Introduction:

At this point in your chemistry career, you should be able to predict the products of chemical reactions, the states of the products, and whether the reaction will occur spontaneously at any given set of conditions.  You should even be able to determine the rate at which the reactants are consumed and predict the amount of time it would take to produce a given amount of product.  While this is extremely useful information, it only applies to a limited set of reactions, namely those that occur in one direction only:

A  +  B    C  +  D

Here the reactants A and B collide with sufficient energy and the proper geometry to form the products C and D.  What about a reaction in which C and D now become reactants in the opposite direction and form the products A and B?

C  +  D     A  +  B

Initially, when A and B were mixed, the reaction proceeds in the forward direction to produce C and D.  However, as time progresses, the concentration of C and D increases causing an increase in the rate of the reverse reaction.  Concurrent with this increased rate of the reverse reaction is a reduction of the forward rate due to the decrease in the concentration of A and B.  At some point, the rate of the forward and reverse reactions will become the same and we will reach a state of dynamic equilibrium:

A  +  B     C  +  D

This state of dynamic equilibrium does not mean that the forward and reverse reactions have stopped.  Molecules of A and B are still reacting to form C and D and molecules of C and D are reacting to form A and B.  However, since the rate of the forward and reverse reactions is the same, it will appear that nothing is happening.  As such, all quantifiable physical and chemical properties such has pH, color, and concentration will remain constant.

For a general equilibrium equation in which a moles of A react with b moles of B to produce c moles of C and d moles of D,

aA  +  bB     cC  +  dD

We can specify an equilibrium constant, Kequil (same as Kc), that relates the concentration of all product and reactant species,

Where [A], [B], [C], and [D] are the molar concentration of all species present at equilibrium.  The exponents, a, b, c, and d represent the stoichiometric coefficients from the balance chemical reaction.  Kequil is a constant for all conditions at a given temperature (normally 25°C unless otherwise noted).

Purpose:

The purpose of this experiment is to familiarize you with the concept of an equilibrium reaction.  The ionization of a weak acid or weak base is a typical example of an equilibrium process.  Consider the reversible ionization of the classic weak acid, acetic acid:

CH3CO2H  +  H2O     H3O+(aq)  +  CH3CO2-(aq)

The reversible arrows tell us that the ionization reaction does not go to completion.  Sometime after the acetic acid (CH3CO2H or HAc) is mixed with water, the reverse of the ionization process (combination) will begin to occur as the concentrations of the hydronium ion (H3O+ or H +) and acetate ion (CH3CO2- or Ac -) increase.  At some time, the opposing reactions will be occurring at the same rate and the concentrations of all reactants and products will remain constant.  Once we have reached this state of dynamic equilibrium, we can define the equilibrium constant as:

In dilute aqueous solutions, the concentration of H2O is essentially constant at 55.5 M.  Since it is a constant, we can rearrange the equilibrium equation and define an new equilibrium constant for the ionization of weak acids,  Ka:

To calculate the ionization constant, Ka, for acetic acid, it is necessary to experimentally determine the equilibrium concentrations of H+, Ac-, and HAc.

Method:

Based on the discussion above, if we want to determine the Ka for any weak acid (HA), we need to determine the equilibrium concentration of  H+, A- , and HA.  The most straight forward of these is [H+ ], because we know that the pH = -log[H+].  So if we measure the pH of the equilibrium solution, we will not only know the concentration of the hydrogen ion, [H+], but the concentration of the weak acid's conjugate, [A-], as well.  As an example, let's assume that the pH of this solution was 2.37.  This means that the value for both [H+] and [A-] is:

[H+]  = [A-] = 10-pH = 10-2.37 =  4.27x10-3 M

However, we still need to determine the equilibrium concentration of HA.  Unfortunately, this is difficult to determine since most methods of analysis will change the concentration of the HA and cause the equilibrium to shift.  Since we cannot directly determine the [HA], we need to find the initial concentration of HA.  To do this we need to neutralize all of the HA present by titrating it with a strong base of known concentration.  As the H+ from the weak acid is neutralized by the strong base, the equilibrium will shift to the right generating more H+.  This process will continue as the strong base is added until all of the HA has been converted to H + and A- (equivalence point).  This is no longer an equilibrium solution, it only contains A-(aq), Na +(aq), and H2O(l).  For example, if 23.6 mL of 0.321 M NaOH were required to neutralize 50.0 mL of the HA solution, then the initial concentration of HA would have been:

MHAVHA =  MNaOHVNaOH

MHA =  MNaOH VNaOH / VHA

MHA =  0.321 x 23.6 / 50.0

MHA =  0.152 M

Now we can calculate the equilibrium concentration of HA, by subtracting the equilibrium [H+] concentration from the initial HA concentration:

[HA] = 0.152 - 4.27x10-3 = 0.148 M

Now we have all of the equilibrium concentrations necessary to calculate the Ka for our weak acid!

