Chap19 Ionic Equilibria (2)

Chap19 Ionic Equilibria (2) - Chapter 19 Ionic Equilibria...

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Chapter 19 Ionic Equilibria in Aqueous Systems 19-1
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Ionic Equilibria in Aqueous Systems 19.1 Equilibria of Acid-Base Buffers q 19.2 Acid-Base Titration Curves 19.3 Equilibria of Slightly Soluble Ionic Compounds 19.4 Equilibria Involving Complex Ions 19-2
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he Henderson asselbalch Equation The Henderson-Hasselbalch Equation HA (aq ) + H 2 O( l ) A - ( aq ) + H 3 O + ( aq ) K a = [H 3 O + ][A - ] A] [H 3 O + ] = K a x [HA] - [HA] [A ] - g[H + =- g g [HA] log[H 3 O ] log K a log [A - ] pH = p K a + log [base] [acid] 19-3
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uffer Capacity Buffer Capacity The buffer capacity is a measure of the “strength” of the buffer, its ability to maintain the pH following addition of strong acid or base. The greater the concentrations of the buffer components, the greater its capacity to resist pH hanges changes. The closer the component concentrations are to each ther, the reater e buffer capacity. other, the greater the buffer capacity. 19-4
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Figure 19.4 The relation between buffer capacity and pH change. When strong base is added, the pH increases least for the most concentrated buffer. This graph shows the final pH values for four different buffer solutions after 19-5 gp p the addition of strong base.
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Figure 19.6 The color change of the indicator bromthymol blue. pH < 6.0 pH > 7.5 pH = 6.0-7.5 19-6
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cid- ase Titrations Acid Base Titrations In an acid-base titration, the concentration of an acid (or a ase) is determined by neutralizing the acid (or base) with base) is determined by neutralizing the acid (or base) with a solution of base (or acid) of known concentration. The equivalence point of the reaction occurs when the number of moles of OH - added equals the number of moles of H O + originally present, or vice versa. 3 gy p , The end point occurs when the indicator changes color. he indicator should be selected so that its color change occurs at a - The indicator should be selected so that its color change occurs at a pH close to that of the equivalence point. 19-7
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Figure 19.7 Curve for a strong acid–strong base titration. The pH increases gradually when excess base has been The pH rises very rapidly at the equivalence point, added. qp , which occurs at pH = 7.00. The initial pH is low. 19-8
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Calculating the pH during a strong acid–strong base titration Initial pH [H 3 O + ] = [HA] init pH = -log[H 3 O + ] pH before equivalence point itial mol H + initial mol H 3 O = V acid x M acid mol OH - added = V base x M base mol H 3 O + remaining = (mol H 3 O + init ) – (mol OH - added ) [H 3 O + ] = pH = -log[H 3 O + ] mol H 3 O + remaining V acid + V base 19-9
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Figure 19.8 Curve for a weak acid–strong base titration. The pH increases slowly beyond the equivalence he curve rises point. The pH at the equivalent point is > 7.00 due to the action of the conjugate The initial pH is higher than r the strong acid solution. The curve rises gradually in the buffer region. The weak acid and its conjugate ase are both present reaction of the conjugate base with H 2 O.
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This note was uploaded on 12/15/2011 for the course CHM 139 taught by Professor Browning during the Spring '08 term at University of Toronto- Toronto.

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Chap19 Ionic Equilibria (2) - Chapter 19 Ionic Equilibria...

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