Experiment 03 Concentration measurement of a solution of sodium hydroxide by titration By Christian Redeker 9.10.2007
Contents 1.) Hypothesis...3 1.1) Acid base reactions...3 1.2) Concentration measurement with the help of titration...4 2.) Diagram...5 3.) Method...5 3.1) The preparation of a standard solution of ethanedioic acid...5 3.1.1) Apparatus...5 3.1.2) Procedure...5 3.2) Acid base titration to find the concentration of a solution of sodium hydroxide...6 3.2.1) Apparatus...6 3.2.1) Procedure Titration...6 4.) Errors...6 4.1) Systematic errors...6 5.) Results...8 6.) Conclusion/Evaluation...10 7.) Bibliography...10 2
1.) Hypothesis Titration is one of the fundamental experimental methods in chemistry. A titration can determine the volume of one solution required to react exactly with a known volume of another solution 1. With the help of molar calculations the concentration of one of the solutions can be determined afterwards. In this experiment, titration is used to determine the concentration of a base (sodium hydroxid). The second substances involved in the titration is ethanedioic acid (C 2 H 2 O 4 ). Therefore the reaction involved in the titration is a acid-base reaction. The knowledge of the fundamentals of mole calculations and of acid-base reactions is a prerequisite for the understanding of the principle of titration as used in this experiment. Therefore the fundamentals of acid-base reactions are briefly shown in the following paragraphs, while in chapter 5 the mole calculations are explained at the example the experiment delivers. Chapter 1.2 gives more detailed introduction into titration and how it is used in the experiment. 1.1) Acid-base reactions Acids are ionic compounds in which at least one single positively charged hydrogen (H + ) ion is present. Due to their ionic nature, acids dissolve in water. The polarity of the water molecules separates the ions of the compounds, so that H + (aq) particles are present in the solution. Due to the positive charge of the H + (aq) ion it combines with one lone pair of a water molecule to form a single coordinate bond (both electrons involved in the bond come from the oxygen atom of the water molecule). The resulting molecule is an oxonium ion ((H 3 O) + (aq)). Therefore it is more correctly to speak about a ((H 3 O) + (aq)) ion than of a H +. However, usually it is spoken of H + ions in combination with acids instead of ((H 3 O) + (aq)) ions. While a single positively charged hydrogen ion is common to all acids, a hydroxide ion, (OH) - is common to all aqueous bases. Not all bases are soluble in water (e.g. Cu(II)O); a base which is soluble in water is called an alkali 2. When an alkali dissolves in water, the ions involved in the ionic compound are separated from each other due to the polarity of the water molecules. Therefore (OH) - (aq) ions are present in an alkaline solution. The concentration of H + ions in a solution can be given by the ph scale. ph is defined as the negative logarithm, to the base 10, of the aqueous hydrogen ion concentration [H 3 O + ] measured in mol/dm 3 3. Because an H + ion is common to all acids, the acidity of a solution is given by the ph value of the solution. The more H + ions there are in a solution, the more acidic is the solution. The ph of pure water is 7. Is the ph of a solution below 7 the solution is said to be an acidic solution, because the solution has a higher concentration of free protons than pure water 4. Is the ph above 7 alkalis have removed H + ions from the solution by neutralization reactions (see next paragraph). A change in ph can be measured with particular substances (indicators) which change their colour when the ph of the surroundings changes. 3
Bases and acids react with each other in so called neutralization reactions. It plays no role if the base is an alkali or non-soluble in water. Common to all neutralization reactions is the combination of the H + ions from the acid with the OH - ions from the base to form water. A typically example of a neutralization reaction is the reaction between hydrochloric acid (HCl (aq) ) and the alkali sodium hydroxide (NaOH (aq) ). The reaction is expressed by the following equation: 1.2) Concentration measurement with the help of titration The titration used in this experiment uses neutralization reactions to determine the concentration of an alkali (sodium hydroxide) in an alkaline solution. Thereby the volume of sodium hydroxide solution is measured which reacts with all molecules of aqueous ethanedioic acid in a solution of known volume and concentration. The ethanedioic acid solution is due to the fact that concentration and the volume of it are known called a standard solution. It is essential to know the concentration or the volume of one of the reagent precisely to be able to calculate the concentration of the other with the help of the results of the titration. The reaction which takes place in the experiment is a neutralization reaction due to reagents being an alkali (sodium hydroxide) and an acid (ethanedioic acid). It is described by the following equation: During the titration sodium hydroxide solution is carefully added to the ethanedioic acid. Before the titration begins, an indicator is given into the ethanedioic acid solution. The type of the indicator depends on the substances involved in the measurements. The indicator gives a visual indication of when the reaction is completed 5. Practically speaking, this means the indicator changes its colour when all ethanedioic acid molecules have reacted with sodium hydroxide molecules. This point is called the equivalence point. Actually, it is not totally correct that the