Hypothesis You are required to prove/disprove the hypothesis that: the use of titration can allow you to determine with high accuracy the number of cocalate ligands present in the iron coxnlate complex synthesised in part A of experiment 1. Introduction Marry metal ions, excopt Group 1A, the alkali metals, form an coxalate preoipitate at low oxalic acid concentrations. In human urine, the most abundarnt precipitating metal ion is calcium, which leacls to kidney stones. compound. it is the potassium sait ( \( \mathrm{K}^{*} \) ) of the (ili). \( \left[\mathrm{Fe}_{2}\left(\mathrm{C}_{2} \mathrm{O}_{4}\right)_{3}\right]^{3} \), in its crystalline form, the complex solidifes with three associated water molecules as a hydrate. This complex will be symthesised in two steps. The first step involves the conversion of ferrous ammonium hecalydrate, a pele green crystalline solid, to ferrous cocalate dinydrate, a yollow solid (Eq. 1 and Figure 1): \[ \begin{array}{l} \underset{\text { (ferrous ammonium }}{\mathrm{Fe}\left(\mathrm{NH}_{4}\right)_{2}\left(\mathrm{SO}_{4}\right)_{2} \cdot 6 \mathrm{H}_{2} \mathrm{O}(\mathrm{s})+\underset{\text { (oxalic acid) }}{\mathrm{H}_{2} \mathrm{C}_{2} \mathrm{O}_{4}} \longrightarrow \underset{\text { (ferrous oxalate dilwdrate) }}{\mathrm{FeC}_{2} \mathrm{O}_{4}-2 \mathrm{H}_{2} \mathrm{O}(\mathrm{s})}+\underset{\text { (ammonium sulfate) }}{\left(\mathrm{NH}_{4}\right)_{2} \mathrm{SO}_{4}}+\mathrm{H}_{2} \mathrm{SO}_{4}+4 \mathrm{H}_{2} \mathrm{O}} \\ \begin{array}{l} \text { (ferrous ammonium } \\ \text { sulfate hexahydrate) } \end{array} \\ \text { (ferrous oxalate dihydrate) (ammonium sulfate) } \\ \end{array} \] Part A: Standardisation of the permanganate solution 1) Thoroughly clean a burette, drain the wash solution and rinse water (once with tap and twice with deionised) through the burette tip. Rinse the burette twice with \( \sim 2 \mathrm{~mL} \) of the \( \mathrm{KMnO}_{4} \) solution and dispense through the burette tip. Fill the burette with \( \mathrm{KMnO}_{4} \) solution. Remove any air bubbles from the tip. Alow the solution to stand for \( \sim 30 \) seconds and read the volume to a precision of \( 0.02 \mathrm{~mL} \) (Failure to record and use the correct number of significant figures in this experiment will be penalised severely). The burette is a precision instrument and can give accurate volume readings if used correctly. Consult your laboratory instructor to make sure you are using and reading the burette correctly. 2) Your inboratory instructor will domonstrate the corroct use of the burotte. 3) Weigh approximately \( 0.11 \mathrm{~g} \mathrm{Na}_{2} \mathrm{C}_{2} \mathrm{O}_{4} \) using an Analytical Balance. Record the weight using all the significant permissible figures. 4) Transfer the sodium cxalate sample to a \( 125 \mathrm{~mL} \). Erlenmeyer flask and add \( -30 \mathrm{~mL} \) of deionised water and \( 5 \mathrm{~mL} \) of \( 2 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4} \). Swirl to dissolve the sample. Wash any remaining sodium oxalate from the glazed weighing paper using a small portion of deionised water from your wash bottle. 5) Heat the sample to about \( 60-70^{\circ} \mathrm{C} \) over a hotplate to speed up the reaction at the beginning. The reaction is later catalysed by the formation of Mn?+ which will holp maintain the reaction speed. Plosse don't boll the solution. Instead, romovo the flask from the hotplate as soon as the proviously mentioned tomporature is roachod. 6) While still warm, titrate the sample with the (-0.02 M) permanganate solution until a pink colour persists for about 30 seconds. Record the endpoint volume to a precision of \( 0.02 \mathrm{~mL} \). 7) Repeat with great care twice more and use the values estimated during thefirst titration. Use two new samples of sodium oxalate. 8) Calculate the concentration of the permanganate solution. Part B: Titration of potassium ferric oxalate trihydrate Transfer the sample to a \( 125 \mathrm{~mL} \). Crienmeyer flack. Add \( 30 \mathrm{~mL} \) deionised water and \( 5 \mathrm{~mL} 2 \mathrm{M} \mathrm{H}_{2} \mathrm{SO}_{4} \). Swirl to dissolve the sample. 2) Hoat the sample to about \( 60-70{ }^{\circ} \mathrm{C} \). 3) Titrate the sample to the pink endpoint. 4) Repeat this for two additional samples. 5) Calculate the percentage of oxalate in the sample. 6) Dispose of any excess permanganate solution in the appropriate waste container. Criteria Level 2 Title of practical task 0.1 points Properly summarise the work performed in this laboratory practice. It is OK to use the same title present in this manual.
