EXPERIMENT 8 A SIMPLE TITRATION

Transcription

1 EXPERIMENT 8 A SIMPLE TITRATION Structure 8.1 Introduction Objectives 8.2 Titration Types of Indicators Types of Titrations Standard Solution 8.3 Titrimetric Experiment: Determination of the strength of given Sodium Hydroxide Solution Principle Requirements Procedure Observations Calculations Result 8.4 Summary 8.1 INTRODUCTION Titrimetric analysis or volumetric analysis is an important chemical method of analysis for determination of the concentration of various solutions. Titrimetric analysis is based on quantitative performance of suitable chemical reactions. Therefore it is a quantitative method of analysis. Titrimetric analysis is performed by accurately measuring the volume of a standard solution which is required to completely react with a known volume of an unknown solution. Therefore, it is also known as volumetric analysis. Laboratory technicians are required to perform a number of titrations in the laboratory. In this experiment, you will perform a titration of oxalic acid with sodium hydroxide. Objectives After performing this experiment, you should be able to: I prepare a standard solution, perform a titration of oxalic acid with sodium hydroxide, determine the concentration of sodium hydroxide solution, perform any other titration provided proper instructions are given, determine end-point in a titration, and classify various types of titrations. 8.2 TITRATION In titrimetric analysis, one determines the volume of a standard solution which is required to react quantitatively with a known volume of the other solution, the concentration of which is to be determined. For this purpose, an aliquot of the solution to be estimated is pipetted out and is transferred into a conical flask. The standard solution is added dropwise from a burette to the solution in the conical flask.

2 Basic Experiments in Chemistry The conical flask is continuously shaken to enable the two solutions to mix thorougldy. Standard solution (Unit 7 of Block 2) is added till the two solutions react quantitatively. This process is called tiration. The solution in the conical flask is called the titrand and the one in the burette is called the titrant. The total volume of titrant used in the reaction is called the titre. We have said above that in a titration, the titrant is added till it reacts quantitatively with the titrand. Such a stage, at which the quantities of titrant and titrand are in their stoichiometric proportions (in terms of equivalents or moles), is called the equivalence point. A question arises now, as to how do we know that the equivalence point has been reached? At what stage shall we stop adding the solution from the burette? Essentially we need some substance which can indicate this stage by a change in a physical property like colour. A substance which is used to indicate the equivalence point of a titration through a colour change is called an indicator. Equivalence point so obtained is called end point. It is not necessary that the end point is coincident with the equivalence point, because of the delay in getting the indicator to show the change, and other factors. Ideally end point and equivalence point should be as close as possible. The indicator, to be used in a given titration, would depend on the nature of the chemical reaction involved between the two reacting solutions. The basic requirement for an indicator is that it should have distinctly different colours before and after the end point because we need to know the end point visually. If no visual indicator is available, the detection of equivalence point can often be achieved by following the course of the tihation by measuring the potential difference between an indicator electrode md a reference electrode or the change in the conductivity of the solution Types of Indicators The indicators can be of three types depending upon their usage: i) Internal indicators: These have to be addeq! into the reaction solutian. Examples are: phenolphthalein, methyl oranpe, diphenylamine, etc. ii) iii) External indicators: These are not added into the solution. The indicator is kept out on a plate. A drop of the solution being titrated is taken out with the help of a rod and put on the indicator. A change in colour indicates the end point. Potassium ferricyanide is one such example. Self-indicators: Sometimes either the titrand or the titrant changes its colour at the end point and acts as a self-indicator. The example is potassium permanganate used in permmgmatometry Types of Titrations Depending upon the nature of the chemical reaction involved in a titration, the latter can be classified into the following types: i) Acid-base titrations or Neutralisation titrations: The reaction in which an acid reacts with a base to give salt and water is called a neutralisation reaction and the titration involving such a reaction is called an acid-base titration or a neutralisation titration. An example is the reaction between NaOH and HCI : NaOH + HC1 ---b NaCl + H20

