Inorganic Chemistry Experiments PDF

Preparation of [NiII(NH3)6]Cl2 and its analysis by complexometric titration and spectrophotometry  Synthesis of Hexamminenickel(Il) chloride 1. Transfer 10 mL solution of [NiII(OH2)6]Cl2 (contains 6 g of NiCl2) in a 250 mL beaker. 2. Take 15 mL solution of aqueous ammonia in a...

Preparation of [NiII(NH3)6]Cl2 and its analysis by complexometric titration and spectrophotometry  Synthesis of Hexamminenickel(Il) chloride 1. Transfer 10 mL solution of [NiII(OH2)6]Cl2 (contains 6 g of NiCl2) in a 250 mL beaker. 2. Take 15 mL solution of aqueous ammonia in a measuring cylinder. 3. Add the ammonia solution drop wise to the solution of nickel chloride with constant stirring till the colour of the solution has changed from pale green to intense violet. 4. Allow the solution to stand at room temperature for 5 minutes, cover with watch glass. Then cool it in an ice bath for about 15 minutes. 5. Filter the solution and wash the crystals with 3-5 mL ammonia solution. 6. Dry the crystals using filter paper. 7. Report the weight and yield of the dried complex. calculate the %yield  Spectrochemical Series An arrangement of ligands according to their increasing ability to split the d- orbitals is termed as the spectrochemical series. This splitting is quantified using the crystal field splitting parameter (Δ). The parameter is determined experimentally. Weak Field I-  Br- S2- SCN- Cl- NO3- F-  C2O42- H2O NCS- CH3CN  NH3 en  bipy phen NO2- PPh3 CN- CO Strong Field d orbital splitting for an octahedral geometry The splitting parameter ∆o can be related to the color of a complex  Spectrochemical Series 1.0 en > NH3 > H2O 0.8 Absorbance 0.6 0.4 0.2 0.0 300 400 500 600 700 Wavelength (nm) λmax in cm-1 [Ni(OH2)6]2+ [Ni(NH3)6]2+ [Ni(en)3]2+ Band 1 25300 28200 29000 Band 2 13800 17500 18650 Beer-Lambert Law for quantitative analysis  I1  A = − log(T ) T =    Io  A=ɛCl A = absorbance ɛ = molar absorptivity (L/mole/cm) l = cell pathlength (cm) C = concentration of analyte (mol/L) 5 cm Increasing conc. 0.5 cm 1 cm 5 Estimation Ni complex by spectrophotometry Several solutions (of NiCl2) of known concentration and one solution of Find the unknown concentration will be conc of provided to you. unknown 1. Measure the absorbance of all the sample solutions at 395 nm using a UV- Visible spectrophotometer. 2. Plot absorbance versus mg/mL of nickel. Determine the concentration of nickel present in the unknown solution in g/L. S. Conc. of Conc. of Abs. at No. Solution Solution 395 nm (M) (mg/mL) 1 0.01 0.585 2 0.04 2.34 3 0.05 2.925 4 Unknown Unknown An overlay of UV/Vis absorption spectra of Nickel solutions with concentrations: 0.01 M, 0.04 M, 0.05 5 0.06 3.51 M, 0.06 M and an unknown solution  Complexometric titration with EDTA Ni2+ + EDTA4- [Ni(EDTA)]2- Ni Indicator used is Murexide Color change for Ni: Yellow-green to violet https://chem.hbcse.tifr.res.in/  Complexometric titration with EDTA 1. Take 80 mL of 0.05 M EDTA solution in a 250/500 ml plastic beaker and fill it in a clean burette up to the mark. 2. Weigh accurately 1.15 g of [Ni(NH3)6]Cl2 complex and transfer this to a 100 mL volumetric flask. Now add 50 mL of 1 N H2SO4 to dissolve it and makeup the solution to the mark with distilled water. 3. Pipette out 10 mL of the complex solution in a 250 mL conical flask and dilute it with 15 mL of distilled water. 4. Add 2-3 drops of murexide indicator and 5 mL NH4Cl solution (0.5 M) to the conical flask. Now add ammonia solution (7-10 drops) to maintain a pH 7 (light green color of the solution). 5. Titrate it with EDTA solution till the endpoint is near, add 3 ml of ammonia solution and continue the titration till the endpoint (bluish violet color appears). 6. Repeat the titration and get concordant values. 7. Calculate the amount of Ni present in the complex.  