Equilibria: Acids and Bases
Group Work.6
Zen Moment
In silence contemplate the dynamics of dry ice in water.
During the first minute, allow your mind to go where it wants.
During the second minute, become aware-think-about what the water/dry ice solution-is.
Topic: Pure Water. Is it most like Lemons, Doctors Without Borders, or Antacid?
A. What is the pH of the distilled water that you use in lab?
7.0
B. The pH of neutral water is 7. What is the concentration of H3O+ in the flask? What is the pOH?
[H3O+] = 1x10-7 M pOH = 14-7 = 7
C. Each group has an Erlenmeyer flask with water and an indicator (can you tell which one?) Add a small piece ( ~ size of a nickel) of dry ice to the water and swirl. Given that the color change for bromothymol blue indicator occurs between pH = 6.0 to 7.6, what do you estimate the pH of the solution is after the addition of dry ice?
5.5 (< 7.0)
D. Write down the net reaction of CO2(aq) dissolved in water. (hint: recall the geochemistry group work with acid rain and marble) Look at the product of this reaction. Could it be an acid or a base? Write an equation to show its acid/base reaction. How can these two-coupled reactions be used to corroborate your answer to part C?
CO2(aq) + H2O(l) ó H2CO3(aq) Kform = 3.16 x 10-2
H2CO3(aq) + H2O(l) ó HCO3- (aq) + H3O+ Ka1 = 4.2 x 10-7
(Ignore second deprotonation of carbonic acid Ka2 = 10-11 << 4.2 x 10-7)
E. Write down the net reaction of the two-coupled reactions.
CO2(aq) + 2H2O(l) ó HCO3- (aq) + H3O+
F. If you added a 1 g chunk of CO2 (s) to 100 mL of water, what is the pH of the solution? The pKa of carbonic acid is 7.88. (hint: This is a weak acid pH problem and is just a disguised ICE equilibrium problem. Look at the question this way: if initially 1g of CO2 is dissolved in 100 mL of water, what is the equilibrium concentration of products?)
How good was your estimate in step C?
Find the initial moles of CO2
1g/44 gmol-1 = 0.023 mol CO2
i) 0.023 -- 0 1x10-7
c) -x -- +x +x
e) (0.023 — x)/0.10L -- x/0.10L (1x10-7+ x)/0.10L ~ x/0.10L
Ka = 10(-7.88) = 1.33x10-8
1.33x10-8 = (x2/0.10)/(0.023-x) (assume x is negligible compared to 0.023)
x = [H3O+] = 5.5x10-6 (assumption o.k.)
pH = -log [2.76x10-6] = 5.25
G. Would the definition of a Bronsted acid that you learned in lecture on Monday predict the affect of CO2 on the pH of water? Why or why not?
Bronsted Acids release protons, which CO2 does not but a product of CO2 in water does — so sort of.
H. Now add 2 pipets of 1.0 M NaOH to the flask, dropwise. How has the pH of the solution changed qualitatively? What has happened to the hydronium ion concentration? Write down the reaction of the H3O+ with OH- including the color change.
It has become more basic indicated by a change in color from yellow to blue. The base has neutralized (decreased the concentration of H3O+) the hydronium ions.
H3O+ + OH- ó H2O
I. CO2 is an example of an acid anhydride, which is an "acid without water." The Bronsted acid is formed by adding 1 molecule of H2O to the acid anhydride. (You did this in step D. )
Other acid anhydrides from automobile exhaust that are major contributors to acid rain are NO2 and SO3. What common strong acids are NO2 and SO3 related to? (hint: Show reaction of each with H2O.)
NO2 + H2O ó HNO3 & SO3 + H2O ó H2SO4