But wait!  What if we don't know the concentration of the strong base we used to titrate the weak acid?  No problem...we will standardize it!  Standardization is a process of comparing an unknown against a known or standard.  In this case, we will titrate a known quantity of standard acid with our unknown base.  Using our M1 V1 = M2V2 relationship, we will be able to determine the exact concentration of our base.

As with all standardization procedures, the real problem is picking an appropriate standard.  A primary standard is a substance that is readily available in a pure form (<0.02% impurities), it is stable, easy to dry, is not hydroscopic, and should have a fairly high equivalent weight to minimize the consequences of errors in mass determination.  We are fortunate that such a standard exists for our situation, the mono potassium salt of the organic di-acid, phthalic acid (KHC8H4O 4, or KHP, mw = 204.223 g/mole).

For example, if we dissolve 1.000 g of KHP in 50 mL of water and titrate this solution with 31.6 mL of our unknown base, what is the molarity of our base?  First we need to remember that at the equivalence point (where the indicator changes color); the moles of KHP equal the moles of NaOH:

moles KHP  =  moles NaOH  =  1.000g / 204.223g/mol  =  0.0049 moles

Since 31.6 mL of our base solution contained 0.0049 moles, the molarity of our base is:

MNaOH  =  0.0049 moles / 0.0316 liters  =  0.155 M

Now that we know the concentration of our base, we can titrate our unknown weak acid to determine its initial concentration, and use the pH meter to determine the equilibrium [H +], and [A-].  With these measurements, it is a simple matter to calculate the Ka for any weak acid.  [Pssst....there are also other ways of determining the Ka for a weak acid, but that is a story for another day.]

Lab Tips:

Although they seem simple, many people initially have trouble with titrations.  There is a good deal of eye-hand coordination involved and a lot of small errors than can creep  in to ruin your experiment. The following are some tips that should help you be successful:

1. Make sure you have studied the video, before coming to lab: An Overview of Titrations
2. Add your base to the burette over the sink.  If you try to add it in the burette clamp, some might spill into your acid sample and you would have to start over.
3. Don't forget to put the indicator in your sample.
4. Make sure you setup your burette so the tip is below the top of the beaker.
5. Set your magnetic stirrer as fast as possible, BUT no splashing.  If any of the acid splashes on the side, you need to wash it down with distilled water.
6. Make sure you have no bubbles in the tip of your burette.  This probably causes 75% of the problems students have with getting titrations to be reproducible.
7. When you see that the pink color start to persist, slow the addition of base to a drop at a time.
8. To get the best equivalence points, you will need to 'cut' drops.  Barely open your stopcock and let less than a drop form on the tip.  Then use your distilled water bottle to squirt it into the beaker.

1. Your burette is a direct read delivery burette.  This means, 'What You See Is What You Get'.  If you start at 0.00mL and stop at 23.56mL, you used 23.56mL.
2. Very important:  read all burette readings to 0.01mL!  Remember, you always read to one place past what is marked on the measuring device.  Also make sure you look directly at the burette at eye level, do not look down or up to read the meniscus, this will cause parallax errors.  The following figure shows an initial volume of 9.62mL and a final volume of 24.16mL:

Procedure:

Preparing the Sodium Hydroxide Solution:

1. Clean and dry a 600 mL beaker .
2. Weigh out approximately 6 grams (to 0.001 g) of sodium hydroxide pellets directly to the 600 mL beaker.  NOTE:  Handle the sodium hydroxide pellets with care. Sodium Hydroxide is very hydroscopic and can cause burns if it comes in contact with your skin.  Be sure to use weighing boats or weighing paper to determine the mass and to deliver the sodium hydroxide pellets to the beaker.  When removing the sodium hydroxide pellets from the reagent container replace the cap of the container as quickly as possible.  Clean up any spilled sodium hydroxide pellets immediately.
3. Add approximately 400 mL of distilled water to the beaker and stir the solution until the pellets have completely dissolved.
4. Put this beaker of sodium hydroxide on a paper towelS so it does not mare the bench top.  Make sure you  immediately clean up any spills.