equivalence point and the point at which the indicator changes its colour is the same, because the indicator changes its colour when the reagent which is in excess in the solution changes. That in turns means for the experiment it changes its colour when one drop too much of sodium hydroxide solution is added to the ethanedioic acid solution. However, within the volume of one drop of reagent added from the burette, the end point (the point at which the reagent in the burette is in excess in the solution) is usually the same as the equivalence point, if the indicator is chosen carefully 6. When the indicator changes its colour the adding of sodium hydroxide is stopped. The volume of the solution which was added to the ethanedioic acid solution is measured with the scale of the burette. Afterwards the concentration of the sodium hydroxide can be determined using mole calculations as lined out in chapter 5. 4
2.) Diagram 3.) Method The procedure of the experiment can be divided into the two following steps: 1.) Preparing of a standard solution of ethanedioic acid as the raw material for the titration (see chapter 1) 2.) The actual acid-base titration to find the concentration of the sodium hydroxide solution 3.1) The preparation of a standard solution of ethanedioic acid 3.1.1) Apparatus 250 cm 3 beaker, ethanedioic acid powder, scale, distilled water, filter funnel, 250 cm 3 standard flask, pipette, 3.1.2) Procedure Approximately 0.77 g ethanedioic acid powder was given into the 250 cm 3 beaker. Approximately 100 cm 3 distilled water was added to the ethanedioic acid in the beaker. The 5
beaker was carefully shaken, so that the ethanedioic acid dissolved in the water. A filter funnel was used to transfer the solution into the 250 cm 3 standard flask. The beaker was rinsed carefully in order to make sure that almost all the solution went into the standard flask. Distilled water was added until the level was close to the mark on the neck of the flask. Afterwards, the filter funnel was removed and using a pipette water was added to bring the bottom of the meniscus up to the mark. Therefore 250 cm 3 of ethanedioic acid solution were in the flask. 3.2) Acid-base titration to find the concentration of a solution of sodium hydroxide 3.2.1) Apparatus Burette, sodium hydroxide solution, ethanedioic acid solution (preparation of ethanedioic acid solution see chapter 3.1), 100 cm 3 beaker, 25 cm 3 pipette, 250 cm 3 conical flask, phenolphthalein indicator 3.2.1) Procedure Titration The principle set-up of the titration is shown in figure 1. The burette was rinsed out with sodium hydroxide solution. Afterwards, the burette was filled with the same solution, and then the tap was opened, so that the solution could flow out of the burette. The burette was closed when the bottom of the meniscus reached the zero mark on the burette. Some of the ethanedioic acid solution was transferred into a 100 cm 3 beaker. With a 25 cm 3 pipette, 25 cm 3 of the solution were transferred to a clean 250 cm 3 conical flask. Subsequently, three drops of phenolphthalein indicator were added to the solution in the conical flask. Sodium hydroxide solution was given into the flask until the liquid in the conical flask turned pink; it should be a pale pink. Because too much sodium hydroxide solution was given into the ethanedioic solution in the first attempt (dark pink colour), a second reading was carried out which gave a more precise result. The volume of the sodium hydroxide solution which caused the liquid in the conical flask to turn to a pale pink was noted afterwards. 4.) Errors 4.1) Systematic errors During the experiment volumes and masses were determined several times. Each of the measuring devices had a systematic error. In the following chapter the systematic percentage errors of those measurement devices is calculated. 1.) The volumetric flask was used to determine the volume of the standard solution (250 cm 3 ). The systematic error involved in this measurement was +/- 0.1 cm 3. To obtain the resulting percentage error this value had to be multiplied by 100 divided by the measured volume of the standard solution. The resulting percentage error was calculated as in the following equation: 6
The systematic percentage error made during the volume measurement of the standard solution was therefore 0.04%. 2.) A scale was used to determine the mass of the solid ethanedioic acid (0.77 g) with which the standard solution was prepared. The scale had a systematic error of +/- 0.01 g. To obtain the resulting percentage error this value had to be multiplied by 100 divided by the measured mass of the solid ethanedioic acid as described in the following: The systematic percentage error made during the mass measurement of the ethanedioic acid was therefore 1.3%. 3.) A burette was used to determine the volume of sodium hydroxide (11.5 cm 3 ) which completely reacted with the ehtanedioic acid in the standard solution. The burette had a systematic error of +/- 0.1 cm 3. To obtain the resulting percentage error this value had to be multiplied by 100 divided by the measured volume of sodium hydroxide as described in the following: The systematic percentage error made during the volume measurement of sodium hydroxide solution was therefore 0.87%. 4.) A pipette was used to determine the volume of the ethanedioic acid solution (25 cm 3 ) which reacted with the sodium hydroxide solution. The burette had a systematic error of +/- 0.06 cm 3. To obtain the resulting percentage error this value had to be multiplied by 100 divided by the measured volume of the ethanedioic acid solution as described in the following: The systematic percentage error made during the volume measurement of ethanedioic acid was therefore 0.24%. The total systematic error is given by the square root of the sum of the squares of the percentage errors determined above. This is expressed in the following calculation: The total systematic error was therefore 1.58%. 7
