Additional Aspects of Aqueous Equilibria
Objectives-The successful 1C student will:
1.
2.
3.
4.
5.
6.
7.
•
understand the common ion effect.
understand buffers and buffer properties; be able to calculate the pH of a buffer solution.
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.
understand factors that effect solubility of ionic compounds in water.
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.
understand separation of ions by selective precipitation.
learn about some applications of aqueous equilibria concepts.
Questions to answer:
1.
2.
How can we control the pH of a solution?
How do we determine the pH of a solution when acids and bases are mixed in any
proportion?
3.
4.
5.
6.
What is a titration curve and what information does it give us?
If aqueous solutions of two ionic compounds are mixed, will a precipitation reaction occur?
How soluble are “insoluble” ionic compounds in water?
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
1.
2.
3.
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.
Example of a weak acid ionization:
These are the “common”
ions of the equilibrium
CH3COOH + H2O ⇌ CH3COO– + H3O+
The forward ionization is limited if a significant amount of acetate ion (or hydronium
ion) is already present in solution (Le Chatelier’s principle).
The acetate and hydronium ions are the “common” ions to the equilibrium.
Other Aspects of Ionic Equilibria
2
Common Ion Effect
Example of a weak base hydrolysis and common ion effect:
NH3 + H2O ⇌ 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. Ca(C2H3O2)2 is added to a solution of CH3COOH
2. ammonium nitrate is added to a solution of ammonia
3. NaNO3 is added to a solution of HNO3
Other Aspects of Ionic Equilibria
3
Calculations with & without a Common Ion Present
1. Lactic acid (CH3CH(OH)COOH or HC3H5O3, Ka = 1.4x10–4) is an important biomolecule. Muscle
physiologists study its accumulation during exercise. Food chemists study its occurrence in sour milk
products, fruit, beer, wine, and tomatoes. Industrial microbiologists study its formation by various
bacterial species from carbohydrates. Calculate the pH and percent ionization of 100.0 mL of 0.085 M
lactic acid. No common ion present.
2. Repeat the above calculation assuming 1.00 g of sodium lactate is also dissolved in the solution. What
was the effect of adding the common ion?
Other Aspects of Ionic Equilibria
4
Calculations with a Common Ion Present
1. A solution is prepared by mixing 38.0 mL of 0.200 M (CH3)3N with 50.0 mL of 0.200 M
(CH3)3NHCl and then diluting the mixture to a total volume of 100.0 mL. Calculate the pH
of the resulting solution.
Other Aspects of Ionic Equilibria
5
Prelude to Buffers: Weak + Strong Acid/Base Rxns
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.
Here are the steps to determine the equilibrium state of the reaction.
1. Write the acid base reaction.
2. Take the reaction forward 100% using the strong species as the limiting reactant.
a) 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 Ka (or Kb) reaction (ICE) table for the remaining weak species in solution. Initial
equilibrium concentrations are from the limiting reactant calculation. Calculate the final
pH.
Other Aspects of Ionic Equilibria
6
Acid-Base Reactions: Weak + Strong
1. What is the pH of a solution prepared by mixing 50.0 mL of 0.200 M H2C2O4 (oxalic acid) with 15.0 mL
of 1.00 M NaOH? For H2C2O4 Ka1 = 5.9x10–2 and Ka2 = 6.4x10–5
Other Aspects of Ionic Equilibria
7
Buffer Solutions - Criteria
Buffers are a special case of the common ion effect involving acids and bases.
Requirements for a buffer solution:
1.
2.
3.
An acid and base that do not react to neutralize each other must be present in
solution. The acid and base are usually a weak acid and its conjugate base; these
are the buffer components.
The concentration ratio of the buffer components should be within a 10/1 (or 1/10)
ratio. (This is directly related to the effective pH range of the buffer.)
The buffer components must have sufficient capacity to neutralize the added strong
acid or base. Usually 0.1 M concentrations or above are used.
A Ka (or Kb) equilibrium is established between the buffer components.
Buffer solutions resist changes in pH even when a strong acid/base is added.
The pH is stabilized by the conjugate acid/base pair present.
1.
2.
The weak acid neutralizes any strong base; the weak acid/weak base ratio
decreases slightly, but pH changes little.
The weak base neutralizes any strong acid; the weak acid/weak base ratio
increases slightly, but pH changes little.
Other Aspects of Ionic Equilibria
8
Buffer Solutions
Importance: Many reaction systems depend on specific pH range to function properly.
Our bodies’ metabolic activity operating at a BUFFERED pH of 7.4:
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
many different blood buffering systems.
H2CO3/HCO3– buffer (human red blood cell buffer)
H2PO4-/HPO42– buffer (intercellular fluid buffering)
CH3COOH/CH3COO– buffer (blood plasma buffering)
Note: HCHO2 and CHO2— are also commonly written as HCOOH and HCOO—.
