Introduction
An acid that gives two hydrogen ions for every acid molecule is known as diprotic acid. Examples of such acids are sulfuric acid (H2SO4), hydrogen sulfide (H2S), oxalic acid (H2C2O4), chromic acid (H2CrO4) and carbonic acid (H2CO3). Dissociation of a diprotic acid in water takes place in two stages. In the first stage, the diprotic acid dissociates to give one hydrogen ion and a monoprotic form of the acid.
H2X(aq) H+(aq) + HX-(aq)
The second stage involves the dissociation of the monoprotic form of the acid to give the second hydrogen ion and fully dissociated form of the acid.
HX-(aq) H+(aq) + X2-(aq)
The successive dissociation of the diprotic acids causes their titration curves to have equivalence points.
An acid refers to a substance that donates a proton while a base is the substance that accepts a proton. The acids that donate only one proton are referred to as a monoprotic acid while the base that can accept only one proton is referred to as a monoprotic base. Polyprotic acids are those acids that donate more than one proton while a polyprotic base is a base that accepts more than one proton. Diprotic acids and bases and triprotic acids and bases belong to the class of polyprotic acids and are capable of donating or accepting two or three protons respectively. In a titration process, polyprotic acids show many equivalence points as the number of their protons when titration curves are drawn. This means that a triprotic acid shows three equivalence points while a diprotic acid shows two equivalence points.
The Ka value numbers are equal to the number of the protons with the first Ka having the largest value followed by the second and so on. Reduction in Ka value is an indication that the protons produced are successfully less acidic as the acid donates them. This is due to the fact that the conjugate bases become less stable than their former state.
The titration process refers to the quantitative determination of an analyte that is in solution. In an acid and base reaction, the reaction continues until one of the reactants is totally consumed (Brian, 2000). Using a base solution with a concentration that is known can thus be used in titrating an acid solution whose concentration is not known. Through the known values of the base solution, the unknown values of the acid are determined. In a similar manner, an acid solution whose concentration is known can be used in titration of a base solution whose concentration is not known. Titration process makes use of the neutralization reaction that occurs between the acids and bases to determine the amount of the titrate needed to neutralize the solution whose concentration is known. Titration of acid and bases can also be used to calculate the percent purity of chemicals. The percent of purity of a specified element or compound that is in an impure substance may be determined by calculating the amount of the pure substance that is present in the impure substances. The two values are then used to determine the percentage purity of the chemical.
The process of determining the concentration of the substances that are basic in nature using the specialized analytic tool of the acid–base titrations is referred to as alkalimetry. The process of determining the concentration of the substances that are acidic in nature using the specialized analytic tool of the acid–base titrations is referred to as acidimetry or acidometry. During titration, the point where equivalent quantities of the reactants have reacted is referred to as equivalence point. This point is usually indicated by a chemical indicator that is added to the reaction mixture (Helmenstine)
In order to detect the end point in a titration process, an indicator is used. The indicator is a halochromic chemical compound which is added in small quantities to a solution in order to determine the acidity or basicity of the solution visually. The chemical indicator detects the presence of hydronium or hydrogen ions in a solution. Normally, adding the indicator in a solution causes a change in the color of the solution depending on the level of hydrogen ions in the solution. The indicators in undissolved form have a different color from the indicator in dissolved form of the indicators. Change in color occurs over a range of H+ concentrations.
Several requirements must be met for an analytical titration to take place. These requirements include a fast and complete reaction and the reactants need to be measured accurately. A good example of a titration process is the titration of sulfuric acid with sodium hydroxide that has no known concentration. Sulfuric acid is classified as a strong acid since the value of its Ka after losing the first proton greater than one. Expressing this acid-base reaction using a balanced equation gives
2 NaOH + H2SO4→Na2SO4 + 2H2O
In such a reaction, two moles of NaOH reacts with 1 mole of sulfuric acid to give 1 mole of sodium sulfate salt and two moles of water. The mole ratio in this reaction is different from the ones in a reaction between sodium hydroxide and hydrochloric acid. As shown in the equation below, one mole of sodium hydroxide reacts with one mole of hydrochloric acid to give one mole of sodium chloride salt and one mole of water.
NaOH + HCl→NaCl+H2O
The aim of the experiment was to determine the percentage of sulfuric acid by calculating the grams of the sulfuric acid in the unknown sample. The diprotic acid was titrated with NaOH solution whose concentration was first determined. The moles of the acid were determined using the moles of NaOH which was equal to the moles of NaOH at the first equivalence point.