Estimation of Nickel(ll) by EDTA Concentration of EDTA solution = 0.05 M S. No. Volume of EDTA used from the burette (X) 1 10.0 mL 2 10.1 mL 3 10.1 mL 1 mole Ni reacts with 1 mole of EDTA to form of the Ni-EDTA complex. Moles of Ni = Moles of EDTA (Molarity x Volume) of Ni = (Molarity x Volume) of EDTA Molarity of Ni x 10 mL = 0.05 M x Volume of EDTA (Burette Reading) Suppose Burette Reading is X mL (use concordant value) Molarity of Ni = (0.05 x X)/10 = Y M Thus, the solution contains Y moles/Liter of Nickel Find X, Y and Z Grams/Liter of Ni = Y x 58.69 (Atomic Weight of Ni) = Z g/L Hence, in 100 mL of solution Z g of Ni is present. We started with 1.15 g of [Ni(NH3)6]Cl2 , thus Z g of Ni is present per gram of [Ni(NH3)6]Cl2  The lab report 1. Synthesis of Hexamminenickel(Il) chloride Experimental yield is 3.41 g, report the %yield in the lab report S. Conc. of Conc. of Abs. at 2. Estimation Ni complex No. Solution Solution 395 nm by spectrophotometry (M) (mg/mL) Plot the absorbance at 395 1 0.01 0.585 nm vs. concentration of Ni 2 0.04 2.34 solution (normal paper 3 0.05 2.925 is also fine) 4 Unknown Unknown 5 0.06 3.51 3. Estimation of nickel(ll) by EDTA Find X, Y and Z in the previous slide and report the calculations in the lab record  Conclusions [NiII(NH3)6]Cl2 was synthesized from [NiII(H2O)6]Cl2 The synthesized complex was then used to estimate the amount of Nickel present Ni2+ + EDTA4- [Ni(EDTA)]2- Concentration of a given solution was determined using solutions of known concentration with the help of UV/Vis absorption spectroscopy 1.0 en > NH3 > H2O 0.8 Absorbance 0.6 0.4 0.2 0.0 300 400 500 600 700 Wavelength (nm)  Estimation of Iodine in Iodized Common Salt Iodine is an essential element for life and one of the heaviest elements required by living organisms. However, around 1/3 of the world’s population lives in areas of iodine deficiency. The practice of adding iodine to salt is a safe, easy and effective way of overcoming iodine deficiency in our diet. Globally two chemical forms of iodine are used for iodization; Iodates (IO3-) and Iodides (I-). The iodides degrade more readily in presence of impurities, exposure to sunlight, moisture and exposure to heat, whereas the iodates remain stable under extremes of weather and handling. USA uses potassium iodide (77 mg/Kg) while Germany and India use potassium iodate (25-20 mg/Kg) for iodine fortification. www.consumeraffairs.in, www.hrt.org  Theoretical Background The direct iodometric titration method called as iodimetry, refers to titrations with a standard solution of iodine. The indirect iodometric titration method, termed iodometry, deals with the titration of iodine liberated in chemical reactions. Experiment involves three steps: Determining whether iodate or iodide anion is present in the given salt sample Standardising a solution of sodium thiosulphate Estimating the amount of iodine in given by performing an iodometric titration using the standardized sodium thiosulphate solution Iodate anion IO3-  Theoretical Background Estimation of Iodine 2 Na2S2O3 + I2 Na2S4O6 + 2 NaI Test for Iodate: Iodate in presence of free hydrogen ion, oxidizes added iodide to give free iodine; which turns starch blue IO3- + 5 I- + 6 H+ 3 I2 + 3 H2O Test for Iodide Iodide is oxidized to free iodine with an acidic solution of sodium nitrite. The free iodine turns starch blue 2 NaNO2 + H2SO4 2HNO2 + Na2SO4 2 HNO2 + 2 I- I2 + 2 NO + H2O Starch used as an indicator for titrations with iodine  Theoretical Background Indicator : Starch Starch-iodine complex has a strong colour; blue-purple-brown-grey-black. Starch binds very strongly to iodine (binds to the I3- ion in solution). We add starch when the iodine concentration is a little less because the starch- iodide complex has less solubility and can stop iodine from reacting completely with thiosulphate. Some of the iodine may bind to copper in solution To make sure that all iodine reacts with sodium thiosulphate, we add KSCN. KSCN binds to copper thereby releasing iodine and making it available to react. Solution containing Solution containing Iodine Starch-Iodine complex  Experimental Protocol Test for iodate solution: Moisten a pinch of salt with 2-3 drops of the given solution (mixture of starch (A), KI (D) and (HCl) E). If iodate is present the salt will turn blue/grey and the color will be retained for several minutes before turning brown. Test for iodide solution: Moisten a pinch of salt with 2-3 drops of the given solution (mixture of starch (A), NaNO2 (B) and H2SO4 (C)). If iodide is present the salt will turn blue and remain blue for several minutes before fading. Determination of Iodate Content (if test A is positive) 1. Weigh X g of the given salt sample and transfer into a 250 mL conical flask and dissolve it in 50 mL water. 