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Chapter16IM

# Chapter16IM - Chapter 16 Acid-Base Equilibria and...

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Chapter 16 Acid-Base Equilibria and Solubility Equilibria This is the third of the three-chapter sequence which discusses equilibrium concepts. Upon completion of this chapter, your students should be able to: 1. Describe the common ion effect as a special case of Le Châtelier’s principle. 2. Formulate the Henderson-Hasselbalch equation. 3. Use the Henderson-Hasselbalch equation to determine the pH of a solution containing a weak acid (weak base) and its conjugate base (conjugate acid). 4. Describe a buffer solution and its importance in chemical and biological systems. 5. Calculate the pH of a buffer solution using the Henderson-Hasselbalch equation. 6. Calculate the pH of a buffer solution after the addition of H + or OH - . 7. Describe how to prepare a buffer of a desired pH. 8. Predict the pH profile of a strong acid-strong base titration and calculate the pH at any stage of the titration. 9. Predict the pH profile of a strong acid-weak base (or strong base-weak acid) titration and calculate the pH at any stage. 10. Distinguish between end point and equivalence point of a titration. 11. Describe common acid-base indicators and name the correct method of selection for a specific titration. 12. Use the concepts of equilibrium to relate ion product, Q, with K sp to predict if a solution is unsaturated, saturated, or supersaturated. 13. Explain the solubility of a substance in terms of molar solubility and solubility in terms of mass. 14. Use the concept of fractional precipitation to predict concentration of insoluble ions in a solution. 15. Calculate the solubility of an insoluble ion when a common ion is present. 16. Describe how changing pH can affect solubility. 17. Use the concepts of equilibrium, formation constant (K f ) and complex ion formation to predict solubility and ion concentration on solubility.

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18. Describe how solubility product principle is used in qualitative analysis. Section 16.1 Homogeneous versus Heterogeneous Solution Equilibria This chapter is split into two main topics: homogeneous equilibria and heterogeneous equilibria. Examples of heterogeneous equilibria include the solubility of slightly soluble salts such as AgCl and BaSO 4 . Section 16.2 The Common Ion Effect Le Châtelier’s principle states that a system at equilibrium will shift to relieve any stress that is applied to the system. For a weak acid HA ↔ H + + A the amount of H + present in an aqueous solution is dependent upon the K a value for the specific acid. In many cases we can correctly assume that [H + ] << [HA] or that the percent of ionization is very small. If the stress of adding the common ion A - is placed on the equilibrium system, then the equilibrium will shift in the direction of HA and even less of the weak acid will ionize. Therefore, if equal concentrations of a weak acid and its conjugate base are present in a solution, it is a safe assumption that the weak acid does not ionize and that [HA] = [A ] The Henderson-Hasselbalch equation a [conjugate base] pH pK log [acid] = + is easy to derive and is often referred to as the buffer equation. It should be noted that
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