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).
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)
reversible arrows tell us
ionization reaction does not go to completion. Sometime after
acetic acid (CH3CO2H or HAc) is mixed with water,
reverse of the ionization process (combination) will begin to occur as
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
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:
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.]
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:
Preparing the Sodium Hydroxide Solution:
(Updated 1/1/13 by C.R. Snelling)