Aqueous Chemistry: Polyprotic Acids and Problem Solving
Explore the world of polyprotic acids in aqueous chemistry, including their conjugate bases, acid structure dependence, and problem-solving techniques. Get ready for exams with practice questions and examples. Understand the complexities of polyprotic acids and how to navigate them effectively.
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Announcements I Second Mastering Set due today (should cover material to do most problems today) Exam #1 Next week on Thursday (9/29) Same format as last year (1 12 point work out problem, 22 multiple choice problems each worth 4 points plus one bonus multiple choice problem) Covers material covered through Thurs. (last topic = buffers)
Announcements II Today s Lecture Polyprotic acids Polyprotic acids conjugate bases Dependence of Kaon acid structure Buffers (Chapter 16)
Chem 1B Aqueous Chemistry Some Practice 1. Which solution will have a greater fraction of ionization? 0.10 M HClO vs. 0.10 M HF Ka(HClO) = 2.9 x 10-8 (did already) 2. An unknown base is dissolved in water so that its initial molarity is 0.050 M. The pH is measured and found to be 10.13. What is its Kbvalue? 3. The Kbfor NH3is 1.76 x 10-5. What is the pH of a solution initially made to 0.10 M NH4Cl? Ka(HF) = 3.5 x 10-4
Chem 1B Aqueous Chemistry Polyprotic Acids Generic Example: H2A has two protons that can be lost through acid reactions (diprotic) Some Examples: H2SO4 (sulfuric first H+loss is strong acid) H2SO3 (sulfurous) H2CO3 (carbonic) H3PO4 (phosphoric triprotic) Reaction of generic diprotic example 1) H2A(aq) H+(aq) + HA-(aq) K = Ka1 2) HA-(aq) H+(aq) + A2-(aq) K = Ka2
Chem 1B Aqueous Chemistry Polyprotic Acids in Problems Solving polyprotic acid problems can be challenging (the concentrations of the products from the first reaction affect the equilibrium in the second reaction) To simplify the problem, we assume the two reactions occur independently (valid if Ka1>> Ka2) Example Problem: calculate [H2CO3], [HCO3-], pH, and [CO32-] for a 1.0 x 10-3M solution of H2CO3
Chem 1B Aqueous Chemistry Polyprotic Acids Salts of While conjugate bases of monoprotic weak acids can only be basic, conjugate bases of polyprotic acids may be acidic or basic Example: from H2CO3(carbonic acid), we have HCO3-and CO32-as conjugate bases (from 1st and then 2ndweak acid reactions) 1) H2CO3 (aq) H+(aq) + HCO3-(aq) K = Ka1 2) HCO3-(aq) H+(aq) + CO32-(aq) K = Ka2 Salts allow us to start in the intermediate or basic form
Chem 1B Aqueous Chemistry Polyprotic Acids Salts of cont. The most basic form (CO32-) can only be basic (it has no H+to lose), while the intermediate form (HCO3-) can react as an acid or as a base Acid reaction: HCO3-(aq) H+(aq) + CO32-(aq) Base reaction: HCO3-(aq) + H2O(l) To determine if the intermediate form is acidic or basic, we must compare K values for the acid and base reactions Acid reaction: K = Ka2= 4.7 x 10-11 Base reaction: K = Kw/Ka1= 2.2 x 10-8> Ka2so basic H2CO3(aq) + OH-(aq)
Chem 1B Aqueous Chemistry Polyprotic Acids Salts of cont. Rank the following salts from most acidic to most basic: KHSO4 Na3PO4 KHCO3 KHC2O4 Acid Ka1 Ka2 Ka3 H2SO4 H3PO4 H2CO3 H2C2O4 >> 1 1.2 x 10-2 7.11 x 10-3 6.32 x 10-8 4.5 x 10-13 4.45 x 10-7 4.69 x 10-11 5.60 x 10-2 5.42 x 10-5
Chem 1B Aqueous Chemistry Molecular Structure Acidity Relationship Acid strength depends on ability for H bond to break and on stability of conjugate base formed More stable conjugate bases means stronger acid For example, what makes ethanol (C2H5OH) neutral while acetic acid (CH3CO2H) is acidic? O- H3 C O H3 C H - H O acetate: stabilized by delocalized electrons ethanol anion (not very stable)
Chem 1B Aqueous Chemistry Molecular Structure Acidity Relationship General Rules for Simpler Structures: Binary Acids: e.g. HCl more electronegative element makes for stronger acid longer (and weaker bond) makes for stronger acid (HCl is stronger than HF due to bond strength) Oxyacids: e.g. HClO2 more oxygens make acid stronger (HClO4 is a strong acid, HClO is a very weak acid) Other effects Adding other electronegative elements often makes conjugate base ions more stable Compound pKa CH3CO2H 4.75 CClH2CO2H 2.86 CCl3CO2H 0.66
Chem 1B Aqueous Chemistry Buffers (Chapter 16) We have discussed some mixtures briefly (e.g. strong acid + weak acid) One particular type of mixture: weak acid + conjugate base (or weak base + conjugate acid) makes a solution called a buffer Buffers are desirable because they keep the pH nearly constant even if an acid or base is added Buffers are very important in biology because many enzymes (protein catalysts) will only work over a narrow pH range
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Example: Determine pH of a mix of 0.010 M HC2H3O2 and 0.025 M Na+C2H3O2- solution (Ka of HC2H3O2 = 1.8 x 10-5)
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Buffer Solutions: Question: Was the ICE Problem set up needed? Answer: No. The assumption of x << [HA], [A-] is valid for all traditional buffers Traditional Buffer Weak acid (3 < pKa < 11) Ratio of weak acid to conjugate base in range 0.1 to 10 mM+ concentration range
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Buffer Solutions: Since ICE not needed, can just use Ka equation Ka = [H+][A-]/[HA] = [H+][A-]o/[HA]o (always valid) (valid for traditional buffer) But log version more common pH = pKa + log([A-]/[HA]) Also known as Henderson-Hasselbalch Equation
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Addition of small amounts of acid to a buffer: Example: let s say we have a buffer made to be 0.050 M NH3 + 0.100 M NH4Cl in 1.00 L Calculate the pH Now lets add 0.005 moles of HCl. What is the new pH?
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Change of pH and Buffer Capacity A buffer is designed to minimize the change in pH from addition of base or acid What is the most effective pH? (show next slide) Can we add too much acid or base to a buffer? Does the absolute concentration of acid/base affect pH? Why is a higher concentration better? Buffer Capacity is the ability of the buffer to absorb acid or base without the pH changing significantly
Chem 1B Aqueous Chemistry Buffers (Chapter 16) Most Effective pH Range NH3 + NH4Cl in 1.00 L Example [NH3] + [NH4+] = 0.150 M Addition of 0.005 moles HCl pH -2.75 -0.09 -0.07 -0.06 -0.07 -0.08 -0.32 [NH3] (M) 0.005 0.030 0.050 0.075 0.100 0.120 0.145 [NH4+] (M) 0.145 0.120 0.100 0.075 0.050 0.030 0.005 pH 7.78 8.64 8.94 9.24 9.55 9.84 10.71 Best region (equal moles of WA and CB