2. Add 1 mL 2 N H2SO4 (use dropper, do not pipette by mouth), then add 5 mL of 10% KI solution using a measuring cylinder. The solution will turn yellow. Wrap the mouth of conical flask with a piece of filter paper and keep it in cupboard for 10 minutes. 3. Take 60-80 mL solution of (approx.) 0.005 M Na2S2O3 in a 250/500 mL plastic beaker and use for titration. Rinse burette and fill it in and adjust zero level. 4. Remove flask from cupboard and titrate with Na2S2O3 solution until the solution turns pale yellow. Now add approx. 10 drops of starch indicator. The solution will turn dark purple. Continue titrating until the solution becomes colorless. 5. Record the volume of titrant (Na2S2O3 solution) used and calculate the amount of iodine present in part per million (ppm).  Theoretical Background Standard Solution: A solution for which the concentration is known accurately. The process of determining the exact concentration of a solution is termed as standardization. In a titration, it is critical to know the exact concentration of the titrant (the solution in the burette which will be added to the unknown) in order to determine the concentration of solutions being tested. Standardization of Sodium Thiosulfate Sodium thiosulphate is titrated against a standard solution of copper sulphate. 2 Cu2+ + 4 I- Cu2I2 + I2 I2 + 2 S2O32- S4O62- + 2 I- Net Reaction: 2 Cu2+ + 2 I- + 2 S2O32- Cu2I2 + S4O62-  Experimental Protocol Standardization of Sodium Thiosulfate 1. Pipette out 10 mL copper sulfate of concentration 0.005 M in a conical flask and add 5 mL of 5% KI solution. The solution will turn yellow in color. 2. Titrate with Na2S2O3 solution until the solution turns pale yellow. Add 7-8 drops starch indicator solution at this stage. 3. Continue the titration until the purple color fades, then add 5-6 drops of KSCN solution and titrate again. The end point gives a colorless solution. CAUTION: To ensure that you have obtained the true end point, stir the flask for 20 seconds and then wait for 20 seconds to make sure that the purple color does not reappear. 4. Repeat the titration to get concordant readings. Calculate the molarity of the given sodium thiosulfate solution.  Observations and Calculations Standardization of Sodium Thiosulfate S. No. Volume of sodium thiosulphate used from the burette 1 9.4 mL 2 9.5 mL 3 9.5 mL Given concentration of copper sulphate = 0.005 M Equivalents of Copper = Equivalents of Thiosulphate (Normality x Volume) of Copper = (Normality x Volume) of Thiosulphate 0.005 x 10 mL = Normality of Thiosulphate x Volume of Thiosulphate (Burette Reading) Suppose Burette Reading is 9.5 mL Normality of Thiosulphate = (0.005 x 10)/9.5 Normality of Thiosulphate = 0.0053 M  Observations and Calculations Estimation of Iodine S. No. Volume of sodium thiosulphate used from the burette 1 2.4 mL 2 2.5 mL 3 2.5 mL IO3- + 6S2O32- + 6H+ 3S4O62- + I- + 3H2O Consider the above reaction of iodate ions present in salt sample reacting with sodium thiosulphate 6 x M1V1 Iodate = M2V2 Thiosulphate 6 x M1 x 50 = 0.0053 x Burette Reading Value obtained by standardizing sodium thiosulphate against a standard copper solution Suppose Burette Reading is 2.5 mL 6 x M1 x 50 = 0.0053 x 2.5 Molarity of Iodate solution = 4.4 exp -6  Observations and Calculations Thus, the solution contains 4.4 exp -6 moles of Iodine per Liter of solution The solution contains 4.4 exp -6 x 126.90 grams Iodine per Liter of solution Atomic Weight of Iodine = 126.90 g Solution contains 5.6 exp -3 grams Iodine per Liter of solution = 5.6 mg of Iodine per Liter of solution