|When a solution of an acid is mixed with a solution of a base, a chemical reaction occurs that is called a neutralization reaction. We can represent this reaction by the following net ionic equation:
This reaction is commonly used to change the acidity of solutions. The reaction can be considered to proceed completely to the right since the equilibrium constant for the reaction is about 1014 at room temperature. The limiting reagent is consumed, leaving the solution either acidic or basic, depending on whether H+ or OH- ion was in excess.
Since the reaction is essentially quantitative, it can be used to determine the concentrations of acidic or basic solutions. The process in which this reaction is used to determine the specific concentration of acid or base from a known volume of a standard solution is called titration. A frequently used procedure involves the titration of an acid with a standard solution (accurately known concentration) of base. In the titration, the standard solution is added from a buret to a measured volume of acid solution until the number of moles of OH- ion added is just equal to the number of moles of H+ ion present in the acid. At that point the volume of basic solution that has been added is measured. Therefore, if the molarity of either the H+ or the OH- ion in its solution is known, the molarity of the other ion can be found from the titration.
Traditionally, the end point in the titration is determined by using a chemical called an indicator. The indicators are selected because they have the characteristic of changing color at at a given pH range. Common indicators used in acid-base titrations are weak organic acids or bases that change color when they are neutralized. One of the most common indicators is phenolphthalein, which is colorless in acid solutions but becomes red when the pH of the solution becomes 9 or higher. When using a pH electrode, the pH of the solution is measured, a significant inflection in pH of the solution occurs at the end point.
When a solution of a strong acid is titrated with a solution of a strong base, the pH at the end point will be about 7. At the end point a drop of acid or base added to the solution will change its pH by several pH units, so that phenolphthalein can be used as an indicator in such titrations. If a weak acid is titrated with a strong base, the pH at the equivalence point is somewhat higher than 7, perhaps 8 or 10, and phenolphthalein is still a satisfactory indicator although not the best. If, however, a solution of a weak base such as ammonia is titrated with a strong acid, the pH will be a 4 - 6 at the end point, and phenolphthalein will not be a satisfactory indicator for that titration as, for example, methyl red, whose color changes from red to yellow as the pH changes from about 4 to 6. Ordinarily, indicators will be chosen so that their color change occurs at about the pH at the equivalence point of a given acid-base titration.
Rather than using an indicator to detect the end point of the titration, we will use a pH electrode coupled to a computer to detect the stoichiometric point. In this experiment, you will experimentally determine the molarity of OH- in an NaOH solution by titrating that solution against a solid monoprotic acid, potassium hydrogen phthalate, KHC8H4O4 (abbreviated KHP), dissolved in an accurately known volume of water. Once the NaOH solution's concentration is experimentally determined the basic solution is referred to as the standard NaOH solution.
In the second part of this experiment you will use your standardized NaOH solution to titrate a sample of a pure solid organic acid. By titrating a weighed sample of the unknown acid with your standardized NaOH solution you can find the number of moles of H+ ion that your sample can furnish. From the mass of your sample and the number of moles of H+ it contains, you can calculate the number of grams of solid acid, equivalent mass, that would contain one mole of H+ ion.
The equivalent mass of the acid is equal to the number of grams of an acid per mole of available H+ ion:
The equivalent mass, EM, of an acid may or may not equal the molar mass of the acid, MM. The reason is almost, but not quite, obvious. Let us consider three acids, with the molecular formulas HX, H2Y, and H3Z; presuming that in the titration all of the H+ ion in the acid is neutralized:
In all cases the molar mass and the equivalent mass are related by simple equations, but in order to find the molar mass from the equivalent mass we need to know the molecular formula of the acid.
|WEAR YOUR SAFETY GLASSES
WHILE PERFORMING THIS EXPERIMENT
B. EXPERIMENTAL PROCEDURE
|Note: This experiment is relatively long unless you know precisely what to do. Study the experiment carefully before coming to class so that you don't have to spend a lot of time finding out what the experiment is all about.|
|So that this experiment will consume less time, students will work in pairs. Both the standardization of the NaOH and the determination of the equivalent mass of the unknown acid using your standardized base solution will be determined using LabWorks II to monitor the pH change during the titration. The instructions for using LabWorks II with the pH electrode to measure the change in pH during the titration can be found in this link, instructions.
