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A Titration Experiment to Determine the Formula of a Copper (II) Hydroxide Polyatomic Ion Objective: In this titration experiment you will react sodium hydroxide with a copper (II) salt. The reaction will create a polyatomic ion containing both the copper (II) ion(s) and the hydroxide ion(s). There are three possible copper (II) salts available, and you will have one assigned to each of you. From the titration, you will determine the mole ratio between the copper (II) ion and the hydroxide ion, and then determine what cation or anion the polyatomic ion matches with in order to make a neutral compound, that is insoluble and precipitates out of solution. For your pre-lab: Read through this handout and 1) outline/summarize the procedure as usual and 2) and make a data table for this experiment in your lab notebook. [I will not accept photocopied tables, but typed tables are ok.] The point of the pre-lab is to read through the lab, and figure out what measurements you will need to take. I will be marking off for putting down RESULTS instead of DATA in your data tables. I will also be marking off if you forget to include some of the DATA you will need. 3) Review the Solution and Dilution and Titration learning labs to remind yourself of the proper way to prepare your glassware and techniques for making solutions and doing titrations. Outline this information as part of your procedure. 4) You will need to make approximately a 0.1 M solution of the copper (II) salt. Set up this calculation in your notebook ahead of time, as part of your pre-lab. Leave space in your calculation for the molar mass of the salt, which you will be able to write in, once you are assigned a salt.

During the Lab This lab will be done individually. When you have received your copper (II) salt, write down the formula and its molar mass. If your salt is a hydrated salt, your molar mass includes the water molecules; however, when writing out your chemical reaction, omit the water molecules. So for Cu(NO3)2 · 21/2 H2O, all you need to include as a reactant is Cu(NO3)2. Put the molar mass into your set up calculation, and figure out how much you need to weigh out (approximately) to make a 0.1 M solution. Use the correct number of significant figures in your calculations. If your calculated mass is different from the actual mass that you weighed, redo the calculation so you know just what you have in your volumetric flask. Label the flask with colored tape; the label should include the concentration, the name of the salt, your name, and the date. This titration is a little different than most titrations; you know the concentration of both solutions ahead of time! That is because you need to figure out moles of Cu2+ that have combined with moles of OH-, and since the ratio results in a polyatomic ion, rather than a neutral compound, you need concentrations of both solutions.


To prepare your titration sample you always need three things, the original solution being analyzed, the indicator (remember what happens when you forget to add the indicator?), and some distilled water. After preparing your glassware appropriately, pipet about 20 mLs of your copper (II) salt solution into a 250 mL Erlenmeyer flask, add about 50 mL of DI water, and 10 drops of the phenol red-Nile blue indicator. Mix thoroughly! Once the buret has been prepared, fill it with the NaOH solution. Record the concentration of the NaOH solution. Titrate. This is a very pretty titration; but the colors make things a bit confusing. Your precipitate will be one color, and your solution above the precipitate will be a different color. As you titrate, you need to swirl the solution regularly, but as you are waiting for your indicator to turn the solution yet another color, you need to look at the color of the solution above the precipitate, so keep that in mind. The endpoint occurs when the solution is a homogeneous violet color. Concentrate on how YOUR solution color changes as you titrate. To do this, you need to titrate slowly, giving your precipitate time to settle (after you swirl) so you can see what color you really have, and how it is changing. The indicator will turn color when all of the Copper ions have been used up and have become part of the precipitate. You will probably need to do four titrations to obtain three good titrations. After you have completed your three good titrations, put the chemicals in the appropriate waste containers, and rinse /wash all of your glassware. Your buret is the buret you will use for the Lab Practical, so make sure it is cleaned well! This is the last titration you will do before the Lab Practical. The practical is to show (me) that you can make a solution and then titrate it, just like in this lab, and then you will be expected to complete all of the calculations necessary to find the concentration of the unknown before you leave lab. The lab practical is an exam, and so you have to do that experiment and the calculations silently, as you are being tested. This is the quietest lab day ever! Make sure you have any questions about titrations answered before or during this lab, and make sure that you understand the calculations involved in a regular titration. The problems at the end of chapter 5 combined with the information given in the two learning labs are excellent preparation for this exam. Lab Report ­ This is a Formal Lab Report · Objective, as usual. · Procedure, as usual. · Observations, as usual. · Reactions: Figure out the molecular reaction that has occurred. Hints: 1) Consider the conservation of mass and 2) there are only two products. You need to figure out the mole ratio of your product (the polyatomic ion) AND figure out what cation or anion from the reactants could combine with the polyatomic ion to make it neutral and allow it to precipitate. Then, write the molecular reaction, full ionic reaction, and net ionic


reaction that occurred during this titration. Each of the three chemical equations should be balanced, and include the states. · Data Table, as usual. · Calculations, written as usual. o Show the calculation of the actual solution that you made. o Calculate the number of moles of Cu2+ used in forming the precipitate. o Calculate the number of moles of OH- used in forming the precipitate. o Calculate the ratio: moles of OH- / moles of Cu2+ (this is as simple as it looks), find the number, and as if it were an empirical formula problem (look this up if you forget), the numbers should be whole numbers. All calculations should be done to the correct number of sig figs, THEN round to whole numbers at the end. o Make sure to show the rounding to whole numbers step. o Write the formula of the polyatomic ion in the following format: CuxOHy and what do all IONS have???? Make sure to include it! o Calculate the average ratio, as usual o Calculate the average deviation, as usual. · Results Table, as usual. · No graphing section. · Discussion, as usual. · Conclusion, as usual.



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