Researching Chemistry

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product, side reactions, position of the equilibrium and the purity of the reactants. Examples of results obtained from an experiment: Mass Weight Mass of round bottomed flask 40.25g Mass of round bottomed flask + ethyl benzoate 45.61g Mass of ethyl benzoate 5.36g Mass of clock glass 9.62g Mass of clock glass + benzoic acid 12.86g Mass of benzoic acid 3.24g From the two-step balanced equations for this reaction it can be noted that 1 mole of ethyl benzoate is required to produce 1 mole of benzoic acid. From this we can work out the theoretical yield of benzoic acid. 1 mole of ethyl benzoate → 1 mole of benzoic acid 150.0g → 122.0g 5.36g (quantity used in experiment) → = 4.36g 4.36 g is the theoretical yield of benzoic acid that will be produced from 5.36 g of ethyl benzoate according to the molar ratios of the balanced equations. As reactions are rarely 100% effective due to the reasons given previously, we can calculate the actual percentage yield of benzoic acid by comparing the theoretical yield to the mass actually produced in the experiment. Theroretical yield = 4.36g Actual yield = 3.24g percentage yield = percentage yield = 4) Researching Chemistry actual yield theoretical yield x 100 = 3.24 4.36 (5.36x122.0) 150 x 100 2 percentage yield = 74.3% Reactant in Excess In the reaction between calcium carbonate and hydrochloric acid they react according to the following equation: CaCO3 + 2HCl → CaCl2 + CO2 + H₂0 The calcium carbonate and hydrochloric acid react in a molar ratio of 1: 2 where 1 mole of calcium carbonate (CaCO3) reacts with 2 moles of hydrochloric acid (HCl ). In this case both reactants would be used up completely in the reaction. If there was a shortage of calcium carbonate for example then the reaction would stop when it runs out and there would be hydrochloric acid left over. In other words the hydrochloric acid would be in excess. Excess Calculation Example Calculate which reactant is in excess, when 10g of calcium carbonate reacts with 50cm³ of 2 mol 1-¹ hydrochloric acid. This calculation involves the reactants only. Firstly work out the number of moles of each reactant involved. Calcium Carbonate n = n = mass gfm 10 100.1 n = 0.1moles Hydrochloric Acid n = concentration x volume (litres) n = 2 x 0.05 n = 0.1moles From the balanced equation 1 mole of calcium carbonate reacts with 2 moles of hydrochloric acid and therefore 0.1 moles of calcium carbonate would react with 0.2 moles of hydrochloric acid. From the previous calculation of the number of moles of hydrochloric acid you can see that there is only 0.1 moles available which is insufficient and therefore the calcium carbonate is in excess and some will be left over at the end of the reaction which will terminate once the hydrochloric acid has been used up. Calculations Associated With Pharmaceuticals Percentage solution by mass is the mass of solute made up to 100cm³ of solution. 4) Researching Chemistry 3 %mass 100 %mass 30% - Example 1 100g of salt solution has 30g of salt in it %mass=(30) x 100 mass of solute mass of solution %volume Percentage solution by volume is the number of cm³ of solute made up to 100cm³ of solution. x 100 = 4) Researching Chemistry volume of solute) × 100 volume of solution Example 2 Wine with 12ml alcohol per 100ml of solution. 12 %volume (¹20) × 100 %volume = 12% Example 3 Making 1000ml of 10% ethylene glycol solution. 10% = 0.1 x 1000 10% 100ml of ethylene glycol The other 900 ml would be made up using water (although this would need to be topped up due to the two liquids not mixing to form exactly 1000 ml). Gravimetric Analysis Measurements for gravimetric analysis are masses in grams and are typically determined using a digital balance that is accurate to two, three or even four decimal places. Gravimetric analysis is a quantitative determination of an analyte based on the mass of a solid. The mass of the analyte present in a substance is determined by changing that chemical substance into solid by precipitation (with an appropriate reagent) of known chemical composition and formula. The precipitate needs to be readily isolated, purified and weighed. Often it is easier to remove the analyte by evaporation. 