Molar Relationships Lab Part 3

Moles Lab Activity 3: Compounds—Chalk Materials Sidewalk chalk Balance Procedure 1. Find the mass of a piece of sidewalk chalk. 2. Write your name on the sidewalk or parking lot with the piece of chalk, but do not use it up completely. 3. Find the mass of your piece of chalk following your writing. 4. Answer the following questions, taking the necessary measurements and recording them with units. Show all your calculations, rounding your answers to the teacher-specified number of significant digits and labeling units clearly. What is the molar mass of chalk (calcium carbonate [CaCO3])? What is the mass of your piece of chalk before writing? What is the mass of your remaining chalk after writing? How many grams of chalk did you leave on the ground? Is the amount left on the ground more than, equal to, or less than one mole of chalk? How many moles of chalk did you leave on the ground? How many molecules of chalk did you leave on the ground? How many individual atoms of oxygen did you leave on the ground? 5. Check your answers with the student calcium carbonate experts, and ask them to initial your original data to certify that they are correct. Extension 1. Calculate the percent oxygen by mass of chalk (calcium carbonate [CaCO3]). 2. Answer the following questions: How many calcium (Ca) atoms did you leave on the ground? How many atoms (total of all types) did you leave on the ground? Is chalk an ionic or covalent compound? Why? Moles Lab Activity 3: Compounds—The Fictitious Compound “Candium” Materials Sample of “candium” (mixture of 12 M&M’s and 8 Skittles) 2 small paper cups Balance Procedure Take the necessary measurements, and record them with units. Show all your calculations, rounding your answers to the teacher-specified number of significant digits and labeling units clearly. 1. Assume that each kind of candy (“Mm” and “Sk”) in the sample of “candium” represents a different type of atom and that the sample is the compound. Write the formula in the format MmxSky for the compound “candium,” using the smallest whole-number ratio of Mm to Sk. 2. Find and record the molar mass of candium. (Assume your sample is one mole.) 3. Find and record the mass of the Mm in your sample. 4. Find and record the mass of the Sk in your sample. 5. Answer the following questions: What is the average mass of an Mm (the molar mass of Mm)? What is the average mass of an Sk (the molar mass of Sk)? What is the percent by mass of Mm in candium? Use this equation: % by mass Mm = molar mass Mm × number of Mm atoms in candium formula × 100 molar mass of candium What is the percent by mass of Sk in candium? Use a similar equation. 6. Check your answers with your teacher, and ask him/her to initial your original data to certify that they are correct. Extension “Candium Isotope” Activity 1. Assume your candium sample is a sample of an element, not a compound. In this element, Mm and Sk are the two naturally occurring isotopes of the element. What is the percent abundance (by number) of each isotope in your sample? 2. What is the average atomic mass of the element candium? Use the average mass of each isotope as its mass number and its percent abundance to calculate, using the following formula: [Mass #(Mm) (%Mm) + Mass #(Sk) (%Sk)] ÷ 100 = average atomic mass 3. Compare the average atomic mass to the molar mass (total mass of all candies in the sample). What is your percent error? Moles Lab Activity 4: Solutions—Aqueous Copper (II) Sulfate Pentahydrate Materials Beaker or plastic cup Copper(II) sulfate pentahydrate (CuSO4 • 5 H2O) crystals Graduated cylinder Water Filter paper Funnel Ring stand and ring Watch glass or Petri dish Procedure Take the necessary measurements, and record them with units. Show all your calculations, rounding your answers to the teacher-specified number of significant digits and labeling units clearly. 1. Find and record the molar mass of hydrated copper(II) sulfate (copper(II) sulfate pentahydrate [CuSO4 • 5 H2O]). 2. Find and record the mass of 0.0125 moles of copper(II) sulfate pentahydrate. 3. Weigh out 0.0125 moles of copper(II) sulfate pentahydrate. Using a graduated cylinder to measure, place 25.0 mL of water into a plastic cup or beaker. 4. Record the color changes you observe as you add the 0.0125 moles of copper(II) sulfate pentahydrate to the water and stir to dissolve completely. 5. Answer the following questions: Is the mixture heterogeneous or homogeneous? What is the solute in this solution? What is the solvent? How many copper atoms are in one mole of copper(II) sulfate? How many copper atoms are in your 25.0 mL of copper(II) sulfate solution? What is the percent water by mass in copper(II) sulfate? What is the percent copper in the 0.0125 moles of copper(II) sulfate? Extension 1. Predict the appearance of the filtrate (the substances that pass through the filter) if you were to filter the copper(II) sulfate solution. 2. Assemble a filtration apparatus, using a ring stand and ring, a plastic funnel, a beaker, and a piece of filter paper. Sketch and label the apparatus in your data book. 3. Prime the filter with tap water, and filter the copper(II) sulfate solution. 4. Describe the filtrate, and compare it to your prediction. Account for similarities and/or differences. What substances are found in the filtrate? 5. How many oxygen atoms are in one mole of copper(II) sulfate (CuSO4)? Calculate the number of oxygen atoms in your 25.0 mL of saturated copper(II) sulfate solution. 6. Draw a model of an aqueous copper(II) ion by drawing Cu2+ surrounded by water molecules. Which end of the water molecules is attracted to the Cu2+ ion? 7. Based on the copper(II) ion (Cu2+) and the formula for copper(II) sulfate, what is the charge of the sulfate ion (SO4)? 