Lecture03

# Lecture03 - Autoprotolysis 2H2O currentpoint currentpoint...

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Autoprotolysis : 2H 2 O H 3 O + + OH - "Autoprotolysis Constant" K w = 1.0 x 10 -14 M 2 = [H 3 O + ][OH - ] (25 °C) In pure water, [H 3 O + ] = [OH - ], so: 1.0 x 10 -14 M 2 = [H 3 O + ] 2 [H 3 O + ] = 1.0 x 10 -7 M = [OH] - (25 °C) pH = -log[H 3 O + ] = -log(1.0 x 10 -7 ) = 7.00 (Neutral)

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Add Acid to Water : e.g., [H 3 O + ] increases to 1.0 x 10 -3 M. pH = -log(1.0 x 10 -3 ) = 3.00 In general (at room temperature): Acidic solutions have [H 3 O + ] > 10 -7 M and so pH is < 7 Basic solutions have [H 3 O + ] < 10 -7 M and so pH is > 7
Strong Acids : Dissociate completely into A - and H 3 O + . [H 3 O + ] = C HA HX (X = Cl, Br, I), HNO 3 , HClO 4 , H 2 SO 4 (1 st ) Strong Bases : Dissociate completely into B + or BH + and OH - . [OH - ] = C B MOH (M = Li, Na, K), Ba(OH) 2 (both) Weak Acids and Bases : Most others [H 3 O + ] << C HA (names end in “acid”) e.g. Acetic Acid 0.1 M HOAc --> 0.0013 M H 3 O + [OH - ] << C B (names end in "ine") e.g. ammonia, pyridine

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pH vs pOH : Example - In 0.020 M HCl (strong acid)
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## This note was uploaded on 04/09/2008 for the course CHE 1316 taught by Professor Gipson during the Spring '08 term at Baylor.

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Lecture03 - Autoprotolysis 2H2O currentpoint currentpoint...

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