4 The final product has to be dried completely which is done by 'heating to constant mass'. This involves heating the substance, allowing it to cool in a desiccator (dry atmosphere) and then reweighing it all in a crucible. During heating the crucible lid should be left partially off to allow the water to escape. A blue flame should be used for heating to avoid a build up of soot on the outside of the crucible which could affect the mass. Heating should be started off gently and then more strongly. The mass of the crucible is measured before adding the substance and the final mass of the substance is determined by subtracting the mass of the crucible from the mass of the crucible and dried substance (weighing by difference). This process of heating, drying and weighing is repeated until a constant mass is obtained. This shows that all the water has been driven off. Experiment 1: Gravimetric analysis of water in hydrated barium chloride This experiment determines the value of n in the formula BaCl₂.nH₂O using gravimetric analysis through accurate weighing. The hydrated barium chloride is heated until all the water has been removed as in the equation: BaCl2.nH,O→ BaCl2+nH,O From the masses measured in the experiment it is possible to calculate the relative number of moles of barium chloride and water and hence calculate the value of n (must be a whole number). The experiment is carried out in a crucible which is heated first to remove any residual water, cooled in a desiccator and then weighed accurately. Approximately 2.5g of hydrated barium chloride is then added to the crucible and again weighed accurately (crucible plus contents). The hydrated barium chloride is heated in the crucible using a blue Bunsen flame after which the crucible and contents are placed in a desiccator to cool (desiccator contains a drying agent which reduces the chances of any moisture from the air being absorbed by the sample). After cooling the crucible and contents are reweighed accurately. The heating, cooling and reweighing is repeated until a constant mass is achieved and the assumption that all the water has been removed. Example of results obtained Mass of empty crucible = 32.67g Mass of crucible + hydrated barium chloride = 35.03g Mass of hydrated barium chloride = 35.03 - 32.67 = 2.36g 4) Researching Chemistry 5 Mass of crucible and anhydrous barium chloride = 34.69g (heated to constant mass) Mass of anhydrous barium chloride = 34.69 32.67 = 2.02g Mass of water removed = 2.36 2.02 = 0.34g 0.34 Moles of water removed = = 0.0189 moles 18 Number of moles of barium chloride BaCl2 = Ratio of moles = 0.00970 : 0.0189 BaCl2 H₂O=1:2 Formula of BaCl2.nH₂O is therefore BaCl₂.2H₂O where n = Experiment 2: Determination of nickel using butanedioxime (dimethylglyoxime) Dimethylgloxime is used as a chelating agent in the gravimetric analysis of nickel. It has the formula CH3C(NOH)C(NOH)CH3. 2 H₂C 4) Researching Chemistry H3C Dimethylglyoxime HO CH3 + Ni²+ 2.02 208.3 CH3 Dimethylglyoxime H₂C H3C OH = 0.00970 moles CH3 +2H CH3 Insoluble complex The nickel is precipitated as red nickel dimethylglyoxime by adding an alcoholic solution of dimethylglyoxime and then adding a slight excess of aqueous ammonia solution. When the pH is buffered in the range of 5 to 9, the formation of the red chelate occurs quantitatively in a solution. The chelation reaction occurs due to donation of the electron pairs on the four nitrogen atoms, not by electrons on the oxygen atoms. The reaction is performed in a solution buffered by either an ammonia or citrate buffer to prevent the pH of the solution from falling below 5. The mass of the nickel is determined from the mass of the precipitate which is filtered, washed and dried to constant mass as in the other examples. 