5.) Results This chapter shows how the concentration of the sodium hydroxide solution used in the experiment was determined by using mole calculation. First the concentration of the ethanedioic acid solution had to be determined with the measurements from the experiment. The concentration determination of the ethanedioic acid is not essential to carry out to obtain the concentration of sodium hydroxide; it is only essential to determine the number of moles of ethanedioic acid involved in the reaction between ethanedioic acid and sodium hydroxide. With the stoichiometric equation of the reaction between sodium hydroxide and ethanedioic acid to form and water, the ratio between those two substances can be determined. Therefore the number of moles of sodium hydroxide involved in the reaction can be calculated. With this value and with the help of the volume of sodium hydroxide which reacted with all ethanedioic acid molecules in the standard solution determined by the titration eventually the concentration of the sodium hydroxide can be determined. The concentration (C) of a solution is defined by the number of moles (n) per volume (V). This is expressed by the following equation: To calculate both the number of moles of the solute and the volume of the solution had to be known. Therefore, the volume of the ethanedioic acid solution had to be known to calculate the concentration of the sodium hydroxide solution firstly. The volume of the ethanedioic acid solution was 250 cm 3 as measured with the flasks. Secondly, the number moles of ethanedioic acid in the solution had to be determined. The number of moles (n) is defined as the mass (m) of a substance divided by its molar mass (M r ). The following equation shows this relationship in mathematical terms: The mass of the ethanedioic acid in the standard solution was 0.77 g as measured with the scale. Solid ethanedioic acid has the chemical formula (C 2 H 2 O 4.2H 2 O) and has therefore a molar mass of approximately 126 gmol -1. With those values the number of moles of ethanedioic acid in the standard solution could be calculated as shown in the following: The number of moles of ethanedioic acid in the standard solution is therefore 0.0061 mol. With the values for the number of moles of ethanedioic acid in the solution and the volume of the solution the concentration of the solution could be determined. Due to the fact that the 8
volume of the solution was measured in cubed centimetres and the concentration is usually given mol/dm 3 the value for the volume had to be converted into cubed decimetres. Therefore the value in cubed centimetres was multiplied by 10-3. This is shown in the following equation: The volume of the ethanedioic acid solution was therefore 0.25 dm 3. Afterwards the concentration of the standard solution could be calculated as shown in the following equation: The concentration of the standard solution was therefore 0.0244 mol/dm 3. In the next step the ratio of sodium hydroxide and ethanedioic acid which reacted with each other in the titration had to be determined. Therefore the stoichiometric formula of the reaction was considered, which is the following: The ratio of sodium hydroxide and ethanedioic acid is 2:1. Therefore the number of moles of ethanedioic acid involved in the reaction was multiplied by two to obtain the number of moles of sodium hydroxide involved in the reaction as shown in the following calculation: Therefore 0.0112 mol of sodium hydroxide were involved in the reaction. From the titration it was known that 11.5 cm 3 of sodium hydroxide were responsible for a complete reaction with the aqueous ethanedioic acid. Therefore both values which were required to calculate the concentration of sodium hydroxide in the solution were obtained. As for the concentration calculation of ethanedioic acid the volume of the requested substance, in this case of the sodium hydroxide solution was given in cm 3. Therefore it had to be converted into dm 3. This is shown in the following calculation: Afterwards the concentration of sodium hydroxide in the solution was calculated as shown in the following: 9
Therefore the concentration of sodium hydroxide in the sodium hydroxide solution was 1.06 mol/dm 3. 6.) Conclusion/Evaluation The concentration of sodium hydroxide in the solution could be determined as being 1.06 mol/dm 3. However, it is known that the supposed actual concentration of sodium hydroxide was 0.1 mol dm -3. If this were the actual value of the sodium hydroxide in the solution, the value obtained by the experiment would be very inaccurate. However, in other experiments of the same type, the concentration of sodium hydroxide in the solution was also determined to be roughly 1 mol dm -3. This observation is a hint to the supposed actual concentration of 0.1 mol dm -3 being inaccurate, probably because of a wrong calibration of the scale, which was used during the preparation of the sample of sodium chloride solution. Due to the outcome of the other experiments it is likely that the value for the concentration determined by this experiment is accurate and the supposed actual concentration is wrong. 7.) Bibliography Note 1, 2, 3, 5, 6: Clugston, Michael; Flemming, Rosalind. Advanced Chemistry. (2000). Oxford University Press. pp. 616 Note 4: Kent, Michael. Advanced Biology. (2000). Oxford University Press. pp. 624. 10