Chemical Equilibrium

The Equilibrium Constant Chemical Equilibrium Recall that the chemical and physical properties of analytes are invoked for Chemical Analysis. The Chemical indicates the importance of EQUILIBRIA. Classes of Equilibria (Keq): ACID-BASE (Ka & Kb); COMPLEXATION (Kf ); SOLUBILITY PRODUCT (Ksp); PHASE DISTRIBUTION (KD); Etc. Use Molar concentrations If gases use “bar” Solids are omitted

Limestone helps lakes with This problem CaCO3(s) is a base Aside – Acid Rain Result of nitrogen and sulfur oxides from fuel consumption e.g., NO2(g) + H2O(l) + O2(g)   HNO3 (aq) “nitric acid” Your book gives a similar example on page 99.

Manipulation of Equilibrium Constants Problems – will work these at board Consider a reaction with an Equilibrium Constant, K. If the reaction is written in the OPPOSITE direction (right to left), the Equilibrium Constant is 1/K. If Reaction 1 (K1) and Reaction 2 (K2) are ADDED, the sum of the two Reactions (K3) is: K3 = K1 X K2.

Thermodynamics and Equilibria What does it mean if I write Q instead of K? (for a given reaction, the equations look the same) When I write Q,I am saying that the system is NOT at Equilibrium. To calculate Q, I write in the appropriate concentrations at a PARTICULAR POINT IN TIME. If I wait for a while, the reaction will continue until I get a value of K. Then I am at EQUILIBRIUM!

For SPONTANEOUS reactions, DG is NEGATIVE. NEGATIVEDH favors a reaction – EXOthermic POSITIVEDH – ENDOthermic More entropy (disorder) FAVORS a reaction. Temperature effects depend on the ENTROPY term. If DG0 is NEGATIVE, K is greater than 1. DG0 implies that all reaction components are in the STANDARD State. What is DG at equilibrium??

Le Châtelier’s Principle • When a STRESS is applied to a system at equilibrium the system • shifts to REDUCE the stress (e.g., why does one ice skate with so little friction?) • Of course, we are interested in how the Principle • is applied to traditional equilibria • Add something to one side of reaction – shift AWAY • Take something from one side of reaction – shift TOWARD • Heat used as READTANT or PRODUCT to determine TEMPERATURE effects • COMMON ION effect is really a matter of Le Châtelier’s Principle. • How will adding a common ion effect the concentration of the • other ions in an equilibrium? • E.g., what effect does pH have on the solubility of AgCN • For AgCN, Ksp ~ 10-16 • For HCN, Ka ~ 10-9 • We will work this at the board

Separation of Ions in Solution by Precipitation Example: Ions will precipitate in order of the Ksp values. Ag+ added to 0.1 M Cl-, Br-, I- the first to precipitate is I- KspAgI = 8x10-17; KspAgBr = 5x10-13; KspAgCl = 2x10-10 If Ksp values differ enough the I- concentration will be negligible at the point that the Br- begins to precipitate Problem 6-24: What final concentration of carbonate must Exist to precipitate 99.9% of a 0.10 M Zn2+ soln? Ksp = 1.0 x 10 -10 for ZnCO3 ZnCO3(s)  Zn2+ + CO32- Ksp = [Zn2+] [CO32-] we need zinc conc 0.001 (0.10) M 1.0 x 10-10 = (1.0 x 10-4) [CO32-] [CO32-] = 1.0 x 10-6 M

Bronsted-Lowery Definition: Acids are proton DONORS and bases are proton ACCEPTORS in acid base reactions The reactions are NEUTRALIZATIONS that often form a SALT. HCl + NH3 NH4+Cl- Here is another B-L rxn of hydrochloric acid HCl + H20  Cl- + Acid-Base Basics The products of this reaction are also “CONJUGATES”; i.e. NH4+ is the CONJUGATE ACIDof ammonia and Cl - is the CONJUGATE BASE of hydrochloric acid. What can be said about the strengths of these conjugates? Hydronium ion pH is really -log[H3O+ ]

Acid-Base Basics H20  H+ + OH- Kw = [H+] [OH-] (ok I’m being lazy about Kw ~ 1.0 x 10-14 the hydronium ion thing) pH + pOH = 14 What is the Kb of NH3? (your book only lists Kas)

Know – strong vs weak acids (what it means) and what is a polyprotic acid

TITRATION Basics A TITRATION is a volumetric method of WET chemical analysis using a solution of KNOWN concentration as a TITRANT. The TITRANT must react with the analyte with KNOWN stoichiometry and reasonable KINETICS. why? The process is depicted here 

Titration Basics • Other Terms • STANDARD solution - the titrant (solution concentration known with high accuracy). • EQUIVALENCE point - point in a titration when there are “equivalent” amounts of titrant and analyte • END point - point in a titration when some noticeable signal indicates equivalence point. Note - There can be errors - that is END point and EQUIVALENCE point may not be the same. • Primary STANDARD - highly purified compound that is used as a reference material to “standardize” solutions. Other characteristics are STABLE, READILY AVAILABLE, GOOD SOLUBILITY, and LARGE MOLECULAR MASS. Why? • Normality & Equivalents - we will use molarity and moles only

Titration Basics Shapes (types) of titration curves e.g., pH, pOH, pMetal e.g., colorimetric titration A + B  C