Other Aspects of Ionic Equilibria
9
The Henderson-Hasselbalch Equation: Summary
⎛ [base] ⎞
pH = pK a + log ⎜
⎝ [acid] ⎟⎠
1. The pH is primarily determined by the acid pKa.
2. The ratio of base to acid gives a fine control of the pH.
3. The ratio of base to acid should be in the range 0.1 to 10 for an
effective buffer. This means the pH will be within ≈ ±1 unit
of the pKa.
4. Concentrations should be near or above 0.1 M for an effective buffer
capacity.
5. Since both the acid and base are in the same solution, we can use
moles (or mmol) instead of concentration in the H-H equation:
⎛ mol base ⎞
pH = pK a + log ⎜
⎝ mol acid ⎟⎠
Other Aspects of Ionic Equilibria
10
Buffers and the H-H equation
1. Find the pH of a buffer containing 0.152 M ammonia and 0.234 M ammonium chloride.
2. One physiological buffer is the dihydrogen phosphate/hydrogen phosphate system. Find the pH of a
buffer where 2.0 grams of each sodium salt is dissolved in a 100.0 mL solution.
Other Aspects of Ionic Equilibria
11
Buffer Preparation of Specific pH
When preparing a buffer of a desired pH, you need to chose your acid/base pair carefully.
From the H-H equation we see that the pH of a buffer will be near the pKa for the weak acid.
Thus, you should select an acid that has a pKa close (±1 unit) of your desired pH.
⎛ [base] ⎞
pH = pK a + log ⎜
⎝ [acid] ⎟⎠
Text problem 17.31: 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
12
Buffer Action Calculations
1.
What is the pH of a 0.20 M NaCl solution?
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 NaH2PO4 and 0.20 M Na2HPO4.
13
Other Aspects of Ionic Equilibria
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
NaH2PO4 and 0.20 M Na2HPO4?
4.
What is the pH change when 9.0 mL of 0.80 M NaOH is then added to the results of problem 3?
or new
mmoles
HX and X–
Other Aspects of Ionic Equilibria
or use HH
equation
14
Buffer Action 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
15
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
16
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.
(b) Write the net ionic equation for the reaction that occurs when a few drops of potassium
hydroxide solution is added to this buffer.
Other Aspects of Ionic Equilibria
17
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 1.42 g of
sodium fluoride.
Other Aspects of Ionic Equilibria
18
Biological Applications: pH Control in Blood Plasma
The major buffer system that controls blood pH is the carbonic acid/hydrogen carbonate ion system.
⎛ [HCO3– ] ⎞
2 H2O(l) + CO2(g) ⇌ H2O + H2CO3(aq) ⇌ H3O+(aq) + HCO3–(aq) pH = pK a + log ⎜
⎝ [H 2CO3 ] ⎟⎠
The normal pH for blood plasma is 7.4.
1. What is the ratio of [HCO3–]/[H2CO3] for this buffer system?
(At physiological temperatures pKa1 for carbonic acid is 6.1)
2. What does this ratio tell us about the ability of the buffer system to neutralize excess acid? excess
base?
The buffer component proportions reflect the greater need for acid buffering since breakdown of
foods, like proteins and carbohydrates, results in many intermediate substances that are acidic.
Other Aspects of Ionic Equilibria
19
Biological Applications: pH Control in Blood Plasma
Death may result when the pH falls below 6.8 or rises above 7.8. Respiration via exhalation of CO2(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 HCO3–.
Acidosis: A condition in which the pH of blood decreases. Acidosis can be brought on by heart failure,
kidney failure (not enough HCO3–), persistent diarrhea, a long-term protein diet, or emphysema/
pneumonia (CO2 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 CO2 eliminated) or
exposure to high altitudes (causes rapid breathing).
1. Why does hyperventilation - ridding your body quickly of CO2, change your blood pH?
2. Why does breathing in a paper bag help restore your blood pH to more normal levels?
Other Aspects of Ionic Equilibria
20
pH Control in Blood Plasma
The regulation of the pH of blood plasma also relates directly to the effective transport of O2 to body
tissues by hemoglobin. Hemoglobin reversibly binds both H+ and O2 through a series of equilibria
involving protonation-deprotonation and oxygenation-deoxygenation of the hemoglobin (Hb). The overall
equilibria can be approximated as follows:
HbH+(aq) + O2 (aq) ⇌ HbO2(aq) + H+(aq)
1. The concentration of O2 is higher in the lungs and lower in the body tissues. What effect does high [O2]
have on the position of the equilibrium?
2. During strenuous exercise, as O2 is consumed in body tissues, what direction does this equilibrium
shift? What would happen to your blood pH?
a) Exertion also raises the temperature in body tissues. The temperature change shifts this equilibrium
to the left. What is the sign of ∆H for the equilibrium?
b) Lactic acid is produced during strenuous exercise. Which direction does this shift the equilibrium?
3. Exercise also produces CO2 by metabolism. Which direction does this shift the equilibrium above?
Other Aspects of Ionic Equilibria
21