6 Information mass of crucible + lid = w1 (w = weight) mass of crucible + lid + dried precipiatate = w2 mass of precipitate = w2w1 = w3 GFM of precipitate Ni(C4H7O2 N2)2 = 288.79 GFM of nickel (Ni) = 58.7g (w3x58.7) 288.7g mass of nickel = Example Calculation When 0.968 g of an impure sample of nickel(II) sulfate, NiSO4.7H2O was dissolved in water and reacted with dimethylglyoxime, 0.942g of the red precipitate was formed. Calculate the percentage of nickel in the impure sample. number of moles of nickel dimethylglyoxime mass of nickel = n x GFM mass of nickel = 0.0032629 × 58.7 mass of nickel - 0.192g percentage of ni el = 0.192 0.968 percentage of nickel = = 19.8% x 100 = 4) Researching Chemistry 0.942 288.7 = 0.0032629 Volumetric Analysis Volumetric analysis involves using a solution of known concentration (standard solution) in a quantitative reaction to determine the concentration of the other reactant. The procedure used to carry out volumetric analysis is titration whether in the form of standard, complexometric or back titrations. Titrations involve measuring one solution quantitatively into a conical flask using a pipette. The other solution is added from a burette until a permanent colour change of an indicator is seen in the conical flask. A 'rough' titration is carried out first followed by more accurate titrations until concordant titre values are achieved (titre volumes added from burette should be 7 ±0.1cm³ of each other). The mean or average value of the concordant titres is used in calculations. ● burette Preparation of a Standard Solution A standard solution is one of which the concentration is known accurately and can be prepared directly from a primary standard. A must have, at least, the following characteristics: ● conical flask a high state of purity; stability in air and water; solubility; • reasonably high formula mass; The standard solution is prepared as follows: • calculate the mass of the primary standard required to make the concentration of solution required in the appropriate volume of solution; weigh out the primary standard as accurately as possible; • dissolve the primary standard in a small volume of deionised water in a beaker; • transfer the solution and all the rinsings into a standard flask; 4) Researching Chemistry titration using a burette and conical flask 8 • make the solution up to the mark with more deionised water; • invert the stoppered standard flask several times to ensure thorough mixing. a standard (volumetric) flask Substances that are used as primary standards include: • Oxalic acid (H2C2O4.2H₂O) • Sodium carbonate (Na2CO3) • Potassium hydrogen phthalate (KH (C8H4O4)) • Potassium iodate (KIO3) • Potassium dichromate (K₂Cr₂O7) Sodium hydroxide (NaOH) is not a primary standard as it has a relatively low GFM, is unstable as a solid (absorbs moisture) and unstable as a solution. Sodium hydroxide solution must be standardised before being used in volumetric analysis. Standardisation is the process of determining the exact concentration of a solution. Titration is one type of analytical procedure often used in standardisation. The point at which the reaction is complete in a titration is referred to as the endpoint. A chemical substance known as an indicator is used to indicate the endpoint. The indicator used in this experiment is phenolphthalein. Phenolphthalein, an organic compound, is colourless in acidic solution and pink in basic solution. This experiment involves two separate acid-base standardisation procedures. In the first standardisation the concentration of a sodium hydroxide solution (NaOH) will be determined by titrating a sample of potassium acid phthalate (KHP; HKC8H404) with the NaOH.This is the primary standard. 4) Researching Chemistry 9 Experiment A 0.128g sample of potassium acid phthalate (KHP, HKC8H4O4) required 28.5cm³ of NaOH solution to reach a phenolphthalein endpoint. Calculate the concentration of the NaOH. HKC8H4O4 → NaKC8H4O4 + H₂O mass GFM n of KHP- n of KHP = n of KHP=6.271 × 10-4mol 0.128 204.1 n of NaOH KHP:NaOH=1:1 n of NaOH = 6.271 × 10-4mol concentration of NaOH = moles volume in litres 6.271x10-4 0.0285 concentration of NaOH = concentration of NaOH = 0.0220 mol 1-¹ NaOH Use of Controls in Chemical Reactions The use of a control in chemical reactions validates a technique and may consist of carrying out a determination on a solution of known concentration. In the determination of the percentage of acetyl salicylic acid in commercial aspirin tablets a sample of pure aspirin (100% aspirin) would also be analysed to validate the techniques being used. In the determination of the vitamin C content in fruit juice a sample of ascorbic acid (pure vitamin C) would also be analysed again to validate the techniques being used in the determination and to give a referencing point on which to base all other results from impure samples. Complexometric Titrations Complexometric titration is a form of volumetric analysis in which the formation of a coloured complex is used to indicate the end point of a titration. Complexometric titrations are particularly useful for the determination of a mixture of different metal ions in solution. 