8. Place 1 mL of your filtrate on a watch glass, and allow the water to evaporate (this may take a day or two). Draw the shape of the crystals that form. Moles Lab Activity 4: Solutions—Alum Materials Beaker or plastic cup Alum (hydrated aluminum potassium sulfate [AlK(SO4)2 • 12 H2O]) crystals Water Filter paper Funnel Ring stand and ring Procedure Take the necessary measurements, and record them with units. Show all your calculations, rounding your answers to the teacher-specified number of significant digits and labeling units clearly. 1. Find and record the molar mass of alum (hydrated aluminum potassium sulfate [AlK(SO4)2 • 12 H2O]). 2. Find and record the mass of 0.050 moles of alum. 3. Weigh out 0.050 moles of alum. Using a graduated cylinder to measure, place 25.0 mL of water into a plastic cup or beaker. Add the 0.050 moles of alum to the 25.0 mL of water, and mix thoroughly. Record your observations. 4. Answer the following questions: Is the mixture saturated or unsaturated? What is the solute in this solution? What is the solvent? How many aluminum atoms are in one mole of aluminum potassium sulfate? How many aluminum atoms are in your 25.0 mL of aluminum potassium sulfate solution? What is the percent water by mass in aluminum potassium sulfate? What is the mass of sulfur in the 0.050 moles of aluminum potassium sulfate? 5. Check your answers with your teacher, and ask her/him to initial your original data to certify that they are correct. Extension 1. Predict the appearance of the filtrate (the substances that pass through the filter) if you were to filter the aluminum potassium sulfate solution. 2. Assemble a filtration apparatus, using a ring stand and ring, a plastic funnel, a beaker, and a piece of filter paper. Sketch and label the apparatus in your data book. 3. Prime the filter with tap water, and filter the AlK(SO4)2 • 12 H2O solution. 4. Describe the filtrate, and compare it to your prediction. Account for similarities and/or differences. What substances are found in the filtrate? 5. How many sulfur atoms are in one mole of alum? Calculate the number of sulfur atoms in your 25.0 mL of saturated alum solution. 6. How many oxygen atoms are in one mole of alum? Calculate the number of oxygen atoms in your 25.0 mL of saturated alum solution. 7. Place 1 mL of this solution on a watch glass, and allow the water to evaporate (this may take a day or two). Draw the alum crystals remaining after the water evaporates. Show clearly the shape of one crystal. Moles Lab Activity 5: Synthesis of an Oxide of Copper Materials Hot plate Evaporating dish Balance Granular or powdered copper Pre-Lab Copper has two common ions: copper(I) (Cu+) and copper(II) (Cu 2+). When heated in air, the copper will react with oxygen gas (O2) to form an oxide. The Roman numeral is used in the name to differentiate between the two possible oxides of copper. Remember that the Roman numeral indicates the charge on the metal, not the quantity. 1. Determine the formula for each oxide of copper, and indicate its name. Check with the teacher or the student copper experts for accuracy of these answers before proceeding. 2. Determine the percent oxygen in each oxide. 3. Based on your answers from the previous step, which oxide should form first in this experiment? Why? Write the balanced equation for its formation. 4. The second oxide cannot be formed directly, but is formed from the further oxidation of the first oxide. Write a balanced equation for the formation (synthesis) of the second oxide, using the first oxide and oxygen gas as the reactants. 5. Check that the answers to the pre-lab are correct before performing the experiment below. Procedure Take the necessary measurements, and record them with units. Show all your calculations, rounding your answers to the teacher-specified number of significant digits and labeling units clearly. You must wear goggles when heating. 1. Find and record the molar mass of copper. 2. Find and record the mass of a clean, dry evaporating dish. 3. Determine the mass of 0.30 moles of copper. 4. Place 0.30 moles of copper into the evaporating dish, and record the total mass before heating. 5. Place the evaporating dish on the hot plate, and turn it on to the highest setting. Heat strongly for at least 15 minutes, and record at least four qualitative observations while heating. 6. Use crucible tongs to carefully remove the evaporating dish from the hot plate, and allow the dish to cool for 5 to 10 minutes. 7. Find the mass of the dish and oxide of copper after heating, and record. 8. Determine the mass of the copper(x) oxide formed. 9. Determine the mass of oxygen reacted (the difference in mass before and after heating). 10. After the product has cooled, observe it carefully for any signs of an incomplete reaction. Record observations. Analysis 1. Answer the following questions: What is the experimental percent oxygen in your product? Report this in the class data table. Based on your percent oxygen, which oxide most likely was formed? Based on this result, what is your percent error? Look up the color of copper(I) oxide and copper(II) oxide. Does this information support your data? Why, or why not? Is the class data accurate and/or precise? Why, or why not? What experimental errors would account for any differences you observed? Extension 1. Iron can also form two oxides. Look up the possible oxidation states (charges) for iron, write the formulas for both oxides, and indicate their names. 2. Determine the theoretical percent oxygen of each oxide. 3. In an experiment similar to the one described above, a student reacted 1.20 g of iron and determined that it gained 0.5 grams upon heating. Determine the percent oxygen in the product. 4. Determine the empirical formula for the oxide formed. 5. Identify which oxide was formed in the experiment. What evidence do you have to support this conclusion? 6. Look up the color of each possible oxide of iron. What would be the color of the product in the experiment? Why? 7. In forming the oxides of transition metals with multiple oxidation states by heating them in air, identify the two things that most likely determine which oxide is formed.