Determination of an Equilibrium Constant Lab Report, Lab Reports of Chemistry

Download Determination of an Equilibrium Constant Lab Report and more Lab Reports Chemistry in PDF only on Docsity! Name_________________________ Team Name ______________________ CHM112 Lab – Determination of an Equilibrium Constant – Grading Rubric Criteria Points possible Points earned Lab Performance Printed lab handout and rubric was brought to lab 3 Initial concentrations completed before coming to lab. 2 Safety and proper waste disposal procedures observed 2 Followed procedure correctly without depending too much on instructor or lab partner 3 Work space and glassware was cleaned up 1 Lab Report ICE tables and K calculations complete and shown in detail. 5 Question 1 1 Question 2 1 Question 3 1 Question 4 1 Total 20 Subject to additional penalties at the discretion of the instructor. Determination of an Equilibrium Constant Introduction It is frequently assumed that reactions go to completion, that all of the reactants are converted into products. Most chemical reactions do not go to completion because they are equilibrium systems where the reaction proceeds in both directions. As the reactants are used up, the rate of the forward reaction decreases. Conversely, as the concentrations of the products increase, the rate of the reverse reaction increases. Eventually, the rate of the forward reaction equals the rate of the reverse reaction and the concentrations of the reactants and the products stay constant. The system has reached a state of dynamic equilibrium. At equilibrium, both the forward and reverse reactions are occuring, but no net change is observed. Consider the general reaction: aA + bB ⇌ cC + dD where a,b,c and d are the stoiochiometric coefficicents. ( 1 ) Experimental evidence shows that the ratio of products to reactants (with each product and reactant expressed as a molar concentration and raised to its stoichiometric coefficient) is a constant for a reaction that has reached equilibrium. This constant, which is different for each chemical reaction, is known as the equilibrium constant and is designated with the letter K. There is a separate value of K for each temperature at which the reaction occurs. Thus, at equilibrium, the equilibrium constant K is equal to: K = [ ] [ ] [ ] [ ] C D A B c d a b ( 2 ) where the brackets [ ] imply molarity and the exponents are the stoichiometric coeffients of the balanced chemical equation. The equilibrium constant measures the extent to which a chemical reaction occurs. The larger the value for K, the greater the tendency for the reaction to go to completion is and the more products will be formed relative to the reactants. In this experiment you will determine the equilibrium constant for the following reaction: (Spectator ions are not shown.) Fe 3+ (aq) + HSCN (aq) ⇌ FeSCN 2+ (aq) + H + (aq) ( 3 ) K = [ ][ ] [ ][ ]HSCNFe HFeSCN + ++ 3 2 ( 4 ) Solutions of Fe3+ and HSCN will be mixed and will react to form some FeSCN 2+ and H+. The initial amounts of Fe3+ and HSCN can be calculated. The equilibrium concentration of FeSCN 2+ will be found using its spectroscopic properties – how much light it absorbs at a specific wavelength. FeSCN 2+ is a blood red complex that absorbs the blue- green wavelengths of visible light. Its absorbance is directly proportional to its concentration. The absorbance (a measure of the amount of light absorbed) will be measured by a spectrophotometer. Solutions to be measured are placed in cuvettes; these are sqaure tubes have minimal absorbance in the wavelength range of the spectrophotometer. 10. Measure the temperature of each solution. (Remember that K is temperature dependent.) Spectrophotometric measurement: 11. Obtain 4 cuvettes. Fill one (to the mark) with 0.20 M HNO3(aq). This will be used as the standard/blank. (All solutions have been prepared in 0.20 M HNO3(aq), and the absorbance of 0.20 M HNO3(aq) will be set to zero.) 12. Fill the other three other cuvettes to the line, each with one of the three test tube solutions. 13. Use Kimwipes to carefully wipe off any fingerprints and moisture from the cuvette. Fingerprints or other residue on the cuvette can affect the reading. 14. Open the lid of the spectrophotometer. If two sides of the cuvette are frosted, be sure that light will pass through the clear sides as you insert the cuvettes. Insert the blank in the B slot of the sample holder and insert the other samples in slots 1, 2, and 3 corresponding to the test tube they came from. Close the lid. 15. Press the the leftmost arrow key under the screen where it says “measure blank” to measure the baseline absorbance of 0.20 M HNO3(aq). 16. From the cell position keys, select the sample you want to measure (buttons labeled as 1, 2, and 3), then record the absorbance on your data sheet. Calculations: 1. Taking into account the dilutions in steps 1 – 9 of the procedure, calculate the initial concentrations of [Fe 3+] and [HSCN ] in each of the three test tubes. Show these calculations on the sheet titled ‘Calculation of Initial Concentrations’. Complete the calculations on this sheet before coming to lab. This will count towards your lab grade. Record these initial concentrations on the data sheet. This sheet must be handed in as directed by your instructor. 2. Using the absorbance values obtained and the value of ε, 4400, calculate the equilibrium concentrations of [FeSCN2+] in each cuvette. 3. Use the stoichiometric relationships in equation (3) and ICE chart to find the equilibrium concentrations of [Fe 3+] and [HSCN ]. 4. Plug these values into the equilibrium expression (4) and calculate K. Do this for all three solutions. Disposal The contents of all test tubes, volumetric flasks and beakers may be disposed of in the sink. Flush with a large volume of water Determination of an Equilibrium Constant: Initial Calculations Name ____________________________ Report Page 1 of 4 Calculate the initial concentration of Fe3+ and HSCN in the three test tubes before coming to lab. This will count towards your lab grade. Test Tube 1 Volume of stock Fe used 5.0 mL Concentration of Stock Fe solution 1.0 x 10-2 M Volume of stock HSCN used 5.0 mL Concentration of Stock HSCN solution 1.0 x 10-3 M Total volume after mixing Concentration of “Fe” after mixing (Use M1V1= M2V2) This is the initial concentration of Fe3+ in test tube 1. Record this number on data sheet. Concentration of HSCN after mixing (Use M1V1= M2V2) This is the initial concentration of HSCN in test tube 1. Record this number on data sheet. Test Tube 2 Volume of Fe D1 used: Concentration of Fe D1: Volume of stock HSCN used Concentration of stock HSCN solution Total volume after mixing Concentration of “Fe” after mixing Concentration of HSCN after mixing Test tube 3 Volume of Fe D2 used: Concentration of Fe D2: Volume of stock HSCN used Concentration of stock HSCN solution Total volume after mixing Concentration of “Fe” after mixing Concentration of HSCN after mixing Determination of an Equilibrium Constant: Data and Calculations Name ____________________________ Report Page 2 of 4 Note the initial concentrations from your calculations on this sheet before coming to lab Test Tube [Fe3+] initial [HSCN] initial 1 2 3 Test Tube Absorbance [FeSCN2+] eq (C = A/4400) 1 2 3 ICE Charts (Fill in charts and show all the calculations clearly) 1. From Test Tube 1 Fe3+ + HCSN ⇌ FeSCN2+ + H+ I C E K=