4) Researching Chemistry EDTA (ethylenediaminetetraacetic acid) is a hexadentate ligand and an important complexometric reagent used to determine the concentration of metal ions in solution forming an octahedral complex with the metal 2+ ion in a 1 : 1 ratio. In particular it can be used to determine the concentration of nickel ions in a nickel salt. To carry out metal ion titrations using EDTA, it is almost always necessary to use a complexometric indicator to determine when the end point has been reached. 10 Murexide is used and, compared to its colour when it is attached to the Ni²+ ions, is a different colour when free. Murexide is a suitable indicator as it binds less strongly to the Ni²+ ions than the EDTA does and so is no longer attached to the Ni²+ ions at the end point of the titration where the colour is changed from yellow to blue- purple. Titrations would be carried out until concordant results were obtained. In this experiment ammonium chloride and ammonia solutions are used as a buffer to keep the pH constant as murexide is a pH dependent indicator. Titre H 1 2 3 EDTA H Initial burette reading (cm3) 1.4 20.5 15.3 4) Researching Chemistry H -H H H Example Approximately 2.6g of hydrated nickel sulfate (NiSO4.6H₂O) was weighed accurately and dissolved in a small volume of deionised water, transferred to a 100cm³ standard flask and made up to the mark with deionised water. 20cm³ samples of this solution were titrated with 0.112 mol 1-¹ EDTA solution using murexide as an indicator. The following results were obtained. Mass of weighing boat + NiSO4.6H₂O = 4.076g Mass of weighing boat after transferring NiSO4.6H₂O = 1.472g Final burette reading (cm3) 20.5 38.6 33.5 N Calculate the % of nickel in the hydrated nickel sulfate. Average titre volume = 18.15 cm³ Number of moles of EDTA used = C x V NjO EDTA complexed with nickel 2- Volume of EDTA added (cm) 19.1 18.1 18.2 11 Number of moles of EDTA used = 0.112 × 0.01815 Number of moles of EDTA used = 0.00203 mol Since EDTA complexes with nickel ions in a 1 : 1 ratio, the number of moles of Ni²+ in 20cm³ = 0.00203 mol. Therefore, in 100cm³, there are 0.00203 × 5 = 0.0102 mol Mass of nickel present = n x GFM Mass of nickel present = 0.0102 × 58.7 Mass of nickel present = 0.599g Mass of NiSO4.6H₂O = 4.076 - 1.472 Mass of NiSO4.6H₂O = 2.604g 0.599 % nickel = 2.604 % nickel = 23.0% x 100 Back Titrations and Associated Calculations A back titration is used to find the number of moles of a substance by reacting it with an excess volume of reactant of known concentration. The resulting mixture is then titrated to work out the number of moles of the reactant in excess. A back titration is useful when trying to work out the quantity of substance in an insoluble solid. From the initial number of moles of that reactant the number of moles used in the reaction can be determined, making it possible to work back to calculate the initial number of moles of the substance under test. An experiment which uses a back titration is in the determination of aspirin due to it being insoluble in water. A sample of aspirin of accurately known mass is treated with an excess of sodium hydroxide (the actual volume is known). The sodium hydroxide catalyses the hydrolysis of aspirin to ethanoic acid and salicylic acid and then neutralises these two acids. 4) Researching Chemistry 12 COOH + 2NaOH 1 2 Titre Starting volume (cm3) Rough 0.0 15.6 0.0 CH3 determination of aspirin COO Na* As an excess of sodium hydroxide is used the volume remaining is determined by titrating it against a standard solution of sulfuric acid. The difference between the initial and excess volumes of sodium hydroxide allows the mass of aspirin in the tablet to be determined. 