Titration of NaOH Solution:
1. Obtain a 50 mL burrett, close the stopcock and fill it to the top with distilled water.  Open the stopcock and allow all of the water to drain.
2. Close the stopcock and fill your burrett with 50 mL of your NaOH solution (from above) so that the solution comes in contact with the entire inner surface of the burrett.
3. Open the stopcock and allow all of the NaOH to drain through the tip.
4. Fill the burrett to the top with the NaOH.  Open the stopcock all the way to flush all bubbles out of the tip. When all bubbles have been flushed out (it may take several tries), close the stopcock and refill the burrett.
5. Read the bottom of the meniscus and record the initial reading to the nearest 0.01 mL. The Teflon stopcock should turn smoothly with a little resistance.  If the stopcock is too loose, tighten it a little, otherwise the solution will leak around the stopcock and the titration will be for naught.  A leaking burrett is a major cause of error in titrations.
7. Use a repipetter to deliver 10.00 mL of the standard HCl into a clean 150 mL beaker.  Wash down the sides of the beaker with the wash bottle.  The addition of water at this stage has no effect on the total amount of acid already present in the beaker.  Be sure to record the concentration of the standard HCl!
8. Add 2 drops of phenolphthalein indicator to the acid solution.  The solution should remain colorless.
9. Add a magnetic stir bar to the beaker and place on the heating stir plate.  Adjust the stirring rate to obtain a vortex without any of the solution splattering on the sides of the beaker.  Avoid spilling any of the beaker contents.  Any loss of sample would render the titration worthless.
10. Rinse down the inside of the beaker occasionally and continue, slowly adding NaOH until the first permanent, faint pink color persists for at least 30 seconds.  At this point the titration is complete (the endpoint).
11. Read the final volume of NaOH and record to the nearest 0.01 mL.  Remember these are 'direct read' burretts, what you see is what you get!  Some of you may have used other burretts where they start at 50 mL and go to zero, so you subtract your reading from 50.  DO NOT do that with these burretts!
12. Calculate the normality of the NaOH solution.
13. Repeat Steps #7-12 with 20.00 mL of standard HCl.  Fill the burrett and read it before each titration if there is a  possibility of running out before a titration is complete.
14. Repeat Steps #7-12 with 30.00 mL of standard HCl.
15. The results from these three titrations should agree within ± 0.005 N, otherwise titrate a fourth sample and throw out the value which does not agree.  Average the results and use this value as the concentration of your STANDARDIZED BASE.  Do not discard the base which you have just standardized or clean the burrett.   This standardized NaOH solution will be used to determine the concentration of an unknown acid solution in the next section.

Determination of Initial Weak Acid Concentration:
1. Pick one of the of the unknown weak acids and be sure to record its number.
2. Using the repipetter, transfer 20.00 mL of this acid to a clean 100 mL Erlenmeyer flask.  Use distilled water to make sure all of the acid has been rinsed off the sides of the flask.
4. Put a magnet stirring bar in the flask and set the stir plate to a moderate rate (avoid splashing).
5. Fill your 50 mL burette with your standardized sodium hydroxide solution.
6. Titrate the unknown weak acid with your standardized sodium hydroxide solution until a faint pink color remains for 30 seconds.
7. Repeat Steps 2 - 6, for two additional 20 mL aliquots of the same unknown weak acid and average your results.
8. Once you are satisfied with your results you may pour any remaining sodium hydroxide solution down the sink with copious water.
Determination of Equilibrium Hydrogen Ion Concentration:
1. Add approximately 20 mL of the unknown acid to a clean 50 mL beaker.
2. Remove the pH electrode from its buffer solution and rinse it off with distilled water.
3. Submerge the pH electrode into the unknown acid and wait 10-30 seconds for the meter to stabilize.
4. Record the pH and the temperature of the unknown acid solution.
Clean up:
1. Sodium hydroxide solutions can etch glass and leave white rings on the bench top.  So, it is important to thoroughly clean any glassware that has come in contact with the sodium hydroxide.
2. There is a special procedure for cleaning your burette.  First, open the stopcock and completely drain the sodium hydroxide from your burette.  Then fill it with Burette rinse (squeeze bottle with black tape) and allow it to completely drain.  Finally, fill your burette with distilled water and allow it to drain completely.  Be sure to leave the stopcock open after you are done.
3. Carefully clean the whole bench top including the area around the sink.
Results/Calculations:
• From your titration data, calculate the molarity of your sodium hydroxide solution.
• How reproducible were your results for the molarity?
• From your titration of the weak acid, calculate the it's initial concentration.
• How reproducible were your results for this concentration?
• Knowing the initial concentration of the weak acid and its pH, calculate the Ka for this acid.
• In the back of your textbook is a list of Ka values for several weak acids.  Keep in mind that as careful as you were with your titrations there are several factors outside of  your control so your answer may be a factor of 2, 3, even 5 off from the value in the back of the book.  So use some common sense when deciding which unknown acid you had.
• The Ka values in the back of your textbook have been very accurately determined using a process similar to the one you have used.  One major difference is that all water used in the analysis is freshly boiled before use.  Why was this done?  Don't tell me to make the water pure, I want to know specifically why.
• Phenolphthalein was used as the indicator in this experiment.  What is the role of the indicator?  Why was phenolphthalein used and not some other indicator (see your textbook for other indicators)?

(Updated 8/2/13 by C.R. Snelling)