3 The indicators used in these titrations depend upon the ph at the end point, the familiar examples are phenolphthalein and methyl orange. Experiment 8 t ii) Oxidation-Reduction or Redox titrations: Titrations involving oxidation-reduction reactions, i.e., those in which one component gets oxidised while the other gets reduced are known as redox titrations. An example is the titration between oxalic acid and potassium permanganate in acidic medium, in permanganatometry. In this case, potassium permanganate gets reduced to ~ n while ~ oxalic + acid gets oxidised to CO2 and water. In this titration, potassium permanganate acts as a self-indicator. The following equation represents the reaction: Titrations involving potassium dichromate (Chromatometry) and iodine (iodometry) are also examples of redox titrations. iii) Precipitation titrations: In certain reactions, when the two components react, a precipitate is formed. The end point is indicated by the completion of precipitation. Such reactions are termed as precipitation reactions and the tirations as the precipitation titrations; an example is the titration between potassium chloride and silver nitrate as per the following equation:.. Titrations involving AgN03 are also called argentometric titrations. iv) Coslplexometric titrations: A complexation reaction involves the replacement of one or more of the coordinated solvent molecules, which are coordianted to a central metal ion, M, by some other groups. The groups getting attached to the central ion are known as ligands, L. M(H20), + nl =+ ML, + nhzo The titration involving such type of a reaction is called a complexometric titration. For example, determination of hardness of water using ethylenediaminete&aacetic acid (EDTA) as the complexing agent. The indicator used in thr case is eriochrome black T Standard Solution A standard solution is defined as the one whose concentration (strength) is known accurately, i.e., we know exactly how much of the solute is dissolved in a known volume of the solution. A standard solution may be prepared by dissolving an accurately weighed, pure stable solid (solute) in an appropriate solvent. Preparation of a standard solution is generally the first step in any quantitative experiment, so it is important to know how to prepare a sta~dard solution. Primary and Secondary Standards In titrimetry, certain chemicals are used frequently in defined concentrations as reference solutions. Such substances are classified as primary standards or secondary standards. A primary standard is a compound of sufficient purity

4 Basic Experiments in Chemistry ' from which a standard solution can be prepared by weighing a quantity of it directly, followed by dilution to give a definite volume of the solution. The following specifications have to be satisfied for a substance to qualify as a primary standard: 1. It must be easily available and easy to preserve. 2. It should not be hygroscopic nor should it be otherwise affected by air. 3. It should be readily soluble in the given solvent. 4. The reaction with a standard solution should be stoichiometric. 5. The titration error should be negligible. Few available primary standards for acid-base, redox and comlexometric titrations are given in Table 8.1. Table 8.1: Some primary standards. S.No. 1. Compound Potassium hydrogen phthalate (KHP) Formula unit CsHs04K Relative molar mass, M, Type of titration Acid-base 2. Anhydrous sodium carbonate Na2COj Acid-base 3. I I I I Sodium Salt of EDTA Na2HZCIOH1208N2.2HZ Complexometric 4. Copper(I1) sulphate CuS04.5H Iodometric I I I 1 Potassium dichromate I I I I Arsenic(II1) oxide Potassium iodate K2Cr207 I I I I I I I Sodium oxalate As20; KIO3 NazCz Redox Redox Redox Redox 9. "' Fendus-ammonium FeS04.(NH4)2S04.6H Redox sulphate I I,- r-- I I 10. Oxalic acid (COOH)7.2H20 63.OO Redox/acid- base I 1 Solutions prepared from the primary standards are called primary standard solutions. Substances which do not satisfy all the above conditions, are known as secondary standards. In such cases a direct preparation of a standard solution is not possible. Examples are alkali hydroxides and various inorganic acids. These substances cannot be obtained in pure form. Therefore, concentration of these can be determined by titrating them against primary standard solutions. This process is called standardisation and the solution so standardised is called a secondary standard solution:

5 Preparation of a Standard Solution Experiment 8 To prepare a standard solution of volume, v cm3, of known molarity, M mol dm", the mass of the solute required, m g, of molar mass M, can be calculated as follows: Mass of the solute (m) = M.M,.V 1000 g The solute is then weighed on an analytical balance, transferred into a standard flask and dissolved first in a small quantity of the solvent, the solution is then made up to the mark and shaken thoroughly to get a homogeneous solution. In preparing a standard solution whose concentration is, say, around 0.1 M, the amount of the substance weighed need not be exactly equal to that corresponding to 0.1 M. It can be slightly less or more, but the weighing must be accurate. From the weight of the solute actually taken, molarity of the solution can be calculated using Eq TITRIMETRIC EXPERIMENT: DETERMINATION OF THE STRENGTH OF GIVEN SODIUM HYDROXIDE SOLUTION Having learnt about titration in general, types of titrations and indicators, you would now like to learn how you would do an experiment, make observations, record data and calculate the result. It is also important to examine the result critically, compare it with known or expected value, look for the sources of error so that improvement can be made. We will illustrate all this in the following example. Of course, you will have to perform various experiments according to the procedure given in each case. We consider here a simple titration involving a weak acid and a strong base, viz. oxalic acid and NaOH, using phenolphthalein as the indicator Principle! Sodium hydroxide is not a primary standard. Therefore, it should be standardised with a suitable primary standard such as oxalic acid. Sodium hydroxide reacts with oxalic acid according to the following equation:. From the above reaction, you can see that sodium hydroxide reacts with oxalic acid in 2: 1 molar ratio. Hence No. of molesof oxalicacid No. of molesof sodium hydroxide 2 where 1111 = molarity of oxalic acid Vl = volume of oxalic acid taken MZ = molarity of sodium hydroxide

