Stoichiometry of Iron-Copper

Stoichiometry of Iron-Copper Objective : To apply stoichiometric principles when experimenting the cycle of heating copper, poured it into the iron, cleaned the iron, and then used the respective masses in order to determine the mol ratio between copper and iron, to reach the conclusion of whether iron ion was Fe2+ or Fe3+. Introduction: Through this experiment we sought to use stoichiometric principles which derive from stoichiometry, the study of substance quantities, in ratio or proportion in a chemical reaction. In this lab experiment the principles help distinguish the equation between iron powder and aqueous copper (II) sulfate in the iron-copper (II) sulfate redox reaction. The iron displaces the copper (II) ions from the solution and is identified as more active in the solution. This process seen in the experiment is of the oxidation- reduction, redox reaction is defined with oxidation which is the gain of oxygen and loss of hydrogen while reduction is the loss of oxygen and gain of hydrogen. A redox reaction is a quite significant chemical reaction that occurs around everyday life, for example , in biological processes, metabolism , combustion, etc. In this experiment we see redox with the copper(II) reducing by gaining electrons from the more active metal iron, which is oxidized since it loses electrons to the copper ion and since iron forms in two ions, iron (II) and iron(III) we use this experiment to distinguish which one is formed. Two balanced reactions: 1. Fe(s) + Cu2+ (aq) ---- Fe2+ (aq) + Cu(s) : same number of moles Cu and Fe 2. 2 Fe(s) + 3 Cu2+ (aq) ---- 2 Fe3+ (aq) + 3 Cu(s) : Cu(s) will be 1.5 times the number of moles of Fe(s) Experimental: 1. First, Calculate the the volume of 1.0 M copper(II) sulfate solution required to completely react with 1.0 g of Fe(s) according to equation above. Then add 10% to that number. That is the volume of copper(II) sulfate solution . 2. Obtain directions from the instructor about where and how to safely get the iron powder and copper sulfate solution. 3. Locate an electronic balance scale and place empty weighing paper on the scale and tare the scale to 0.000 grams. 4. Add iron powder on the weighing paper until the scale reads 1.000g, make sure to record the weight to the exact 0.0001 weight on the scale. 5. Get a 150mL beaker from your equipment drawer and transfer the powder into it. 6. Get a graduated cylinder from equipment drawer and safely obtain 1.0M copper (II) sulfate solution based on the calculations made before the experiment 7. Take out Erhlenmeyer flask , tongs, and a wire guaze from equipment. 8. Pour the copper (II) sulfate solution into the Erhlenmeyer flask , and heat it to about 8090 degrees celsius using the hot plate. (BE CAREFUL TO NOT TOUCH IT) 9. After heating, use the tongs to hold the flask and slowly add the Copper (II) solution into the 150mL beaker containing the iron powder, watch out for extreme reaction. 10. Wait for the reaction to stop and the solution to reach ot about room temperature and allow the copper to settle. In notebook write down observations about the color of the liquid. 11. Using another beaker , use a dropper to take out all the decan liquid without removing any of the copper. 12. Now to clean, get 10-15mL of distilled water in your wash bottle and , swirl to clean and using the dropper remove this liquid as well. Repeat this twice. 13. After using the distilled water twice, get 5-10 mL to acetone and use that to clean and then decant into the waste. Repeat this twice. 14. Place the beaker with the copper metal cleaned, on the hot plate until its dry and free flowing. 15. Using weighing paper and directions from step 3, find the mass of the copper produced and record. 16. REPEAT STEP 1-15, again and record results for trial 2. Results / Discussion : Calculations before experiment: 1) Volume: (1g Fe * 1mol Fe/ 55.85gFe) * (3mol/ CuSO4)* (1LCuSO4/ 1molCuSO4) * ( 1mL CuSO4/10^-3LCuSO4) 2) 1.040(1 mol Fe/ 55.85g Fe)( 3mol CuSO4/ 2mol Fe)(1 L CuSO4/ 1 mol CuSO4) ( 1mL CuSO4/ 10^-3L CuSO4) 26.849 mL CuSO4 + 10 % = 31.2 = 30mL Trial 2: 3) 1.022(1 mol Fe/ 55.85g Fe)( 3mol CuSO4/ 2mol Fe)(1 L CuSO4/ 1 mol CuSO4) ( 1mL CuSO4/ 10^-3L CuSO4) 26.849 mL CuSO4 + 10 % = 29.54 = 30mL Calculating moles example: MASS/ Molar MASS Trial 1: 1.090 / 63.55= .0171 TABLE 1: Trial 1 Trial 2 1.040g 1.022g .0186 moles .0183 moles Mass of iron powder Moles of iron Trial 3 Volume of 1.0 M copper (II) sulfate used 30mL 30 mL 1.090g 1.043g .0171 moles .04 moles Mass of copper produced Moles of copper produced Conclusion/ Discussion Ratios: Ratio of moles of copper produced to moles of iron reacted. Trial 1 : .0171/.0186= 0.919 Trial 2: .0164/.0183= 0.896 Expected ratio: Fe(s) + Cu2+ (aq) ---- Fe2+ (aq) + Cu(s) : same number of moles Cu and Fe 2 Fe(s) + 3 Cu2+ (aq) ---- 2 Fe3+ (aq) + 3 Cu(s) : Cu(s) will be 1.5 times the number of moles of Fe(s) Based on these given equations we can assume that the results derived from the experiment match the ratios of equation 1 more as the ratios are more similarly matched with each other. Based on the ratio of (nCu / nFe) could be either 1.0 or 1.5, according to the results the reaction steers closer to 1.0 on the ratio calculated for both trials. Assuming that the copper mass received is exactly right and applying a 5% loss of copper we would get about 1.035 g of copper still for trial 1 and 0.990 for trial 2. This would still carry forward an approximate ratio of 1. Assuming a loss of 10 % in copper we would obtain Trial 1:0.981And Trial 2 : 0.938, both of these would carry out the same approximate ratio as initially obtained by the first equatio of 1 matching the Fe(s) + Cu2+ (aq) ---- Fe2+ (aq) + Cu(s) : same number of moles Cu and Fe Balanced reaction.