Chapter 17 Acid-Base Equilibria and Solubility Equilibria Part I Notes
Chapter 17 Acid-Base Equilibria and Solubility Equilibria Part I Notes CHEM 1312
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This 4 page Class Notes was uploaded by Justin Sequerra on Tuesday March 8, 2016. The Class Notes belongs to CHEM 1312 at University of Texas at Dallas taught by Dr. Sibert in Winter 2016. Since its upload, it has received 17 views. For similar materials see General Chemistry II in Chemistry at University of Texas at Dallas.
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Date Created: 03/08/16
CHAPTER 17: ACID-BASE AND SOLUBILITIY EQUILIBRIA GENERAL CHEMISTRY II This general outline is meant as a supplement to the General Chemistry II (1312) course taught at the University of Texas at Dallas and should not be taken as a standalone study guide for the overall curriculum. However, I do hope that this broad summary of the textbook helps you all in becoming successful undergraduate students here at UTD. Justin Sequerra, “Chemistry is the study of matter. But I prefer to see it as the study of change.” – Walter White, Breaking Bad 1 ACID-BASE AND SOLUBILITY EQUILIBRIA 17.1 THE COMMON ION EFFECT Up to now: Have only dealt w/ solutions w/ only 1 solute Now: We get to add another Think of these questions as Le Chatelier’s principle (Ch. 15 Recap: Any change in equilibrium results in the system undergoing a change to reduce the stress) but w/ more focus on the addition of reactants or products One must also understand acidbase properties of salt solutions (Ch. 16 Recap: Dissolve salt completely (ex: NaCl) and look at the ions (Na+ and Cl) and decide how they may affect the solutions equilibrium and pH (if they do)) Weak electrolyte (weak base or acid) dissolves in solution to form ionssalt dissolved later in solution may also have the same ion(s) (Common Ion Effect: Sidebar Definition) Addition of these “common” salts SUPPRESSES the ionization of the weak electrolyte (Shifts the Equilibrium to the left/decreases percent ionization) Example: 1. HF (weak electrolyte) dissolved in solution H+ and F (has an equilibrium since HF is a weak acid) 2. Addition of NaF (salt w/ common F ion) increases the concentration of F (since NaF ionizes completely) shifts Equilibrium to left decrease percent ionization of HF (decrease concentrations of both H+ and added F ions decrease pH) 2 ACID-BASE AND SOLUBILITY EQUILIBRIA 17.2 BUFFER SOLUTIONS Buffers contain either weak acid and conjugate base ( ex: HF and F) or weak base and conjugate acid (ex: NH3 and NH4+) The conjugate acid or base is added in the form of a salt (previous examples: NaF to introduce the F ion, and NH4Cl to introduce the NH4+ ion) They buffer the change in pH (resist change in pH) (example: when the internet is buffering it doesn’t want to change the screen; takes a while to) Buffer solutions must have comparable (nearly equal) concentrations of both the weak acid/base and its conjugate pair; it is why we add a salt (b/c weak acids/bases only ionize only a little to produce their respective conjugate pair) Addition of a SMALL AMOUNT of strong acid into a buffer converts it into a weak acid. Ex: H+ + F HF Add HI (strong acid) dissociates to H+ and I Addition of H+ Shift Equilibrium right converts into weak acid (HF) Same thing w/ strong base addition (strong base weak base) By making strong acids/bases weak, it buffers the pH The added strong acid/base MUST be of a smaller amount compared to conjugate pair in order for the buffer to work (all added strong acid/base consumed by reaction) Ex: [H+] < [F] [OH] < [HF] May either use Equilibrium Table (longer) or Henderson Hasselbalch (sidebar definition) equation to calculate pH Must do STOICHIOMETRY before doing Equilibrium table or HendersonHasselbalch equation If weak acid/ base [ ] = [conjugate pair] pH = pKa Only a buffer if: 10 > [conjugate pair] / [weak acid/base] > 0.1 Buffer Range: Want a specific pH; pH = pKa +/ 1; so pick conjugate pair w/ a pKa w/in 1 pH unit of the desired pH 3 ACID-BASE AND SOLUBILITY EQUILIBRIA *PRACTICE PROBLEMS* Week! Have a Great Spring Break!
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