6 Basic Experiments in Chemistry V2 = volume of sodium hydroxide taken A slight excess of sodium hydroxide at the end point imparts a distinct pink colour to the solution, when phenolphthalein is used as an indicator Requirements Apparatus Chemicals Analytical balance 1 No. Oxalic acid Beaker 400 c d capacity 1 No. Sodium hydroxide solution Burette 50 cm3 capacity 1 No. (approximately MI10) Burette stand 1 No. Phenolphthalein indicator Conical flask 250 cm3 capacity 1 No. Funnel small 1 No. Pipette 20 cm3 capacity 1 No. Volumetric flask 250 cm3 capacity 1 No. Wash bottle 1 No. Weighing bottle 1 No Procedure 1. Preparation of a standard solution of oxalic acid As the concentration of the given sodium hydroxide solution is approximately 0.1 M, you will have to prepare a standard solution of oxalic acid of about 0.05 M concentration. As the molar mass of oxalic acid is 63, you will require g of oxalic acid for preparing 250 cm3 of 0.05 M solution. This can be calculated as shown below. m = MxM,xV = 0.05 mol dm-3 x g mol-' x 0.25 dm3 = g In Unit 2 of this course, you have studied the handling of different types of analytical balances. Weigh out an empty weighing bottle on an analytical balance and record its mass. Then weigh out the weighing bottle with about 1.60 g of pure oxalic acid accurate1 and record the mass in your note book. Then transfer the solid to a 250 cm Y clean volumetric flask through a glass funnel. Weigh the weighing bottle again accurately and record its mass. Dissolve the solid in cm3 of distilled or deionised water. Make the solution up to the mark wi!h distilled water. Stopper the flask and shake it well to make the solution homogeneous. 2. Standardisation of sodium hydroxide solution First collect the sodium hydroxide solution in a 250 cm3 bottle from your counsellor. Take a clean burette. Rinse the burette with sodium hydroxide solution and fill it up with this solution. Note the initial reading of the burette and record it in the observation Table 8.2 under the initial readin column. Pipette out 2fl cm3 of standard oxalic acid solution into a 250 cm Q conical flakk. Add one or two drops of phenolphthalein indicator. Titrate this solition by slowly adding small amounts of sodium hydroxide solutioa from the burette and continuously shaking the conical flask. Continue adding sodium hydroxide solution until a permanent pink colour appears. This indicates the end point of the titration. Note the burette reading and record it in the observation Table 8.2

7 under the 'final reading' column. The difference of the two readings gives the volume of NaOH used. Experiment 8 t Repeat the titration to get at least two concordant readings to ensure a correct and exact measurement. Record your readings in Table 8.1 and calculate the strength of sodium hydroxide solution. This solution now can be used to determine the strength of other acid solutions Observations Approximate mass of the weighing bottle Mass of weighing bottle + oxalic acid =rnl =... g =m2=... g I I i Mass of weighing bottle after transferring oxalic acid = m3 = g Molar mass (Mr) of oxalic acid = 63.0 g mol-' Volume of oialic acid solution prepared = 250 cm3 = 0.25 dm3 Table 8.2: Titration of oxalie acid with sodium hydroxide solution r S.No. Volume of oxalic Burette Reading Volume of NaOH in acid cm3 (Final-Initial) vt Initial Final v cm cm cm Calculations m Molarity of oxalic acid solution MI = - Mr. V Volume of oxalic acid taken Vl = 20.0 cm3 Volume of sodium hydroxide used = V2 cm3 Molarity of sodium hydroxide solution = M2 Using Eq. 8.2, - (m2 -m3) " 4 mol dm- 63

8 Basic Experiments inshemistry Result Molarity of sodium hydroxide solution = mol dm SUMMARY In this experiment, you have performed the following: prepared a standard solution of oxalic acid, performed an acidimetric titration of a standard solution of oxalic acid against a sodium hydroxide solution, determined the strength of the sodium hydroxide solution.