Additional Aspects of Aqueous Equilibria Objectives-The successful 1C student will: 1. understand the common ion effect. 2. understand buffers and buffer properties; be able to calculate the ph of a buffer solution. 3. know some common applications of buffers. 4. understand titration curves: be able to determine ph of a solution during a titration; be able to use titration curves to determine Ka and Kb values. 5. understand factors that effect solubility of ionic compounds in water. 6. understand and be able to apply Ksp and Kf concepts to solubility of ionic compounds in water; be able to apply these concepts to determine the solubility of ionic compounds in water. 7. understand separation of ions by selective precipitation. 8. know some applications of aqueous equilibria concepts. Questions to answer: 1. How can we control the ph of a solution? 2. How do we determine the ph of a solution when acids and bases are mixed in any proportion? 3. What is a titration curve and what information does it give us? 4. If aqueous solutions of two ionic compounds are mixed, will a precipitation reaction occur? 5. How soluble are insoluble ionic compounds in water? 6. What factors effect solubility of ionic compounds? How do we redissolve a precipitate? How can we separate ions in solution by selective precipitation? Other Aspects of Ionic Equilibria 1 Common Ion Effect The common ion effect occurs when an ion involved in an equilibrium is added to the system from a secondary source. This common ion shifts the equilibrium away from the added ion. Some of this ion is consumed as the equilibrium is shifted. For weak acids and bases, the common ion effect decreases the percent of ionization or hydrolysis. The acetate and hydronium ions are the common ions Example of a weak acid ionization: to the equilibrium. CH 3 COOH + H 2 O CH 3 COO + H 3 O + The forward ionization is limited if a significant amount of acetate ion (or hydronium ion) is already present in solution (Le Chatelier s principle). Other Aspects of Ionic Equilibria 2
Common Ion Effect Example of a weak base hydrolysis and common ion effect: NH 3 + H 2 O NH4 + + OH The forward hydrolysis is limited if a significant amount of ammonium ion (or hydroxide ion) is already present in solution. The ammonium and hydroxide ions are the common ions to the equilibrium. Indicate whether the ph increases, decreases or remains the same when: 1. ammonium nitrate is added to a solution of ammonia 2. NaNO3 is added to a solution of HNO3 Other Aspects of Ionic Equilibria 3 Buffer Solutions - Criteria For a solution to act as a buffer, a significant amount of an acid and base that do not react to neutralize each other must be present in the solution. The acid and base are usually a weak acid and its conjugate base; these are the buffer components. Thus, buffers are a special case of the common ion effect involving acids and bases. The equilibrium established between the buffer components works to stabilize the ph of the solution: HA(aq) + H 2 O(l) A (aq + H3O + (aq) [H 3 O + ] = K a x [HA] [A ] Buffer solutions resist changes in ph when a reasonable amount of strong acid or strong base is added. The ph is stabilized by the conjugate acid/base pair present. The weak acid neutralizes added strong base; the weak acid/conjugate base ratio decreases, but ph is relatively stable (within limits). The weak base neutralizes added strong acid; the weak acid/conjugate base ratio increases, but ph is relatively stable (within limits). Question: Which of the following, when in solution together, will act as a buffer? a) HNO3 and NaNO3 b) H2CO3 and NaHCO3 c) NaHCO3 and Na2CO3 Other Aspects of Ionic Equilibria 4
Buffers: Common Ion Effect When an excess of a weak acid/base reacts with a strong base/acid we can determine the final ph by using stoichiometry and then applying the common ion effect in an equilibrium calculation. The steps to determine the equilibrium state of the reaction are as follows: 1. Write the acid base reaction. 2. Take the reaction forward 100% using the strong species as the limiting reactant. 1. For di- tri-protic acids you may have to react the second or third ionizable hydrogens until all the strong base is consumed. 3. Set up a K a (or K b ) reaction (ICE) table for the remaining weak species in solution. Initial equilibrium concentrations are from the limiting reactant calculation. Calculate the final ph. Question: A solution prepared by mixing 50.0 ml of 0.200 M H 2 C 2 O 4 (What s the name of this acid?) with 15.0 ml of 1.00 M NaOH? Is the resulting solution a buffer? What is the ph of the resulting solution? For H 2 C 2 O 4 K a1 = 5.9x10 2 and K a2 = 6.4x10 5 Other Aspects of Ionic Equilibria 5 IMPORTANT DETAILS! The Henderson-Hasselbalch (H-H) Equation ph = pk a + log [base] [acid] Since both the acid and base are in the same solution, we can use moles (or mmole) instead of concentration in the H-H equation. The ph of a buffer is primarily determined by the weak acid s pk a value. The ratio of base to acid gives a fine control of the ph. The buffer components must have sufficient capacity to neutralize the added strong acid or base. Buffer capacity refers to the ability of the buffer to maintain the ph following addition of strong acid or strong base. The more concentrated the buffer components, the greater the buffer capacity. Do you understand why? (Usually 0.1 M concentrations or above are used for effective buffer capacity.) mol base ph = pk a + log mol acid Remember that [acid] and [base] refer to the INITIAL concentrations. Therefore, this equation is valid ONLY when the percent ionization of the weak acid is negligible (less than 5%). The capacity of a buffer to neutralize added acid versus base depends on the relative concentrations of the buffer components. Do you understand why? Buffers are effective when the base/acid ratio is within the range of 0.1 to 10. Given this, what is a buffer s range, the ph over which the buffer will be effective? Other Aspects of Ionic Equilibria 6
