Ian Jacobi

Group 6

Section 01

September 30, 2004

Strong Acid-Strong Base Titration Lab

One of the most stereotypical chemical reactions in chemistry is that of an acid and a base. In particular, these reactions are very helpful in order to determine molarity of solutions (naturally of bases and acids only, but that is a given, given that the reactants in the reaction are an acid and base). Given that many reactions occur in aqueous environments, determining the molarity of solutions becomes very important in completing stoichiometry equations and the calculations of reactions. This particular experiment attempts to utilize stoichiometry and standard titration methods with a strong base, sodium hydroxide, in order to calculate the molarity of an unknown solution of hydrochloric acid.

When in solution, sodium hydroxide and hydrochloric acid react as follows:

Given this stoichiometric equation, it is fairly easy to see that for every mole of hydrochloric acid reacted, a mole of sodium hydroxide is required. This may be easy to determine, but it is significantly more difficult to determine the equivalence point of the reaction (that is, the pH where the amount of acid in the solution equals the base). In weak-acid/weak-base solutions and strong-acid/strong-base solutions, this is near 7.0, but in weaker solutions that react with strong solutions, due to buffering, in which a solution partially changes into its conjugate , the equivalence point will not be 7.0, but perhaps higher or lower. In this case, an appropriate indicator must be selected. In the case of sodium hydroxide and hydrochloric acid, however, both are strong, and thus, the equivalence point will be near 7.0, and thus bromothymol blue, with a color change between pH 6 and pH 7.5, is an appropriate indicator for the endpoint of a reaction.

In this particular experiment, it was necessary to calculate the molarity of the unknown solution of hydrochloric acid in this titration. In this case, since the equivalence point is nearly 7.0, it can safely be assumed that, at the equivalence point:

[H+] = [OH-]

Given this fact, it can also be safely assumed that the molarity of the final solution must contain the same number of moles of NaOH as HCl, thus making for a relatively easy equation to solve for initial molarity of the HCl solution:

Where M is the concentration as molarity, and V is the volume in liters. The quantity MV in the numerator represents the molar quantity of the hydroxide ion in the NaOH solution, given that the ratio of hydroxide ions in solution to NaOH atoms is 1:1. Likewise, as the ratio of hydronium ions in solution to Hcl is 1:1, there is no other constant coefficient.

Naturally, these quantities are measured in an experimental environment, primarily involved with a known volume of HCl solution (20.0 mL) in a flask. After this has been prepared, and a certain volume of NaOH has been filled into the pipet so as to be able to be titrated with the HCl and measured, bromothymol blue is added to the HCl solution to act as an indicator of the endpoint of titration. Lastly, the titration is performed carefully in such a fashion as to carefully measure the volume of NaOH necessary to neutralize the HCl (from which the initial volume is subtracted to get the volume titrated). Given these measurements, it is fairly easy to thus calculate the molarity of the HCl through the given equation above.

Reaction #	1	2	3
Initial NaOH Volume	0.00 mL	22.13 mL	0.00 mL
Initial Color	Light yellow	Light yellow	Light yellow
Initial HCl Volume	20.00 mL	20.00 mL	20.00 mL
Final Color	Blue	Blue	Blue
Final NaOH Volume	20.01 mL	42.20 mL	20.00 mL

Table 1. Observed Values

The observations of our particular team (table 1) were surprisingly accurate in terms of error, having a less than one-tenth of a milliliter difference between titrations. This safely eliminated most of the error involved in the experiment, making the calculated values for the molarity of the acid very close (table 2). Please note that the given molarity of NaOH was 0.1153 M.

Reaction #	1	2	3
NaOH Titrated Volume	20.01 mL – 0.00 mL = 20.01 mL	42.20 mL – 22.13 mL = 20.07 mL	20.00 mL – 0.00 mL = 20.00 mL
Moles of NaOH	0.1153 M * 0.02001 L = 0.002307 mol	0.1153 M * 0.02007 L = 0.002314 mol	0.1153 M * 0.02000 L = 0.002306 mol
Molarity of HCl	0.002307 mol / 0.02000 L = 0.1154 M	0.002314 mol / 0.02000 L = 0.1157 M	0.002306 mol / 0.02000 L = 0.1153 M

Table 2. Calculated Values

Given the molarity of the NaOH, the identical stoichiometric coefficients, and the titrated volumes of NaOH, it is not difficult to see that the molarity of the HCl would be very close to that of the NaOH, which is supported by all of the calculated values. The average value of the molarity, given these values however, is 0.1155 M.

In addition, given this value for the average, it is relatively easy to show the high accuracy of this experiment (following equation), illustrating errors of > +/- 0.15%, which, given most chemistry experiments, is quite low.

Indeed, given such a low value of error, it is very unlikely that further experiments would likely lower the error, though it is possible that with continued PRECISE experiments, that the standard deviation and average would be very close to calculated molarities, thus possibly decreasing the standard deviation slightly. As for experimental error, the little that was caused would likely be minimalized by an exact measurement of pH, though even this would not decrease standard deviation much simply due to random errors, as opposed to systematic ones which likely caused the minimal deviation of the calculated HCl molarity (in particular that of the second titration which diverges from the very close calculations of the first and third titrations.

In summary, this lab overall showed how titrations can be very accurate methods of determining concentrations of acids and bases, so as to much more easily determine the ultimate outcome of chemical reactions in aqueous solutions. The problem with this, however, is that the accuracy of titrations are only really useful in acid/base reactions, which do not make up a majority of all chemical reactions (especially in organic chemistry, which is important for biological studies). Nevertheless, the accuracy of titrations are extremely significant, in comparison to that of other reactions, thus making acid/base titrations one of the more easily predictable reactions.

http://www.sparknotes.com/chemistry/acidsbases/titrations/section1.html

Ian Jacobi

2004-09-29T17:30:50