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In this lab, you will reduce vanillin (4-hydroxy-3-methoxybenzaldehyde) with sodium borohydride (NaBH4) to produce vanillyl alcohol (4-hydroxy-3-methoxybenzyl alcohol). After isolating the product, you will determine its melting point, determine the experimental yield of the reaction, and calculate a percent yield for the reaction. Background: A common method for preparing alcohols is the reduction of aldehydes to produce primary alcohols (equation 1) or of ketones to produce secondary alcohols (equation 2). Some commonly used reducing agents are lithium aluminum hydride (LiAlH4), sodium borohydride (NaBH4), or catalytic hydrogenation. Vanillin is an aromatic compound that can be isolated from the cured fruit of Vanilla planifolia. It is mostly used in foods and perfumes. However, most commercial vanilla flavoring is derived from the lignin of wood pulp. We will be using vanillin as a starting material for the production of vanillyl alcohol, a compound that can serve as a valuable intermediate in the production of novel flavorings and perfumes. An excess of sodium borohydride, a convenient and mild reducing agent, will be used as the reductant in this reaction. The addition of HCl causes these four reactions to occur: 1. It hydrolyzes the O – B bond in the intermediate and protonates the oxygen atom that was originally the oxygen atom in the carbonyl group. 2. It destroys any excess NaBH4 that may be present in the reaction mixture. 3. It neutralizes any excess NaOH. 4. It protonates the phenolic oxygen. The following shows the detailed mechanistic sequence of events that lead to the final product, vanillyl alcohol. Procedure: 1. Place 3.8 g (25 mmol) of vanillin in a 125 mL Erlenmeyer flask followed by 30 mL of 1 M sodium hydroxide (NaOH). Add a stir bar. Place the flask on a stir plate and commence stirring at room temperature to solubilize the vanillin. If the vanillin does not dissolve after 5 minutes of continuous stirring, add another 5 mL of 1 M NaOH. 2. After the vanillin goes into solution, place the Erlenmeyer flask into an ice bath to cool the solution. Cool the solution to just below 15 °C. 3. In a small beaker, dissolve 0.5 g (12 mmol) of sodium borohydride (NaBH4) in 6 mL of 1 M NaOH. 4. Remove the vanillin solution from the ice bath. Using a glass pipet, slowly add the NaBH4 solution dropwise to the vanillin solution over a period of 5 minutes. [Note: This is an exothermic reaction, so be careful to add the NaBH4 solution slowly. Try to keep the temperature of the mixture within the range of 20 – 25 °C. If it is too cool, the reaction will proceed too slowly. If it is too warm, the product will not crystallize properly.] After the addition is complete, let the solution sit for 30 minutes while frequently stirring. 5. Place the reaction mixture in the ice bath to again cool the mixture. Cool the mixture to below 20 °C. While stirring with the stir bar, add 6 M hydrochloric acid (HCl) dropwise until the evolution of H2 gas stops (You are decomposing the excess NaBH4). [Note: This may require 6 – 10 mL of HCl.] At this point, the yellow color should have faded. Check the pH by removing a drop of the solution and adding it to a strip of pH paper to make sure the reaction is acidic. [Note: Do not dip the pH paper into the reaction mixture.] 6. Stir for an additional 15 minutes while cooling to allow the product to precipitate from the solution. If crystallization has not occurred, it may be necessary to scratch the bottom of the flask with a glass stirring rod or add a few seed crystals of vanillyl alcohol from the reagent bottle. 7. Weigh a watch glass and filter paper. Collect the product by vacuum filtration, using ice-cold water to aid the transfer of the material to the filter funnel. Wash the product twice with ice-cold water. Remove the filter paper and place it on the preweighed watch glass. Store the watch glass, filter paper, and your product in the drying cabinet until the following lab class. 8. After your product has dried, weigh it to calculate the experimental yield and determine the melting point of the product. Post Lab Questions: 1) Draw the structure of the product you would have obtained if: a) you had used NaBD4 as the reducing agent and then DCl in D2O to acidify (D = deuterium, an isotope of hydrogen) b) you had used NaBD4 as the reducing agent and then HCl in H2O to acidify c) Why is the reaction performed in basic solution? Why is this reaction not performed in acidic solution?