4) Researching Chemistry The reaction taking place is: 2NaOH + H₂SO4 → Na2SO4 + 2H₂O Determination of Aspirin Three aspirin tablets of approximately 1.5g were added to a conical flask. Sodium hydroxide (25.0cm³ 1.00mol 1-¹) was pipetted into the flask along with 25cm³ water. The resulting mixture was simmered gently on a hot plate for approximately 30 minutes and after cooling transferred along with rinsings to a 250cm³ standard flask. The solution was made up to the graduation mark with water, stoppered and inverted several times to ensure through mixing of the contents. 25.0cm³ was pipetted into a conical flask; phenolphthalein indicator added and titrated against sulfuric acid (0.05mol 1-1). The end point of the titration was indicated by the colour change pink to colourless and they were repeated until concordant results were obtained (±0.1cm³). OH 15.1 Titration results Average titre volume = 15.15cm³ End volume (cm3) Titre volume (cm3) 15.6 15.6 30.8 CH₂ 15.2 15.1 Number of moles of sulfuric acid used = 0.01515 x 0.05 Number of moles of sulfuric acid used = 7.575 x 10-4 mol O Nat + H₂O 13 Number of moles of NaOH left in the 25.0cm³ hydrolysed solution = 2 × (7.575 × 10-4) Number of moles of NaOH left in the 25.0cm³ hydrolysed solution = 1.515 x 10-3 mol Number of moles of NaOH left in 250.0cm³ of the hydrolysed solution = 1.515 × 10-²mol Number of moles of NaOH added to aspirin initially = 0.025 × 1 = 2.5 × 10-²mol Number of moles NaOH reacted with aspirin = (2.5 × 10-²) – (1.515 × 10−²) Number of moles NaOH reacted with aspirin = 9.85 x 10-³ mol 2moles of NaOH reacts with 1mole aspirin therefore number of moles of aspirin in 3 tablets = = 4.925 x 10-³ mol (9.85x107 2 Number of moles of aspirin in 1 tablet = (4.925x10-3) 3 = 1.642 x 10-³ mol Mass of aspirin in each tablet = n x GFM Mass of aspirin in each tablet = (1.642 × 10-³) × 180 Mass of aspirin in each tablet = 0.296g = 296mg Practical Skills and Techniques Colorimetry Colorimetry uses the relationship between colour intensity of a solution and the concentration of the coloured species present to determine the concentration. A colorimeter consists essentially of a light source, a coloured filter, a light detector and a recorder. The filter colour is chosen as the complementary colour to that of the solution resulting in maximum absorbance. The light passes through the filter and then through the coloured solution and the difference in absorbance between the coloured solution and water is detected and noted as an absorbance value. 4) Researching Chemistry Preparing a series of standard solutions of appropriate concentration o Dissolve the accurately weighed substance in a small volume of water in a beaker. o Pour the solution into a standard flask. 14 o Rinse the beaker with distilled water and add the rinsings to the standard flask. o Add distilled water to the standard flask making the volume up to the mark. • Stopper the flask and invert several times to ensure thorough mixing. • Choosing an appropriate colour or wavelength of filter complementary to the colour of the species being tested • Using a blank • Preparing a calibration graph Colorimetry uses the relationship between colour intensity of a solution and the concentration of the coloured species present. A colorimeter or a spectrophotometer is used to measure the absorbance of light of a series of standard solutions, and this data is used to plot a calibration graph. The concentration of the solution being tested is determined from its absorbance and by referring to the calibration curve. The concentration of coloured species in the solution being tested must lie in the straight line section of the calibration graph. Absorbance 4) Researching Chemistry calibration curve Distillation Distillation is used for identification and purification of organic compounds. 15 Thermometer Boiling point temperature Still head Stirrer bar/ anti-bumping granules 4) Researching Chemistry Cooling water out Still pot Condenser Cooling water in Heating (oil/sand) bath Heat source Stirrer / heat plate Stirrer speed control Vacuum/ gas inlet Distilate/ receiving flask Cooling bath Heat control simple distillation apparatus The boiling point of a compound, determined by distillation, is one of the physical properties that can be used to confirm its identity. Distillation can be used to purify a compound by separating it from less volatile substances in the mixture. Possible experiments for distillation include: • Preparation of benzoic acid by hydrolysis of ethyl benzoate; Preparation of ethyl ethanoate; • Preparation of cyclohexene from cyclohexanol. Refluxing Refluxing is a technique used to apply heat energy to a chemical reaction mixture over an extended period of time. The liquid reaction mixture is placed in a round- bottomed flask along with anti-bumping granules with a condenser connected at the top. The flask is heated vigorously over the course of the chemical reaction; any vapours given off are immediately returned to the reaction vessel as liquids when they reach the condenser. 