Procedure
Materials
The materials that were necessary for the experiment were collected and included the following:
- About 300 mL of approximately 0.1 M NaOH in a 400-mL beaker
- A burette
- A 15.00 mL pipette
- A pipette bulb
- The unknown sample
- Potassium hydrogen phthalate (KHP)
- Phenolphthalein
Part A: Standardization of NaOH
About 0.6 g of potassium hydrogen phthalate or potassium acid phthalate was weighed out and placed in a 250 mL flask. The flask was marked as #1 with a pencil. KPH in the flask was mixed with about 50 mL of deionized water. The procedure was repeated with two other samples of KHP and placed in flasks marked #2 and #3. To each flask, 5 drops of phenolphthalein were added. The burette was cleared and filled with the 0.1 M NaOH after it was rinsed with 2 small samples (about 5 ml each) of NaOH. The starting volume of NaOH on the burette was recorded to the nearest 0.1mL. The volume was targeted to be somewhere from 0 to 5 mL. The KHP solution in flask #1 was titrated with the NaOH until there was a faint endpoint which stayed pink for about 30 seconds to 1 minute. The new volume of the NaOH in the burette was recorded was recorded. More NaOH was added to the burette to get it to between 0-5 mL again. The titration procedure was followed for flask #2 and repeated for flask #3. The three individual molarities were calculated, and the average molarity of the NaOH determined. The determined molarity was used in Part B.
The average deviation for the Molarity of the NaOH was calculated using the formula below
1n|(Xi-X)|n
Where n = number of samples and X is the mean, and Xi are the individual values.
Part B: Analysis of H2SO4 in water
Three separate 15.0 mL samples of the unknown solution were pipetted out into 3 marked Erlenmeyer flasks, and 4 drops of phenolphthalein indicators added to each sample. Each sample was titrated using the solution from Part A and the initial and the final volume of NaOH used for each sample recorded. The burette and the volumetric pipette were cleaned and returned to their original location.
Conclusion
This experiment aimed to determine the percentage of sulfuric acid by calculating the grams of the sulfuric acid in the unknown sample. The diprotic acid was titrated with NaOH solution whose concentration was first determined. From Part A, mole of NaOH was determined as 0.0049 moles. These moles were used to calculate the number of moles that reacted with sulfuric acid to complete neutralization that was shown by a change in color of the phenolphthalein indicator. Since two moles of NaOH reacts with one mole of H2SO4, the moles of the sulfuric acid in the unknown solution was determined to be 0.00246. Using the moles determined and the volume used the molarity of H2SO4 was calculated as 0.169 M. The mass of H2SO4 was also calculated as 0.208 grams. The percentage H2SO4 in the solution was finally determined to be 1.38% H2SO4 by mass. The number of the unknown sample was C-11-10.
The amount of sulfuric acid present in the unknown sample C-11-10 was very low weighing 0.202 grams out of the 15grams of the impure sample. The calculated average deviation for the molarity of NaOH was 0.0123 M. This deviation indicates that there is a minimal deviation from the average molarity calculated, and the data was very close to the mean. This is also an indication that the procedure of determining the molarity of NaOH is a precise method and has the capability to replicate results with a high precision. The procedure may thus be said to be effective in demonstrating titration as an analytical method of determining the molarity of NaOH.
Presence of deviation in the calculated molarity indicates that some errors were present during the experiment. These errors may have occurred in the course of actualizing the objective of the experiment. Some sources of error that may have affected the obtained results include the intrinsic errors where the end point is different from the equivalent point. This results in more of the acid being added and thus errors in the concentration calculated. The other source of error may have been linked to the accuracy of the volumetric glassware used
There are other random errors that may have resulted to deviation detected in the experiment. These errors are mainly human errors and most of them can be reducing by strictly following the laboratory procedure. Some of the errors may include misjudging the indicator color change near the end point. This results in more or less titrant being used. The other source of error may have been as a result of misreading the volume being measured by not measuring the volume at the wrong angle or counting the graduation marks wrongly. Use of contaminated solutions during the experiment such as using the same pipette to transfer two different solutions without rinsing with distilled water also results to errors in experimental results (Kahlert and Scholz).
Use of solutions with the wrong concentration may also result in difference in the results from the expected ones. This mainly occurs during the standardization process, contamination of the stock solution, or keeping the solution in an open bottle. The amount of indicator solution that was added may have shifted the end point resulting to differences in the results obtained.
Errors resulting from the difference in end point from equivalent point may be reduced by first using blind trials in order to offer an estimation of the titrant quantity that is to be used. Errors resulting from the volumetric glass accuracy may be adjusted by carefully calibrating the glassware or by using high standard glassware. These errors can also be reduced by carefully selecting the volumes of the burettes and pipettes that are used in measuring reagents. Human errors can be reduced by strictly following the laboratory procedure, proper storage of the reagents, observing cleanness during the experiment and soberness when doing the experiment.
Titration at the wrong temperature different from the temperature at which the glassware was meant to work may also affect the accuracy of measuring volumes during experiments. The only way to avoid the error may be by calibrating the glassware once again. This exercise consumes a lot of time and may require prior arrangement (Kahlert and Scholz).
Other sources of error may result from losing solution especially when swirling. Vigorous swirling may result in the solution splashing from the flask before the end point is attained. Some reagents such as the indicators are sensitive to changes in temperature and may affect the results when the experiment is conducted in temperatures different from their ideal temperatures (Kahlert and Scholz).
Works Cited
Helmenstine, Anne Marie. Acid-Base Titrations. 2013. 20 November 2013. <http://chemistry.about.com/od/chemistryquickreview/a/titrationcalc.htm>.
Kahlert, Heike and Fritz Scholz. "Titration Errors." Acid-Base Diagrams (2013): 103-111.