Krishna G.
This is from a lab: DETERMINATION OF THE IRON CONTENT OF A DIETARY SUPPLEMENT A.) Calculate the molarity of the standardized permanganate solution. B.) Calculate the mass content of Fe (in mg) per tablet. Procedure of lab: 1. Preparation of 0.002 M potassium permanganate: Use one of the large beakers for this preparation. Permanganate solutions are heated first to destroy reducible substances (traces of organics remaining in deionized water). After cooling, decant the supernatant to an amber bottle to get rid of the manganese dioxide impurities. If the solution is kept in the dark and is not acidified, its concentration will remain stable for several weeks. Procedure: Dissolve about 0.32 g of KMnO4 (F.W. 158.034 g/mol) in 1 L of deionized water using a large beaker. Cover that beaker with a watch-glass and heat to boiling using a hot plate in the hood. Keep the solution at a gentle boil for about 1 hr. Let the solution cool and transfer the supernatant of the solution to a clean amber glass-stoppered bottle (if you do not have it, use a regular glass bottle, then wrap it with aluminum foil); store in the dark when not in use. 2. Standardization of the potassium permanganate solution: Procedure: Weigh about three portions of 0.015 g dry sodium oxalate, Na2C2O4 samples (to the nearest 0.1 mg) and transfer a sample into a beaker or conical flask. Dissolve each sample in about 65 mL of 6M H2SO4. Fill a clean buret with the KMnO4 solution (make sure to rinse the buret three times with a few milliliters of the KMnO4 solution and dispose of the used KMnO4 in the appropriate waste container). Heat each solution to 80-90°C, and titrate with KMnO4 while stirring with a thermometer. The temperature should not drop below 60°C. The end point is marked by the appearance of a faint pink color that persists at least 30 s. Record the titration data, and calculate the concentration of the permanganate solution. Notes: 1. During the titration, promptly wash down any KMnO4 that spatters on the walls of the beaker into the bulk of the liquid using a wash bottle. 2. Finely divided MnO2 will form if the KMnO4 is added too rapidly and will cause the solution to acquire a faint brown discoloration. This is not a serious problem if sufficient oxalate remains to reduce the MnO2 to Mn2+; simply discontinue the titration until the brown color disappears. 3. The surface of the permanganate solution rather than the bottom of the meniscus can be used to measure titrant volumes if you encounter difficulty reading. 4. Partial decomposition of the permanganate to MnO2 may occur if it remains in the buret for a long time. Clean the buret with a dilute sodium bisulfite solution. 3. The Determination of Iron in a supplement tablet: Grind two iron supplement tablets to a fine powder by using a mortar and pestle as fine powder as possible, and transfer to a 250-mL beaker. Add about 50 mL of 1 M sulfuric acid solution and stir for approximately 30 min to dissolve it completely. Quantitatively transfer the resulting solution into a 100 mL volumetric flask and make up to volume with 1 M sulfuric acid solution (be careful not to overfill). Titrate three 10-mL aliquots of the sample solution with the potassium permanganate solution to the first faint pink that persists for 30 s. Do not add the KMnO4 rapidly at any time. Save the rest of the unknown sample solution (in the original 100 mL flask) for Lab 7. Part A was solved: The average Molarity calculated was 0.002099 M. Part B: Results of iron solution titrated (2 iron tablets, 100 mL of sulfuric acid): (Note: only 10 mL of this solution was used) Trial 1: Initial reading, mL: 23.80 mL Final reading, mL: 44.22 mL Vol. of titrant, mL: 20.42 mL Trial 2: Initial reading, mL: 4.85 mL Final reading, mL: 25.30 mL Vol. of titrant, mL: 20.45 mL Trial 3: Initial reading, mL: 25.30 mL Final reading, mL: 45.60 mL Vol. of titrant, mL: 20.30 mL Equation: 5Fe2+ + MnO4- + 8H+ → 5Fe3+ + Mn2+ + 4H2O According to the reaction equation given above, 1 mole of permanganate reacts with 5 moles of iron, and this ratio has to be used for titration result calculation.
Sri K.