16 Dropping funnel ● Condenser Hirsch funnel Cooling water out 4) Researching Chemistry Cooling water in Possible experiments to use this technique include: Preparation of benzoic acid by hydrolysis of ethyl benzoate; • Preparation of ethyl ethanoate. refluxing Vacuum Filtration Vacuum filtration can be carried out using a Buchner, Hirsch or sintered glass funnel. These methods are carried out under reduced pressure and provide a faster means of separating the precipitate from the filtrate. The choice of filtering medium depends on the quantity and nature of the precipitate. Possible experiments include: Thermometer Preparation of potassium trioxolatoferrate (III); Preparation of aspirin; Preparation of benzoic acid by hydrolysis of ethyl benzoate; • Identification by derivative formation. Sinter funnel 17 Buchner flask Buchner funnel Hirsch funnels are generally used for smaller quantities of material. On top of the funnel part of both the Hirsch and Buchner funnels there is a cylinder with a fritted glass disc/perforated plate separating it from the funnel. A funnel with a fritted glass disc can be used immediately. For a funnel with a perforated plate, filtration material in the form of filter paper is placed on the plate, and the filter paper is moistened with a solvent to prevent initial leakage. The liquid to be filtered is poured into the cylinder and drawn through the perforated plate/fritted glass disc by vacuum suction. Hot filtration method is mainly used to separate solids from a hot solution. This is done in order to prevent crystal formation in the filter funnel and other apparatuses that comes in contact with the solution. As a result, the apparatus and the solution used are heated in order to prevent the rapid decrease in temperature which in turn, would lead to the crystallization of the solids in the funnel and hinder the filtration process. Recrystallisation Recrystallisation is a laboratory technique used to purify solids, based upon solubility. The solvent for recrystallisation must be carefully selected such that the impure compound is insoluble at lower temperatures, yet completely soluble at higher temperatures. The impure compound is dissolved gently in the minimum volume of hot solvent then filtered to remove insoluble impurities. The filtrate is allowed to cool slowly to force crystallisation. The more soluble impurities are left behind in the solvent. Recrystallisation can also be achieved where the pure compound is soluble in the hot solvent but not the cold solvent and the impurities are soluble in the hot and cold solvent. The impure compound is dissolved in the hot solvent. As it cools down the impurities stay in the solvent and can be filtered off. 4) Researching Chemistry 18 Selection of a suitable solvent is crucial to achieve a satisfactory recrystallisation. Neither the compound nor the impurities should react with the solvent. Other factors to consider will be the solubility of the compound and the impurities in the solvent, and the boiling point of the solvent (if the boiling point is too low, it will not be possible to heat the solvent to dissolve the impurities as it will evaporate off too easily). Possible experiments include: ● ● Preparation of benzoic acid by hydrolysis of ethyl benzoate; Preparation of potassium trioxalatoferrate (III); Preparation of acetylsalicylic acid. Use of a Separating Funnel Solvent extraction can be an application of the partition of a solute between two liquids. It is based on the relative solubility of a compound in two different immiscible liquids, usually water and an organic solvent. The partition coefficient is expressed as the concentration of a solute in the organic layer over that in the aqueous layer. The two solvents form two separate layers in the separating funnel and the