Titration Basics More terms & concepts • An INDICATOR(e.g. phenolphthalein) changes in an observable way near the EQUIVALENCE point. (Acid-base INDICATORS are themselves acids or bases!) • The difference between the end point (what you detect) and the equivalence point (what you WANT to detect) is called the TITRATION error. (Causes of systematic TITRATION error?) • A BLANK titration may allow us to determine and eliminate the titration error. Simply repeat the experiment in the ABSENCE of analyte. • Detecting the END POINT usually depends on a rapid change in CONCENTRATION of the analyte or titrant near the equivalence point. • BACK titration is going past the equivalence point with excess titrant then • titrating the excess (next page K-method problem)

Kjeldahl Nitrogen Analysis -An old technique that is still used to measure Nitrogen in samples & illustrates some titration concepts • - Convert all the nitrogen containing compounds in the sample to AMMONIUM ION. • Digest the organic compounds in the sample (decompose, dissolve) in boiling H2SO4. Nitrogen in the organic molecules is converted to NH4+. • - Titrate to determine the amount of NH4+ in the sample. • You know the amount of sample you used. You determine the amount (moles, mass) of nitrogen in that sample based on the titration to determine NH4+. • To do this, • - 1. Make the solution BASIC to convert NH4+ to NH3. Then DISTILL the NH3 (volatile) into a flask containing a known amount of HCl (more HCl than NH3!): • NH3 + HCl  NH4+ + Cl- • - 2. There will be HCl left over. Titrate the HCl with NaOH. This is a BACK titration. • - 3. If you know how much HCl remained (by titration) after distillation of NH3, you can calculate how much NH3 must have been distilled into the receiver.

Titration Basics Problems Factor Label Method is Often Useful What is given “x” conversion factors = Units needed mol N = Total mol HCl – Excess mol HCl = (5mLx0.033M) - 6.34mLx0.01M) = 0.105 mMol N = Mol OBr- L OBr- What next? 0.105 mMol N x 14 g/Mol _________________________________ x 100% = 15.2 % 0.256 mL x 37.9 mg Pro/mL

As it turns out SPECTATOR ions can actually influence equilibria. All Concentrations are not Created Equal! Chemical Activity ACTIVITY, not molarity concentration should be used in equilibrium expressions! a = g [ ] e. g., for reaction A  B + C K = gc[C] gB[B] /gA[A] Remember our prior discussion about how matrices influence calibration The activity coefficient depends on (m = 0.5 S cizi2) Fortunately, in relatively dilute solutions g one and a ~ [ ] That’s All For Chapter 8 !!

Systematic Treatment of Equilibrium(Chapter 9) • Write out what you want to know. • Write out all the other information you have that is relevant. • Use this information to answer your questions. • Write CHARGE balance equations. • Write MASS balance equations. • Write any relevant EQUILIBRIUM expressions. • (Any balanced reaction involving real species in solution is • valid – although perhaps not significant.) • Use this information to answer your questions. • (You will need as many independent equations as you have • unknowns – fortunately approximations can be made.)

Charge Balance • The total charge in a solution should be 0. • Write an equation to show that all the charged chemical species in solution add up to a total charge of 0. • The sum of the POSITIVE charged species equals the sum of the NEGATIVE charged species. Mass Balance • Write an equation to show that the SUM of the concentrations of the different forms of a chemical you have added to solution add up to concentration of the PARENT compound you added to solution. Flow Chart A Couple of Problems

Some Chp 9 Problems

Calculate the Concentration of Ca2+ in a solution that is saturated in CaF(s) and buffered to pH 3.0. Write relevant solubility and acid dissociation reactions (how many different ions?) CaF2(s)  Ca2+ + 2F- HF  H+ + F- Write the Ksp (4 x 10-11) and Ka (7 x 10-4) expressions Ksp = [Ca2+] [F-]2 Ka = ([H+] [F-]) / [HF] Write the mass balance equation for Calcium [Ca2+] = ½ ([F-] + [HF])

Calculate the Concentration of Ca2+ in a solution that is saturated in CaF(s) and buffered to pH 3.0. [Ca2+] = ½ ([F-] + [HF]) Use the Ka expression to substitute for HF and then the Ksp expression to substitute for F- [Ca2+] = ½([F-] + ([H+][F-])/ Ka) = ½[F-](1 + [H+]/Ka) [Ca2+] = ½ (Ksp/[Ca2+])1/2 (1 + [H+]/Ka) [Ca2+]1.5 = (Ksp/4)1/2 (1 + [H+]/Ka) Solve for [Ca2+]; how does this differ from value if HF were strong acid? [Ca2+]1.5 = (4x10-11 / 4)1/2(1+10-3/7x10-4) = 7.7x10-6 [Ca2+] = 4x10-4 M (This is the solubility “s”) If no Ka 4x10-11 = s (2s)2 s = 2x10-4

Sec. IIa LEARNING OBJECTIVES! From Section IIA you may need the following: 1. Understand Le Chatelier’s and common ion principles 2. Be able to determine the number of independent equations needed to solve for all the solutes in a solution 3. Be able to write those equations for situations above; Keq (Ka, Kb, Ksp, Kf, Kw), charge balance, and mass balance 4. Be able to identify acids, bases, conjugate acids, conjugate bases and calculate pH for such situations 5. Know how to use Kw; for example to determine the K-value of a conjugate 6. Know the difference between equivalence point and end point 7. Be able to combine reactions to yield a new reaction and its equilibrium constant 8. Work Kjeldahl method problems

Chemical Equilibrium