(aq) + H 2 O(l) ! HCOO - (aq) + H 3 O + (aq) i) Weak base and conjugate acid ii) Weak acid and conjugate base (last lecture) Buffers resist change in pH when acid or base is added Examples of buffers 101 Step 3. Substitute into expression for K b and solve for x. Step 2. Make up I.C.E. table for dominant reaction. I C E 0.100 M 0.100 M 0.00 M +x - x + x 0.100 - x 0.100 +x + x x = 1.8 ! 10 -5 = [OH - ] pOH=-log[OH - ]= 4.74 Step 1. Relevant reactions. K b = 1.8 ! 10 " 5 = NH 4 + [ ] OH " [ ] NH 3 [ ] = (0.100 + x)(x) (0.100 - x) ! (0.100)(x) (0.100) NH 3 (aq) + H 2 O(l) ! NH 4 + (aq) + OH - (aq) pH = 14.00 - pOH = 9.26 8.1.1. Calculation of pH of buffer solutions e.g. Calculate pH of buffer solution A (0.250 M HCOOH and 0.100 M HCOONa) We already calculated pH using approximations : [HCOOH] eq = 0.250 – x ! [HCOOH] initial ! 0.250M [HCOO - ] eq = 0.100 + x ! [HCOO - ] initial ! 0.100M 102 pH = 3.35 Henderson-Hasselbalch Equation
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Lecture 17 - 2/22/10 What is the pH of a solution...

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