Concentration in ppm To report in ppm Suppose we started with 10 g of salt, 10 g salt was dissolved in 50 mL of water In 50 ml of solution; (5.6 mg x 50)/ 1000 = 0.28 mg of Iodine Thus, 0.28 mg of Iodine in 10 g of salt ppm means parts per million; which is mg per Kg (0.28/10) x 1000 = 28 ppm of Iodine was present in our salt sample  Results Iodate anions were identified in the given salt sample Concentration of iodate ions and iodine was determined by performing an iodometric titration Concentration of Iodine in given salt sample = 28 ppm  Cyanotype Blue Printing Cyanotype is a photographic printing process that produces a cyan-blue print. Engineers used the process well into the 20th century as a simple and low-cost process to produce copies of drawings, referred to as blueprints.  Cyano Type Blue Printing It was invented by Sir John Herschel in 1842. The process of blue printing was eminently suited to its traditional role in reproducing technical drawings, its most common use in engineering and architecture until the advent of modern photocopiers. Sir John Herschel 1792-1871  Principle of Blue Printing There are two distinct reactions involved. Ferric (Fe3+) ions in some complexes such as ammonium ferric citrate and ammonium ferric oxalate are reduced by light to ferrous (Fe2+) ions. The ferrous ions are then reacted with potassium ferricyanide to form an insoluble blue compound; Turnbull’s Blue. Turnbull’s Blue is very much like Prussian Blue, the only structural difference being that the position of ferrous and ferric ions are reversed in the cyanide lattice. It should be noted that, once Turnbull’s blue forms, it immediately converts to Prussian Blue – Fe4[Fe(CN)6]3.15 H2O  Prussian Blue The [Fe(CN)6]4- anion in Prussian blue is octahedral. The right-hand drawing shows its unit cell, which has a cubic lattice structure. The name Prussian blue originated in the 18th century, when the compound was used to dye the uniform coats for the Prussian army.  Prussian Blue Prussian blue is a pigment that is used to color paints, inks, textiles, and other commercial products. But during the past decade, Prussian blue found a high-tech use: Its ability to transfer electrons efficiently makes it an ideal substance for use in sodium-ion battery electrodes. Battery producer Natron Energy (Santa Clara, CA) recently concluded a deal with Lonza Group (Basel, Switzerland) to supply 700–1000 t/year of Prussian blue for the Natron BlueTray 4000 battery system to be used for data-center and telecommunications applications. https://www.acs.org/content/acs/en/molecule-of-the-week/archive/p/prussian-blue.html https://cen.acs.org/materials/energy-storage/Natron-picks-Lonza-Prussian-blue/99/i14  Experimental Protocol 1. Immerse pieces of bond paper one by one in the sensitizing solution and keep them immersed for 4-5 minutes. 2. Remove the wet pieces of paper and place them between sheets of filter paper. This should be done as quickly as possible and in a partially closed locker. Dry it for 10-15 minutes. 3. After the paper has dried, place an opaque object on top of the sensitizing paper, compress it between sheets of glass and expose to sunlight for 4-6 minutes. 4. Make 3-4 exposures, varying the time of exposure to optimize the best condition. 5. After the exposure, dip the paper into 0.1 M ferric cyanide. It is important that the paper is immersed all at once, otherwise lines will appear on the blue field of the paper. 6. Remove the paper and dip it in 0.3 M potassium dichromate solution for one minute. Afterwards, wash the paper first with 0.1 M HCl and then tap water and dry.  Reactions Involved 2 (NH4)3FeIII(C2O4)3 2 CO2 + 2 FeII(C2O4) + 3 (NH4)2(C2O4) (UV Light) FeII(C2O4) + K3FeIII(CN)6 KFeIII[FeII(CN)6] + K2(C2O4) 8 KFeIII[FeII(CN)6] 2 FeIII4[FeII(CN)6]3 + 2 K4FeII(CN)6 In our experiment, diammonium phosphate is added to the sensitizing solution to decrease its sensitivity so that the experiment may be performed in diffused light, rather than a dark room. Potassium dichromate helps to improve the contrast of the image.