An Examination of the Process of Titration

Introduction

Titration is defined as being a common lab method of chemical analysis that is used to determine the unknown concentration of an acidic or basic solution. Properties behind titration include a quantity of a base being added in small increments to a quantity of an acid (“6.2 General principles and terms of titration processes,” n.d.) One of the two solutions, the acid or the base, is dropped into the other until the end point or equivalence point is reached. The end point is identified by visually observing the change in the appearance of the solution or by recording a change in a chemical or physical property of the mixture (“Titration: Description,” n.d.) The equivalence point is the point when the amount of base equals the amount of acid, which allows for a neutral solution. The end point and the equivalence point differ in that the end point specifically occurs when the solution changes colour, while the equivalence point occurs when the solution reaches neutrality. The use of titration in Chemistry can be applied to a number of real life situations. Many of these include, but are not limited to: determining unknown concentrations of chemicals of interest in blood and urine, developing new pharmaceuticals, and determining fat content, water content, and concentrations of vitamins in food (“Titration,” n.d.)

The problem of the experiment discussed what methods can be taken to find the concentration of acetic acid by using 0.100 M hydrochloric acid and a solution of sodium hydroxide of an unknown concentration. The hypothesis procured from analyzing the problem was if we used 0.100 M hydrochloric acid and a solution of sodium hydroxide of an unknown concentration, then we can find the concentration of acetic acid in vinegar by finding the concentration of sodium hydroxide through three trials, and using the average concentration of the three trials to calculate the concentration of acetic acid, also known as vinegar.

Materials and Methods

First, the buret and graduated cylinder were rinsed with water. The stopcock on the buret was opened, so water was allowed to pass through the entire buret. The stopcock was then closed, and a funnel was placed on top of the buret. Sodium hydroxide was poured into the buret, until it almost reached the 0 mL mark.

Then, the graduated cylinder was filled with approximately 10 Ml of acid. The first three trials used hydrogen chloride, while the fourth trial used vinegar. The acid was poured into an Erlenmeyer flask, and two drops of phenolphthalein was added to the acid in the flask. The flask was then placed underneath the buret. The initial reading of the buret was recorded. The stopcock was slowly opened, and sodium hydroxide slowly dropped into the flask. As sodium hydroxide dripped into the flask, the flask was constantly swirled until the solution turned a faint pink colour.

After the titration was completed, the base was drained out of the buret and discarded into the sin. The buret was then filled to the top with water and the stopcock was left open to drain. A beaker was placed underneath the buret as it drained. All glassware and equipment were rinsed with water and returned to their original areas.

Results

The question being asked was how the concentration of sodium hydroxide can be found through the use of the concentration of hydrochloric acid. Then by using the calculated concentration of sodium hydroxide as well as the concentration of hydrochloric acid, how can the concentration of acetic acid be calculated?

Data Table 1: Volumes of Solutions per Trial

This table depicts the number of trials taken, the volume of the acid (Va), as well as the volume of the base (Vb) taken at the start of the trial, and the finish of the trial.

Trial #

Va

Vb

1

10 mL

At start: 0.400

At finish: 8.40

Vb: 8.00

2

10 mL

At start: 8.40

At finish: 16.40

Vb: 8.00

3

10 mL

At start: 16.40

At finish: 24.40

Vb: 8.00

Vinegar

10 mL

At start: 24.40 à 40.3à +21.5

At finish: 53.4à 49.7à 29.7

Vb: 46.6

Data Table 2: Calculations

This data table shows the steps taken to calculate the molarity of sodium hydroxide, the average molarity of sodium hydroxide from three trials, the concentration for vinegar, as well as the percent error of our calculations.

MAMB = VAVB

(10.0 mL)(0.100 M) = MB(8.00)

0.125 M = MB

(0.125 + 0.125 + 0.125) / 3

Average MB = 0.125 M

MA(10.0) = (0.125)(46.6)

MA(10.0) = 5.825

MA = 0.583

% Error = 100 * ((0.583-0.800)/0.800)

= (-0.217/0.800) * 100

= -27.1 %

According to Table 1, the volume taken for sodium hydroxide was measured to be the same for each trial (8.00 ml). The volume recorded for vinegar was 46.6 ml. The concentration for vinegar was calculated to be 0.583 M, and the percent error that was calculated was -27.1 %.

Discussion

The concentration of acetic acid was calculated to be 0.583 M. Through this, the percent error was calculated to be -27.1%. The accepted value for the concentration of acetic acid in vinegar is 0.8 M. The percentage of error turned out to be relatively small, so it can be concluded that the experiment was done accurately.

A possible source of scientific error could be that the acids and bases that were used in the experiment, were diluted from the water used to wash the buret and the flask. This error would have affected the results of the experiment because the concentration of the acid would not have been pure. A small amount of water could have diluted the concentration of the acid. The results that were calculated support the hypothesis. The hypothesis was that the concentration of acetic acid in vinegar could be found by using the average concentration of sodium hydroxide. Both the average concentration of sodium hydroxide was found, as well as the concentration of acetic acid in vinegar. The average concentration was 0.125 M and the using that, the concentration of acetic acid was calculated to be 0.583 M.

The sodium hydroxide solution needed to be standardized in order to titrate the solution of vinegar so that the concentration of sodium hydroxide could be found and used to find/compare the solution of vinegar. At the equivalence point, the number of moles of the acid equals the number of moles of the base. The acid is completely neutralized, and this is observed through a colour change in the indicator. Titration could be used to determine the concentration of a basic or acidic solution by using the known concentration and known volume of the acid/base and setting that in proportion to the known volume of the acid/base, so that the unknown concentration could be found. Phenolphthalein is a better indicator for titration than methyl orange because the titration neutralizes the solution at a pH of 8-9, which can be seen by phenolphthalein, but not by methyl orange. Based on the molar concentration of acetic acid in vinegar that was experimentally determined, the vinegar was legal. The molar concentration of acetic acid (0.583 M) was multiplied with the gram formula mass of acetic acid. The product was 34.98 g. The mass of acetic acid was then divided by the product of the total mass of the solution and the mass of water. The quotient of 34.98 g divided by 1034.98 was then multiplied by 100 to find the nearest percent of acetic acid. The end result was 3.8%. Since 3.8% is less than 4%, the vinegar is legal.