Titrimetric Determination Of Ethanoic Acid Content Of Vinegar

Step 1

Rinse a cleaned 50 cm3 burette with about 5 cm3 of standardised 0.1 M NaOH solution. (Be sure to record the exact concentration of the NaOH if it is not exactly 0.1 M.) After rinsing, fill the burette with the 0.1 M NaOH solution about 2 cm3 above the 0.0 cm3 mark. Use a clean and dry funnel for filling. Tilting the burette at a 45° angle, slowly turn the stopcock to allow the solution to fill the tip. Collect the excess solution dripping from the tip into a beaker to be discarded later. The air bubbles must be completely removed from the tip. If you do not succeed the first time, repeat it until the liquid in the burette forms one continuous column from top to bottom. Clamp the burette onto a ring stand (Fig. 1). By slowly opening the stopcock, allow the bottom of the meniscus to drop to the 0.0 cm3 mark. Collect the excess solution dripping from the tip into a beaker to be discarded later. Read the meniscus carefully to the nearest 0.1 cm3 (Fig. 2).

Figure 1. Titration setup.

Figure 2. Reading the meniscus.

17.58 cm3 – incorrect

17 cm3 – incorrect

17.5 cm3 – correct 

Step 2

With the aid of a 5 cm3 volumetric pipette, add 5 cm3 of vinegar to a 100 cm3 conical  flask. Allow the vinegar to drain completely from the pipette by holding the pipette in such a manner that its tip touches the wall of the flask. Record the volume of the vinegar added, and the initial volume of the NaOH in the burette. Add a few drops of phenolphthalein indicator to the flask and about 10 cm3 of distilled water. The distilled water is added to dilute the natural color that some commercial vinegars have. In this way, the natural color will not interfere with the color change of the indicator.

Step 3

While holding the neck of the conical flask in your left hand and swirling it, open

the stopcock of the burette slightly with your right hand and allow the dropwise addition of the NaOH to the flask. At the point where the NaOH hits the vinegar solution the color may temporarily turn pink, but this color will disappear upon mixing the solution by swirling. Continue the titration until a faint permanent pink coloration appears. Stop the titration. Record the volume of the NaOH in your burette. Read the meniscus to the nearest 0.1 cm3 (Fig. 2).

Step 4

Repeat the procedures in steps 1–3 until consistent values are obtained.

Categories Of Electrons

The elements have three categories of electrons:

Inner (core) electrons are those seen in the previous noble gas and any completed transition series. They fill all the lower energy levels of an atom.

Outer electrons are those in the highest energy level (highest n value). They spend most of their time farthest from the nucleus.

Valence electrons are those involved in forming compounds. Among the main group elements, the valence electrons are the outer electrons. Among the transition elements, the (n - 1)d electrons are counted among the valence electrons because some or all of them are often involved in bonding.

Thermometric Titration - Advantages And Applications

A thermometric titration utilizes the enthalpy change of the reaction involved to locate the end point. It has been defined as “a titration in an adiabatic system yielding a plot of temperature vs. volume of titrant.” The procedure consists of delivering the titrant from a thermostated buret into a solution contained within a thermally insulated vessel. and observing the temperature change of the solution either upon continuous addition, or after each successive incremental addition, of titrant.

Advantages

These temperature-volume plots resemble the corresponding graphs obtained from conductometric, photometric, and amperometric titrations. Yet while each of the latter three methods is severely limited to specific kinds of systems - e.g., conductometric titration requires electrolytic solutes in solvents of high dielectric constant, almost all reactions exhibit detectable enthalpy changes (positive or negative). This wide applicability, coupled with simplicity, suggests a potential increase in the use of thermometric titrations in analytical chemistry, particularly in those media in which electrometric and photometric methods are inapplicable.

The analytical sensitivity of the thermometric titration method is linearly related to concentration, in contrast to the logarithmic relation to concentration which exists for many other analytical methods, e.g., potentiometric methods. A linear relation is an advantage when very dilute solutions or solutions with high concentrations of interfering ions are being analyzed. For example, a pH titration of a solution containing pyridine at a concentration below 0.05 M will give a poorly defined end point, while the end point of a thermometric titration is well defined.

Applications

Neutralization

Weak acids and bases have been studied by several workers. Bell and Cowell recommended the use of thermometric titration for the preparation of neutral solutions of ammonium citrate. Linde, Rogers, and Hume titrated both weak and strong acids and bases, showing that clear end points were obtainable even in emulsions and thick slurries, and that a mixture of sodium hydroxide and sodium carbonate could be determined with good accuracy. Jordan and coworkers show that, unlike potentiometric titration which is dependent upon free-energy changes, thermometric titration works very well even for extremely weak acids. End points are precise and accurate for acids as weak as boric acid, because the enthalpy change of neutralization is not very different from that of a strong acid.

Complexation

The greatest number of papers on thermometric titration deal with investigations of complex formation, all of which are of analytical importance either directly or indirectly. For EDTA titrations, an accuracy within 3% is possible with cation concentrations as low as 0.0005 M.

Thermometric Titration - Curves

Thermometric enthalpimetric titrations (TET) are characterized by the continuous addition of the titrant to the sample under effectively adiabatic conditions. The total amount of heat evolved (if the reaction is exothermic) or absorbed (if the reaction is endothermic) is monitored using the unbalance potential of a Wheatstone bridge circuit, incorporating a temperature-sensitive semiconductor (thermistor) as one arm of the bridge. Simple styrofoam-insulated reaction cells will maintain pseudoadiabatic conditions for the short period of a titration.

The heat capacity of the system will remain essentially constant if the change in volume of the solution is minimized and if the titrant and titrate are initially at the same temperature (usually room temperature). The TET ethalpogram (a thermometric titration curve), shown in the figure below. illustrates an exothermic titration reaction. The base line AB represents the temperature–time blank, recorded prior to the start of the actual titration. B corresponds to the beginning of the addition of the titrant, C is the end point, and CD is the excess reagent line. In order to minimize variations in heat capacity during titrations, it is customary to use a titrant that is 50 to 100 times more concentrated than the specimen being titrated. Thus the volume of the titrate solutions is maintained  virtually constant, but the titrant is diluted appreciably. Correction for the latter is conveniently made by linear back-extrapolation CB’. Under these conditions, the extrapolated ordinate height, BB’, represents a measure of the change of temperature due to the titration reaction.

Typical thermometric titration curves for an exothermic process: (a) idealized curve, and (b) actual curve, illustrating extrapolation correction for curvature due to incompleteness of reaction. AB—temperature–time blank, slope due to heat leakage into or out of  titration cell; B—start of titration; BC—titration branch; C′—end point; CD—excess reagent line, slope due to heat leakage and temperature difference between reagent and solution being titrated; ΔT— corrected temperature change; ΔV—end-point volume. (From L. Meites, ed., Handbook of Analytical Chemistry, McGraw-Hill, New York, 1963.)