lower layer is run off into one container and the upper layer is poured out into another container. The quantity of solute extracted depends on the partition coefficient and on the number of times that the process is repeated. Again, selection of a suitable solvent is crucial to achieve a high concentration of the solute after the extraction. The solute should be more soluble in the solvent than in the aqueous solution and it should not react with the solvent. Possible experiments include: ● Preparation of ethyl ethanoate; • Extraction of caffeine from tea. The largest risk when using a separating funnel is that of pressure build-up. Pressure accumulates during mixing if gas evolving reactions occur. This problem can be easily handled by simply opening the stopper at the top of the funnel routinely while mixing. This should be done with the top of the funnel pointed away from the body. When shaking, hold the stopper in place or it can become dislodged causing the liquids will spill. To account for this, simply hold the stopper in place with one hand. 4) Researching Chemistry 19 separating funnels Thin-Layer Chromatography (TLC) Thin-layer chromatography can be used to assess product purity. Instead of chromatography paper, thin-layer chromatography (TLC) uses a fine film of silica or aluminium oxide spread over glass or plastic. When setting up a TLC plate, a pencil line is drawn, usually 1 cm up from the bottom of the plate. The sample solution is then spotted on several times to get a concentrated spot on the plate. The plate is then placed in a suitable solvent (which acts as the mobile phase) making sure that the solvent is below the level of the spot. The solvent then travels up the plate and the components of the sample separate out according to their relative attractions to the stationary phase on the plate and the mobile phase (see the section on chromatography in 'Chemical equilibrium'). Retardation factor (Rf) values (distance travelled by compound/distance travelled by solvent) can be calculated and under similar conditions a compound will always have the same Rf value within experimental error. Rf values can be used to follow the course of a reaction by spotting a TLC plate with the authentic product, the authentic reactant and the reaction mixture at that point. Comparison of the Rf values of all three spots will allow the progress of the reaction to be determined. Since a pure substance will show up as only one spot on the developed chromatogram, TLC can be used to assess the purity of a product prepared in the lab. Possible experiments include: Preparation of aspirin; • Hydrolysis of ethyl benzoate. Melting Point and Mixed Melting Point 4) Researching Chemistry 20 The melting point of an organic compound is one of several physical properties by which it can be identified. A crystalline substance has a sharp melting point falling within a very small temperature range. Determination of the melting point can also give an indication of the purity of an organic compound, as the presence of impurities lowers the melting point and extends its melting temperature range. Since impurities lower the melting point, the technique of mixed melting point determination can be used as a means of identifying the product of a reaction. In this case the product can be mixed with a pure sample of the substance and a melting point taken of the mixture. If the melting point range is lowered and widened, it means that the two are different compounds. If the melting point stays the same it means that the two compounds are likely identical. The melting point of an organic solid can be determined by introducing a tiny amount into a small capillary tube and placing inside the melting point apparatus. A window in the apparatus allows you to determine when the sample melts in the capillary tube and the temperature can be determined from the thermometer. Pure samples usually have sharp melting points, for example 149.5 - 150 °C or 189 - 190 °C; impure samples of the same compounds melt at lower temperatures and over a wider range, for example 145 148 °C or 186 - 189 °C. Possible experiments include: • Preparation of benzoic acid by hydrolysis of ethyl benzoate. • Identification by derivative formation. • Preparation of aspirin. 4) Researching Chemistry 21