EXPERIMENT: Estimation of acetic acid in vinegar by titrimetry

AIM :

To estimate the amount of acetic acid in vinegar by titrimetry.

THEORETICAL BACKGROUND

Acetic acid CH₃COOH (systematic name ethanoic acid) is generally called as vinegar if obtained by oxidative fermentation of ethanol-containing solutions by acetic acid bacteria. Although vinegar does not entirely exclude diluted chemically produced acetic acid this term is used here to define acetic acid that is produced from ethanol by a primary microbial metabolism, the so-called acetic acid fermentation or vinegar fermentation.

The trivial name, acetic acid. is derived directly from the word for vinegar, which, for example, in Italian Is called aceto The word vinegar itself derives originally from the Latin vinum aegrum, meaning "feeble wine ' In fact, when a wine has "gone off" and has acquired a sour taste, this is due to the oxidation of the ethanol in the wine to acetic acid.

Vinegar fermentation is one of the oldest fermentations known to man occurring naturally as unwanted spoilage of wine. Thus, vinegar Is also known as 'sour wine'. It contains not less than 4 grams of acetic acid per 100 ml along with traces of alcohol, glycerol, esters, reducing sugars and salts.

A number of methods are used to carry out the fermentation. Acetic acid fermentation (acetous fermentation) is an oxidative fermentation by which solutions of ethanol are oxidized to acetic acid and water by acetic acid bacteria using atmospheric oxygen. The oxidation proceeds according to the basic equation

C₂H₅OH + O₂ -> CH₃COOH + H₂O

During acetic acid fermentation, ethanol is almost quantitatively oxidized to acetic acid.

Vinegar has been used since ancient times as an important cooking ingredient, e.g. in salad dressings and on fish and chips. Vinegar derived from red or white wine is the most commonly used form of vinegar in Mediterranean countries and central Europe. The acetic acid content of vinegar can vary widely but for table vinegar it typically ranges from 4 to 8 % v/v. When used for pickling, the acetic acid Content can be as high as 12 %. The purpose of this experiment is to determine the acetic acid content of a commercial vinegar by volumetric analysis.

Volumetric analysis is a technique that employs the measurement at volumes to determine quantitatively the amount of a substance in solution. In any reaction between two or more species, the reaction equation shows the stoichiometric ratio of reacting species. Hence, if the concentration of one of the solutions is known, the concentrations of the others can be determined from the volumes used.

Take, for example, the reaction being investigated in this experiment, that between solutions of acetic acid [CH₃COOH] and sodium hydroxide [NaOH].

The reaction equation is:

CH₃COOH (aq) + OH^- → H₂O + CH₃COOH (aq)

The OH^-(from sodium hydroxide) is present at a known molar concentration. A fixed volume of OH solution is taken and the CH₃COOH solution progressively added until the point at which complete reaction of the substances is reached. This is called the equivalence point, i.e., the point at which the moles of CH₃COOH added equals the number of moles of OH present in the original solution. The incremental process of adding the CH3COOH solution is called titration. It enables the concentration of CH₃COOH to be determined from the stoichiometric equation and the volume of the CH₃COOH solution needed to reach the equivalence point.

Up to the equivalence point the reading solution contains an excess of one reactant (here OH); after the equivalence point the solution contains an excess of the other reactant (here CH₃COOH). We need some means of observing this switch-over. In acid/base titrations (to which the acetic acid/hydroxide titration belongs), the change in the color of a pH indicator is often used to identify the switch-over. The point at which this observed change occurs is called the end point of the reaction. An appropriate indicator is one in which the end point and the equivalence point are as close together as possible.

pH indicators themselves are generally both weak acids and strongly colored dyes, which respond with a dramatic change in color when they react with base. i.e. OH. Because they react with base, in principle they compete with the acid whose concentration you are trying to measure (i.e. CH₃COOH). However because the indicators are so highly colored very low concentrations are required and perturbation of the titration by the indicator in the calculation can be ignored. However the concentration of the indicator has to be kept low.

This experiment uses the indicator, phenolphthalein indicator, which undergoes sudden change in color from pink in basic solution to colorless in acidic solution. The mid-point of the color change occurs at pH of 9.5, i.e. slightly basic. This is a perfect choice as an indicator for the titration of acetic acid with hydroxide ion because the solution at the equivalence point of this titration is slightly basic.

Solutions of sodium hydroxide reacts slowly with carbon dioxide gas present in the atmosphere. Consequently, the solution of sodium hydroxide which you will use in today experiment needs a be standardized. Its concentration needs to accurately determined before use . This is done by titrating with, a solution of a stable acidic compound of high purity which can be very accurately prepared to a known concentration. This is known as the primary standard. In this experiment you will use potassium hydrogen phthalate, KC₈H₅O₄, for this purpose. The equation for its reaction with hydroxide ion is

C₈H₅O₄ + OH^- → H₂O + C₈H₄O_4^-

 

In this experiment you will perform a series of acid/base titrations. In the course of these titrations you will become familiar with the technique of titration and the calculations associated with stoichiometric analysis.

REQUIREMENTS

0.1 M Potassium hydrogen phthalate solution [C₈H₅KO₄, MW = 204.22 ) (500 ml),

0.1 M NaOH solution (250 ml) test vinegars, phenolphthalein indicator ( 1 ml), volumetric flask (250 ml), conical flask (250 ml x 2). pipette 10 ml x 2. 1 ml - 1 ).

burette (50 ml . 1 funnel, distilled water 1000 ml).

 Procedure:

Part A: Standardisation of 0.1 M sodium hydroxide with the primary standard potassium hydrogen phthalate :

1. Use a pipette to transfer 10 mL of 0.1 M NaOH solution to a 100 ml conical flask Add drops of phenolphthalein indicator and swirl to mix Carry out a rapid titration with 0.1 M KC₈H₅O₄ solution in the burette, gently swirling the conical flask as you titrate. Record the rough volume required to reach the endpoint (i.e. the volume required to just observe a persistent color change).

2. pipette another 10 mL of 0.1 M NaOH into a clean conical flask. Add the Indicator. Fill the burette again with the KC₈H₅O₄ solution. Then titrate the KC₈H₅O₄ solution from the burette into the NaOH solution, slowing down to drop by drop as you reach the rough titration volume from your last titration.

3. Record the final volume of KC₈H₅O₄ in the burette at the endpoint and hence calculate the volume of KC₈H₅O₄ solution added (the titre value) in your rough journal.

4. Repeat accurate titrations (in Step 3) until you have constant titre values. Calculate the exact concentration of the 0.1 M solution of NaOH.

Part B : Titration of vinegar with standard sodium hydroxide:

1. Dilute one of the vinegars by a factor of twenty by pipetting 25 mL of it into a 250 ml volumetric flask and making it up to the mark with deionised water. Label the solution, e.g. dilute vinegar . Record in your rough journal which particular type of vinegar you used.

2. Perform a series of titrations with 10 ml of your vinegar solution in the conical flask and the standardized 0.1 M NaOH in the burette (thoroughly rinsed from the last of titrations). Add 4 drops of indicator to your vinegar solution before starting the titration. As in Part A. the solution should start pink and your endpoint will be when the solution in the conical flask is persistently colorless on swirling

3. Repeat your titrations until you have 3 constant readings. Record your titre values in your rough journal

4 Calculate the molar concentration of acetic acid in the vinegar. Don't forget that vinegar was diluted by a factor of 20. Also calculate the concentration of acetic acid in the vinegar as a volume percentage and determine if it falls within the standard range of 4-8 % W/v for table vinegars.