Buffer ph and Capacity The molecular scenes below represent samples of four HA/A - buffers. (HA is blue and green, A - is green, and other ions and water are not shown.) QUESTIONS: Two of these buffers have the same ph. Which two? Do these two buffers have the same buffer capacity? Which buffer has the highest ph? Which buffer will have a greater capacity to neutralize added acid, buffer 2 or buffer 4? Should we add a small amount of concentrated strong acid or strong base to convert buffer 1 to buffer 2 (assuming no volume changes)? Other Aspects of Ionic Equilibria 7 Biological Applications: ph Control in Blood Plasma Buffer solutions are necessary to keep the correct ph for enzymes in many organisms to work. Many enzymes work only under very precise conditions; if the ph moves outside of a narrow range, the enzymes slow or stop working and can denature (structure is disrupted). In many cases denaturation can permanently disable their catalytic activity. Human blood transports dissolved gases, nutrients, and wastes from one location to another within the body. Introduction of these materials into the bloodstream can cause ph changes that interfere with the metabolic activity of blood cells. This interference is minimized by the action of blood buffering systems. The major buffer system that controls blood ph is the carbonic acid/hydrogen carbonate ion system. 2 H 2 O(l) + CO 2 (g) H 2 O + H 2 CO 3 (aq) H3O + (aq) + HCO 3 (aq) QUESTION: ph = pk a + log [HCO 3 ] [H 2 CO 3 ] At physiological temperatures pk a1 for carbonic acid is 6.1. The normal ph for blood plasma is 7.4. What is the ratio of [HCO3 ]/[H2CO3] for this buffer system? What does this ratio tell us about the ability of the buffer system to neutralize excess acid? excess base? Other Aspects of Ionic Equilibria 8
Biological Applications: ph Control in Blood Plasma Death may result when human blood ph falls below 6.8 or rises above 7.8. Respiration via exhalation of CO 2 (g) provides a mechanism for adjusting the carbonic acid/carbonate ion buffer system. In addition, properly functioning kidneys work to maintain the required reservoir of HCO 3. Acidosis: A condition in which the ph of blood decreases. Acidosis can be brought on by heart failure, kidney failure (not enough HCO 3 ), persistent diarrhea, a long-term protein diet, or emphysema/ pneumonia (CO 2 not eliminated sufficiently). Persistent, intense exercise can cause temporary acidosis due to lactic acid build up. Alkalosis: A condition in which the ph of blood increases. Alkalosis can be caused by severe vomiting, hyperventilation (rapid breathing, sometimes caused by anxiety, resulting in too much CO 2 eliminated) or exposure to high altitudes (causes rapid breathing). QUESTIONS: Why does hyperventilation - ridding your body quickly of CO2, change your blood ph? Why does breathing in a paper bag help restore your blood ph to more normal levels? What are some other applications of buffers? (Look them up; available online!) Other Aspects of Ionic Equilibria 9 Buffers and the H-H equation 1. A solution is prepared by mixing 38.0 ml of 0.200 M (CH 3 ) 3 N with 50.0 ml of 0.200 M (CH 3 ) 3 NHCl and then diluting the mixture to a total volume of 100.0 ml. Calculate the ph of the resulting solution. 2. You have to prepare a ph 3.50 buffer, and you have the following 0.10 M solutions available: HCOOH, HC2H3O2, H3PO4, NaHCOO, NaC2H3O2, and NaH2PO4. Which solutions would you use? How many milliliters of each solution would you use to make approximately one liter of the buffer? Other Aspects of Ionic Equilibria 10
Preparing a Buffer - Three Methods 1. Preparing a Buffer - Mix Acid + Conjugate Base A biochemist needs 750 ml of an acetic acid/sodium acetate buffer with ph 4.50. Solid sodium acetate and glacial acetic acid are available. Glacial acetic acid is 99% acetic acid by mass (1% water) and has a density of 1.05 g/ml. If the buffer is to be 0.20 M in acetic acid, how many grams of sodium acetate and how many milliliters of glacial acetic acid must be used? Other Aspects of Ionic Equilibria 11 2. Preparing a Buffer - Mix Excess Weak acid + Strong Base Calculate the ph of a buffer solution prepared by mixing 21.7 ml of 2.00 M NaOH with 9.33 g of sodium bicarbonate in 250.0 ml of solution. (a) Write the net ionic equation for the reaction that occurs when a few drops of nitric acid solution is added to this buffer. Other Aspects of Ionic Equilibria 12
3. Preparing a Buffer - Mix Excess Weak Base + Strong Acid Calculate the ph of a buffer solution prepared by mixing 40.0 ml of 0.444 M HCl(aq) with 2.52 g of sodium hydrogen carbonate. Other Aspects of Ionic Equilibria 13 1. What is the ph of a 0.20 M NaCl solution? Buffer Action Calculations 1.1.Determine the ph change of a 0.20 M NaCl solution when 4.0 ml of 1.0 M HCl is added to 100.0 ml of the solution. 2. Determine the ph of a solution containing 0.20 M NaH 2 PO 4 and 0.20 M Na 2 HPO 4. Other Aspects of Ionic Equilibria 14
Buffer Action Calculations 3. What is the ph change when 4.0 ml of 1.0 M HCl is added to 100.0 ml of a solution that is 0.20 M NaH 2 PO 4 and 0.20 M Na 2 HPO 4? or use H-H Equation or new mmoles HX and X Other Aspects of Ionic Equilibria 15 Buffer Action Calculations 4. Determine the ph change when 9.0 ml of 0.80 M NaOH is then added to the results of problem 3. Other Aspects of Ionic Equilibria 16
Buffer Calculations 1. How many grams of sodium hydroxide must be added to 50.0 ml of a buffer that is 0.100 M in acetic acid and 0.200 M in acetate ion to increase the ph by 0.70 units? Other Aspects of Ionic Equilibria 17