CHEM 1212L: Principles of Chemistry Laboratory II

1. Understand the relationship between accuracy and precision.
2. Explain and calculate standard deviation.
3. Perform procedures involving common lab equipment.
4. Recognize the linear relationship between parameters related by direct proportionality.
5. Use Excel to draw a graph from which results may be interpreted by maximizing the use of the graphing area while also choosing reasonable unit divisions for the x and y-axes.
6. Determine the slope from points on the best fit line on a graph to the correct precision.
7. Review and practice significant figures and scientific notation.
8. Use simple calorimetry equipment to determine the molar enthalpies of three different reactions.
9. Understand that in a closed system, the heat of a reaction is equal in magnitude but opposite in sign to change in heat of the surroundings.
10. Directly determine the molar enthalpy of dissolution of NaOH_(s) in water and use Hess’s Law to indirectly determine the thermodynamic value from two different reactions. Compare molar enthalpy results with literature values using a percent error calculation and compare the molar enthalpy of dissolution as determined by two different methods using a percent difference calculation.
11. Use mass and molarity as needed to mathematically determine the limiting reactant for each reaction so enthalpy per mole can be determined.
12. Conduct basic measurements, mathematical calculations, and conversions for measurement of volume, pressure, and temperature. Calculate percent error and account for sources of error in an experimental result.
13. Conduct laboratory experiments for a redox reaction by using equipment identifiable by name and by following safety procedures.
14. Perform chemical calculations using mass, mole, and volumes. Identify limitations of measuring devices to state the uncertainty/significant figures in measurements and calculation results.
15. Demonstrate a proficiency in relating the chemistry of gases to everyday phenomena.
16. Understand the linear relationship between light absorption and concentration for dilute solutions of dyes, as described by Beer’s Law.
17. Recognize that the amount of light absorbed by a dye is determined by wavelength and the electronic characteristics of each specific dye.  The wavelength at which the maximum light is absorbed by a particular dye is called the lambda max (_max).
18. Practice using a buret to deliver volumes to the precision of 0.02 mL.
19. Practice using the dilution formula to determine the concentrations of solutions diluted from a more concentrated “stock” solution.
20. Calibrate a spectrophotometer.
21. Use your spreadsheet graphing skills to obtain the equation for the linear relationship between absorption and concentration for your dye.
22. Using the equation for the best fit line for your data, determine the unknown concentration of a sample and the absorptivity, with correct units, of your dye.
23. Apply Beer’s Law to solve a real-world application.
24. Use slope and intercept from the equation of a line to determine enthalpy of vaporization, and entropy of vaporization, of water.
25. Calculate Gibbs free energy, of water.
26. Recognize the normal boiling point is the boiling point at 1 atmosphere and calculate that value for water.
27. Use percent error calculation to compare experimental values with literature values.
28. Calculate grams required to make a solution of a desired molarity.
29. Become proficient at making a solution using a volumetric flask.
30. Use mass of a solid along with volume data to determine the molarity of the solution.
31. Dilute a stock solution to a desired molarity.
32. Titrate a solution to determine the molarity of the solution.
33. Use the molarity of a known solution along with volume data and reaction stoichiometry to determine the molarity of a second solution.
34. Practice efficient experimental technique while collecting temperature and time data with the correct precision and best possible accuracy.
35. Use algebra and substitution to demonstrate how molar mass can be determined using the freezing point depression formula.
36. Expand your spreadsheet graphing skills by extracting lines with two different slopes from one data set.
37. Use the intersection of lines on a graph of decreasing temperature with time to determine freezing points of pure solvents and solutions and from those values determine the freezing point depression to the correct precision.
38. Determine molar mass from freezing point depression data.
39. Analyze technique to recognize sources of experimental error.
40. With the best possible precision and accuracy, use a variety of calibrated glassware to make dilute solutions with known concentrations of reactants.
41. Use color changes to monitor the rate of concentration decrease of a reactant.
42. Practice using the dilution formula to determine the concentrations of solutions diluted from a more concentrated “stock” solution.
43. Recognize that for a reaction for which the reactant coefficient is one, the initial rate of the reaction can be determined from the negative slope of a plot of reactant concentration near the start of a reaction as a function of time.
44. Determine the order for each reactant in a rate law by using a series of trials with different initial reactant concentrations and reaction rates.
45. Determine the rate constant for each trial from initial concentration, rate and order of each reactant.
46. Provide the overall rate law for the reaction studied.
47. Make several dilutions of stock solution and use the dilution formula to determine concentrations.
48. Calibrate a spectrophotometer and record percent transmittance (%T) and convert correctly to absorption.
49. Create a calibration curve by plotting absorption values vs known concentrations for the complex, Fe(SCN)²⁺.  Using the equation of the best fit line of the calibration data and absorption values of equilibrium solutions to determine their concentrations.
50. Create an ICE table from initial concentrations of reactants and equilibrium concentration of product to determine the equilibrium constant for the reaction.
51. Demonstrate your ability to use the quadratic formula in a sample calculation using the average equilibrium constant with initial concentration data.
52. Recognize that not all reactions go to completion and that often after a reaction, both products and reactants are present at equilibrium.
53. Learn to recognize and interpret the effects caused by a disturbance to a system at equilibrium, as explained by LeChatlier’s Principle. Disturbances can take the form of changes in concentration, both addition and removal; changes in pH; changes in temperature; changes in volume and pressure; etc.
54. Practice identifying the formulas of precipitates formed in double displacement ionic reactions.
55. Predict how product yields for chemical reactions might be improved using Le Chatelier’s principle.
56. Learn to calibrate and use a pH meter to record pH values to the precision of 0.01 pH unit during an acid-base titration.
57. Create a plot of pH vs volume of base added such that the pH can be read from the y-axis with a minimum precision of ± 0.05 pH unit and volume of base can be read from the x-axis with a minimum precision of ± 0.1 mL.
58. Using the appropriate technique outlined in the procedure, to a minimum precision of ± 0.05, determine the pKa of acetic and an unknown acid and compute Ka to the correct precision.
59. Determine the concentration of acid by using the molarity of the base and its volume to the equivalence point.
60. Use the quadratic formula, an ICE table, and the experimental Ka for the unknown acid to determine pH of a solution of given molarity.
61. Prepare a buffered solution with a desired pH from a weak acid or base and its salt.
62. Evaluate the specific pH for each solution after the addition of acid or base.