Applications of Aqueous - General Chemistry and Qualitative Analysis - Lecture Slides, Slides of Chemistry

Download Applications of Aqueous - General Chemistry and Qualitative Analysis - Lecture Slides and more Slides Chemistry in PDF only on Docsity! 1 Chapter 15, Applications of Aqueous Equilibria We will focus on 3 areas: 1) titrations 2) buffers (incl. the Henderson- Hasselbalch Transformation), 3) solubility equilibria. 2 I. Neutralization Reactions A. Strong acid-strong base 1. Let’s start by looking at an example: HCl(aq) + NaOH(aq) W H2O(l) + NaCl(aq) 2. Can you write a net ionic equation for the above? + W 3. Will the above rxn have mostly products or reactants present at equilibrium? Logic? B. Weak acid-strong base 1. Again, let’s start by looking at an example: 5 component is weak enough to ignore. 2. Look at the rxn. of acetic acid with ammonia: CH3COOH(aq) + NH3(aq) W NH4+(aq) + CH3COO!(aq) What rxns. will sum to give us that rxn.? + W + Ka = + W + Kb = + W = CH3COOH(aq) + NH3(aq) W NH4+(aq) + CH3COO!(aq) Kn = Ka x Kb x 6 Kn = 3. Will the system contain mostly products or reactants at equilibrium? 4. Perform a similar analysis with HCN (a weaker acid) as the acid instead of CH3COOH on your own. Try prob. 15.2 (a & d?), p. 590. II. The Common Ion Effect (build up to buffers) A. In Chapter 14 we looked at solns. of pure acid or base. We now consider what happens when you look at mixed systems. 7 1. Consider mixing acetic acid and sodium acetate: CH3COOH(aq) + CH3COO!(aq) W CH3COOH(aq) + CH3COO!(aq) Can we calculate [CH3COOH], [H3O+], & [CH3COO!] at equilibrium? ([Na+] usually not of interest.) 2. Use the same approach we developed in Chapter 14 (Fig. 14.7). Main difference: [CH3COO!]initial … 0. a) Step #1, Identify reactive (interesting) species: CH3COOH Na+ H2O CH3COO! acid inert acid/base base b) Step #2-3, Identify principle reaction: 10 HCN + OH! ÿ H2O + CN! b) What happens if you add a H3O+ to the buffer? CN! + H3O+ ÿ H2O + HCN 2. Buffer capacity. There is a limit to how much acid or base the buffer can absorb. a) The amount of acid that can be absorbed is related to how much basic component (CN! above) of the buffer is present. b) The amount of base that can be absorbed is related to how much acidic component (HCN above) of the buffer is present. 11 E. Where (on the pH scale) does a buffer work? 1. Recall the Ka expression: [H3O+] [A!] Ka = [HA] 2. This can be rearranged to obtain: Ka [HA] [H3O+] = [A!] 3. This tells us: a) The [H3O+] (and therefore pH) is determined by the ratio of acid and conjugate base. b) The pH of effective buffering depends on Ka. Do Key Concept Prob. 15.6, p. 598. 12 IV. Henderson-Hasselbalch Transformation Some concepts are much more clear if you look at them from a specific point of view. H-H Transformation makes some aspects of buffers more clear. This is a transformation because you are just rearranging the Ka expression. 15 1. Buffers are most effective buffering against both H+ & OH! addition when buffer pH = pKa of that HA. 2. Look at the [base] ' [acid] ratios on p. 599. 3. Try Prob. 15.10, p. 601. 4. Look at the Normal Values section of: http://www.nlm.nih.gov/medlineplus/ency/article/003855.htm a) Do any of: pH'PaCO2'PaO2'SaO2' HCO3! relate to variables in the H-H transformation? Which relate to acid-base chemistry? 16 pH pH = pKa + log([A!]'[HA]) PaCO2 PaO2 SaO2 HCO3! b) Is HCO3! acting like an acid or a base? See pKa values above, think of CO2 leaving the body. 17 V. pH Titration Curves A. Titration: quantitative analysis method in chemistry. Based on chemical rxns. To do one, you need to know: 1. The stoichiometry for the reaction. 2. The concentration of the “known” component. 3. The volume of known component added. B. If you know these things, you can calculate the quantity of unknown present in a sample. C. You can also get pKa information from a titration. 20 3. Comment on the shape (vs. pH location) of the Fig. 15.9 curve for different weak acids. If you understand this figure, you are in good shape re. acid-base chemistry and buffers. To reinforce see Key Concept Prob 15.15, p. 607 VIII. Weak Base-Strong Acid Titrations (VII) IX. Polyprotic Acid-Strong Base Titrations A. Analogous to VII, above. Fig. 15.11, p. 610. B. Try Prob. 15.19, p. 612 on your own. 21 X. Solubility Equilibria A. Examples of biological solubility problems: 1. tooth decay 2. Atherosclerosis 3. Kidney stones (calcium oxalate) B. Consider the equilibrium: CaF2 (s) WCa2+ (aq) + 2 F!(aq) Can you write a Ksp expression for this: Ksp = 22 1. Ksp is called the solubility product. 2. Different salts have different (sometimes very different) Ksp values. Qual scheme? Try Prob. 15.20 c), p. 613. XI. Measuring Ksp, Calculating Solubility from Ksp A. Two ways to approach this problem: 1. Add increasing concentrations of components of interest until you see a ppt. (Example?) 2. Form a saturated soln. and then measure concentrations of ions in soln. B. We will examine some of the reasons for 25 Change in acid-conjugate base ratios as a function of pH pH 65432 7 8 9 10 H2O H3O+ OH- 1 1 10 1 100 1 1,000 1 104 1 105 1 1 10 1 100 1 1000 R C O O H + H2O Aspirin, HA form R C O- O Aspirin, A- form H3O++ pKa = 3.0 Aspirin HA A- 1 1 10 1 1 10 1 100 1 1000 104 1 105 1 H2CO3 pKa: 3 HA A- pKa: 6.37 + H2O Carbonic acid, HA form Carbonic acid, A- form H3O++ pKa = 6.37 C O OO HH C O O-O H Final thoughts on acid-base/buffer items: If the HA form of aspirin crosses readily membranes, where in your g.i. tract will aspirin be absorbed? 26 Change in protein activity as a function of pH pH 65432 7 8 9 10 + H2OHProtein H3O++ pKa = 6.0 HProtein 1000 1 100 1 10 1 1 1 10 1 100 1 Most proteins have acid/base groups. Assume here that a protein has no activity in its HProtein f orm & is 100% active in the Protein- form. Would the activity be higher at pH 7.4 or 6.0??? What would a graph of activity vs. pH look like? Protein- Protein- inactive active 1000 1 Can we start to consider the molecular basis for respiratory acidosis? 27 ac tiv ity (% ) 20 40 60 80 100 pH 642 8 10 pH Activity Relationship Draw the activity vs. pH graph below.