Water Hardness Lab BACKGROUND: In some regions of the country, water is considered "hard." Hard water contains high levels of dissolved calcium and magnesium ions (Ca+2 & Mg+2). These alkaline earth metal cations precipitate (form a solid) with soaps, forming "soap scum." Since much of the soap precipitates and is thus removed from the water, there is less available to form lather. The amount of Ca2+ and Mg2+ cations can be determined using a chelating agent and an indicator. Chelating agents (also known as sequestering agents) are chemical compounds that react with metal ions to form a stable, water-soluble complex. We will use a method called titration to determine the Ca2+ and Mg2+ content of a standard and several unknown samples. The chelating agent we'll use is ethylenediaminetetraacetic acid (EDTA), and the indicator we'll use is calmagite. EDTA chelates divalent cations, removing them from solution. The calmagite indicator will change from a muddy green to a clear blue color when all of the Ca2+ and Mg2+ have been removed from the solution. MATERIALS: - 2.0 mM Na4EDTA - 0.50 mg/mL calcium carbonate (CaCO3) - 0.1% calmagite - pH 10 buffer - deionized water - 3 water samples - a 24-well plate - 3 test tubes - a test tube rack - disposable pipets PROCEDURE: Set-up 1. Obtain 3 clean, dry test tubes and label them "EDTA," "CaCO3," and "buffer." 2. Fill each ~1/3 full with the appropriate solution. 3. Obtain a new 24-well plate and several disposable pipets. Titration of the Color Standard and Reference Standard 1. Add 20 drops of distilled water to well A1. 2. Add 2 drops of pH 10 buffer then 1 drop of calmagite indicator (in that order). This is the color that you're aiming to achieve in the titrations. 3. Add 10 drops of 500 mg/L CaCO3 standard to wells A3-A5, then add 2 drops of pH 10 buffer and 1 drop of calmagite indicator (in that order). 4. Add EDTA solution, drop-wise, stirring briefly after each drop, until the sample turns the same blue color as the color standard. 5. Record the number of drops added to each well in Table 2. Titration of Unknown Samples 1. To wells B3-B5 add 10 drops of "unknown 1," add 10 drops of "unknown 2" to well C3-C5, and 10 drops of "unknown 3" to wells D3-D5. 2. Add 2 drops of pH 10 buffer and 1 drop of calmagite indicator (in that order) to each of the above 9 wells. 3. One well at a time, add EDTA solution, drop-wise, stirring or agitating briefly after each drop, until the sample turns the same blue color as the color standard. 4. Record the number of drops added to each well in Table 3. Equations: Eq 1: (0.5 mg/mL CaCO3 * Av drops EDTA to titrate unknown) / Av drops EDTA to titrate standard = mg/mL unknown Eq 2: (mg hardness of unknown * 1000 mL) / 1 mL unknown = mg hardness of unknown / 1 L unknown Table 1: Water Hardness Terminology Ca+2 concentration Description >300 mg/L Very Hard 150 to 300 mg/L Hard 50 to 150 mg/L Moderately Hard <50 mg/L Soft TABLE 2: RAW DATA - Record your data in the table below. Trial 1 Trial 2 Trial 3 mL of CaCO3 reference standard 0.5 0.5 0.5 drops of EDTA 30 30 20 Average drops of EDTA mL of unknown #1 0.5 0.5 0.5 drops of EDTA 31 17 17 Average drops of EDTA mL of unknown #2 0.5 0.5 0.5 drops of EDTA 30 30 37 Average drops of EDTA mL of unknown #3 0.5 0.5 0.5 drops of EDTA 35 40 33 Average drops of EDTA TABLE 3: CALCULATED DATA: Unknowns Unknown 1 Unknown 2 Unknown 3 Eq Av drops EDTA (from Table 1) mg hardness per 1 mL sample 1 mg hardness per 1 L sample 2 hardness rating (Ref table) LABORATORY QUESTIONS: Answer the following questions in complete sentences or complete thoughts. Show your work and report all answers with correct units. 1. Calcium carbonate (CaCO3), the standard used in this lab, is the major component of chalk, limestone, marble, and the shells of aquatic animals. Before being dissolved in water, what type of solid is solid CaCO3? 2. Do you think pond water would be more or less hard than bottled water? Explain your answer. 3. The molecular formula of ethylenediaminetetraacetic acid (EDTA) is C10H16N2O8. Calculate the molar mass. 4. Calcium carbonate will react with a strong acid according to the following balanced equation. Balance the equation. CaCO3 (s) + 2 HCl (aq) → CaCl2 (aq) + H2O (l) + CO2 (g) 5. Consider a sample of water ice - frozen H2O. What type of solid is water ice? What type or types of intermolecular force (IMF) is/are present in a sample of water ice?
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