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Empirical Formula of a Hydrate

Empirical Formula of a Hydrate. In addition to this presentation, before coming to lab or attempting the prelab quiz you must also:

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Empirical Formula of a Hydrate

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  1. Empirical Formula of a Hydrate • In addition to this presentation, before coming to lab or attempting the prelab quiz you must also: • Read the following pages in the “Laboratory Handbook…” by Stanitski et al. before coming to lab: pages 39-40 (using a bunsen burner) page 42 (heating a solid) • Watch the prelab video for this experiment • Review the video on use of the balances if needed • Read the introduction to the lab in the coursepack

  2. What’s the point? • In this lab you will • Gain more practical experience - bunsen burners, heating apparatus, balances • Determine formula stoichiometry by experiment- find the molar ratio of water to salt in a hydrate • Practice empirical formula calculations

  3. Hydrates • Ionic compounds with water molecules incorporated into their crystal structures • Waters are shown in the chemical formula • The number of moles of water (“waters of hydration”) can be anywhere from 1 to 12 • Example formulas are: • ZnSO47H2O (7 mol H2O for every for 1 mol of ZnSO4) • CuSO45H2O (5 mol H2O for every for 1 mol of CuSO4)

  4. Naming Hydrates To name a hydrate, name the ionic compound, then indicate the amount of water of hydration by adding… …a greek prefix to indicate the number of waters… 1 = mono; 2 = di; 3 = tri; 4 = tetra; 5 = penta; etc... …and then the word hydrate e.g., CuSO45H2O = copper (II) sulfate pentahydrate Note: There is no way to predict the number of waters of hydration for a given ionic compound This must be obtained from experiment

  5. Measuring the Amount of Water • The hydrate molar mass includes waters of hydration ZnSO47H2O would be, • 65.37 + 32.06 + (4 x 16.00) + (7 x [2 x 1.008 + 16.00]) = 287.57 g / mol • The moles of water in a hydrate can be determined by measuring the mass of water lost upon heating (a) Original mass = mass of salt + water (b) Mass after heating = mass of salt (c) The difference gives the mass of water (d) To get the formula, convert masses to moles

  6. Bunsen Burners Open flame heat source Air mixes with gas here.The collar is adjustable to let more or less air in. Gas goes in here Use this valve to adjust gas flow Don’t adjust from nozzle on bench

  7. To use a Bunsen burner • Connect rubber hose to the lab gas valve and burner, checking for holes or cracks in the rubber. • Turn on the main gas valve (on bench). • Light burner with striker. • Adjust the flame as desired by turning the needle valve (under burner tube) and/or adjusting the collar height. • When finished heating, turn off the burner at the main valve (on bench). • Never try to turn off the burner from the needle valve.

  8. A hot blue flame can be obtained by raising the Bunsen Burner collar to allow more air in A cooler, yellow flame can be obtained by lowering the Bunsen Burner collar to allow less air in This type of flame is good if you need a low temperature, but is harder to control This is generally an easier type of flame to control Less Air “reducing” More Air “oxidizing”

  9. Experimental Details • Measure mass of hydrate sample • Heat in crucible to remove water • Cool the crucible in a dry location, the desicoolerkeeps salt from absorbing atmospheric water vapor • Measure mass of water lost You must heat the sample until no more mass is lost This is to ensure all water is removed Verify by several heating steps until constant mass • Measure mass of dry salt • From mass of water and mass of salt, determine molar ratio of water/salt

  10. Clay triangle, on iron ring, to support the crucible Iron ring, clamped to post Post on ceramic stand Bunsen burner, Adjusted to hot, blue flame Experimental Set-Up

  11. Use of Balances (Review) • Use Analytical Balances for All Measurements In This Lab • precise to +/- 0.001 g (1 mg) • RULES • never weigh a hot/warm object on balance • keep the balance area clean

  12. Necessary Calculations • calculating hydrate empirical formulas • moles of salt = grams dry salt / molar mass • moles of water = grams water lost / molar mass Heating removes 0.438 g of H2O from a hydrate, leaving behind 0.562 g dry salt. If the salt has a molar mass of 161.26 g mol-1, what is the hydrate empirical formula?

  13. salt mass = 0.562 g • water mass = 0.438 g 0.562 g = 0.00349 mol salt 161.26 g/mole 0.438 g = 0.0243 mol H2O 18.02 g/mole • Divide the larger number by the smaller: (0.0243 mol H2O) / (0.00349 mol salt) = 6.96

  14. This means, 6.96 mol H2O per mol salt It would seem safe to say that there are 7 mol water for each mol salt • the empirical formula, written in the appropriate form for a hydrate, isSalt7H2O • this is a heptahydrate

  15. Safety • Lab goggles and coat must be on • Use care with the bunsen burnertie back long hair secure loose clothing keep your notebook away from the work area • Do not handle hot objects by hand • TURN OFF THE GAS USING THE MAIN VALVE ON THE BENCH WHEN EXTINGUISHING BUNSEN BURNERS

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