In traditional titrations, the technique is to add a substance called an indicator. The indicator is selected so that it changes color at the stoichiometric point of the reaction. If one used a buret to add the titrant, the volume of a standard solution need to reach the end point can be measured. We will not use an indicator in this experiment. Instead, we will use a pH electrode coupled with LabWorks II to determine and record when the stoiciometric point has been reached.
|Obtain the following at the general distribution counter: solid potassium hydrogen phthalate, KHC8H4O4 (abbreviated KHP), watch glass, 500-mL Florence flask, rubber stopper, and the solid unknown. The pH electrode will be supplied by your the instructor.|
| 1. Standardization of NaOH Solution
Into a small graduated cylinder pour about 8 - 9 mL of the stock 6 M NaOH (dilute NaOH). The 6M NaOH solution is found on the counter top along with other acids and bases. Dilute the the 8 - 9-mL of 6M NaOH to about 500 mL with distilled water in a 500-mL Florence flask. Stopper the flask tightly and mix the solution thoroughly at intervals over a period of at least 2 minutes before using the solution.
Weigh accurately a clean and dry watch glass to ten thousands of a gram and record the weight. Then add 1.8 grams of KHP to the watch glass and again re-weigh to the ten thousands of a gram and record the weight. Transfer the KHP to a clean 100-mL volumetric flask using deionized water to wash the KHP into the volumetric flask. Then add sufficient deionized water until you have filled the volumetric flask to the calibration mark. You have now prepared 100-mL of KHP solution of known concentration. This amount should provide all the KHP you will need to standardize the NaOH solution you have prepared; do not waste either solution. The 100-mL of KHP solution should be sufficient for three titration's plus the three rinses.
Clean the two burets, then rinse each with distilled water, and rinse each buret three times with the solution to be used. To rinse the buret for the KHP solution, use no more than 2-mL of KHP solution for each rinse, in each case thoroughly wetting the walls of the buret with the solution and then letting it out through the stopcock. Fill the buret with KHP solution; open the stopcock momentarily to fill the tip. Proceed to clean and fill the other buret with your diluted NaOH solution in a similar manner. Put the KHP buret, A, on the left side of your buret clamp, and the base buret, B, on the right side. Check to see that your burets do not leak and there are no air bubbles in either buret tip. Practice read the burets to 0.02 mL, then read and record the levels in the two burets to 0.02 mL.
Draw about 25 mL (record the initial and final volume to 0.02-mL) of the KHC8H4O4, KHP, solution from the buret into a clean 600-mL beaker; add to the beaker about 250 mL of distilled H2O. Set-up the pH electrode magnetic stirrer, stirring, bar, and the buret containing the base as shown in Figure 1.
You must record the initial and final volume NaOH solution added during each run. LabWorks will measure the TOTAL number of drops of base. Hence, you can calculate the volume of solution per drop of NaOH solution, expressed as mLs per drop of NaOH solution. From LabWorks II you can obtain the number of drops of solution used to reach the equivalence point; hence, the volume of NaOH solution used to reach the equivalence point be calculated.
At this time follow the instructions for LabWorks II to perform and record the pH - total number of drops of NaOH solution used.
In the final stages of the titration notice the very rapid change in pH on the graph as seen on the monitor. Continue adding base until the pH again changes only slightly for about sixty seconds. Repeat this process two more times.
Calculate the molarity of your base, MOH- for each of the three experimental runs and calculate an average molarity.
| 2. Determining the Equivalent Mass
of an Unknown Acid
Weigh the vial containing your solid acid on the analytical balance to ± 0.0001 grams. Carefully pour out about 0.20 - 0.25 grams of sample into a clean but not necessarily dry 600-mL beaker. Re-weigh the vial accurately. Add about 250 mL of distilled water to the 600-mL beaker. The acid may be relatively insoluble; don't worry if it doesn't all dissolve.
Fill your NaOH buret with the (now standardized) NaOH solution.
Set-up the pH electrode magnetic stirrer, stirring bar, and the buret containing the base as shown in Figure 1. At this time follow the instructions for LabWorks II to perform and record the pH vs number of drops. Titrate the solution of the solid acid with standardized NaOH solution. As the acid is consumed by the base, the acid will tend to dissolve. If the acid appears to be relatively insoluble, add 25 mL of ethanol to increase the solubility.
In the final stages of the titration notice the very rapid change in pH on the graph as seen on the monitor. Continue adding base until the pH again changes only slightly for about sixty seconds. Repeat this entire process two more times.
- Computer-based measurements (must be completed during the laboratory period).
i. Opening LabWorks II
ii. Calibration Set-up
2nd Obtain and set-up the two pH reference solutions.
iv. Acquire Set-up
v. Analysis Set-up
We must accomplish a number of tasks:
Make certain the Spreadsheet and Graph view is on the computer screen. To print your titration curve, click on the File button found on the far left of the menu toolbar.