1. How many cans were required to make a mole of aluminum?
In this lesson students will review all quantitative measurement units used in chemistry they have learned thus far in order to compare and contrast microscopic units and macroscopic units. They should come to the conclusion that both moles and grams are macroscopic units while atoms and molecules are microscopic units. They will end by converting between macroscopic units of grams and moles.
Unit designed by Carole Namowicz a teacher at Lindblom.
When are grams an appropriate measurement, when are moles necessary, and how can we convert between the two units?
You will need the following resources to complete this assignment.
In the following activity, you will measure out a mole, in grams, of several elements in order to see how much a mole really is. Make sure your group has a periodic table, you're going to need it.
Create a 1 mole sample of aluminum using aluminum pop cans. It should be close to a mole, but it does not need to be exact.
1. How many cans were required to make a mole of aluminum?
2. Now weigh a single pop can and record its mass.
3. Calculate how many moles are in a single pop can.
Create a 1 mole sample of iron using paperclips. It should be close to a mole, but it need not be exact.
1. Estimate the total number of paperclips used (don't waste time counting them out one by one).
2. Now weigh 5 paperclips and record their mass.
3. Calculate how many moles are in 5 paperclips.
Create a 1 mole sample of water. It should be close to a mole, but it not be exact. Measure out the amount of water needed in the graduated cylinder. Place the small 50 mL beaker on the scale and zero it. Then pour the water in the small beaker and record the mass of the water. (Hint: the density of water is 1 g/mL. Density is mass/volume. Get the molar mass and solve for the proper volume).
1. Record the mass of one mole of water.
Create a 1 mole sample of salt (salt is sodium chloride = NaCl). Do not place salt directly on the scale, measure it inside of a weighing boat provided. It should be as close to a mole as possible, but it need not be exact.
1. Approximately how many scoops of salt were needed to make a mole of salt?
2. Now weigh one teaspoon of salt and record its mass.
3. Calculate how many moles of salt are in one teaspoon of salt.
In the conservation of mass labs we performed in class, all measurements were taken in grams.
Brainstorm other units chemists use to measure quantities in. Write as many as you can think of below.
Enter all of your units in the table below and categorize each of the units as macroscopic or microscopic.
Which of those units do you think chemists use in the lab (you may include more than one unit in your answer)? Explain why.
Chemical equations can be used to represent what happens on either the microscopic or macroscopic scale.
C (s) + O2 (g) CO2 (g)
This equation can be read in either of the following ways.
1. Which reading of the chemical equation C (s) + O2 (g) → CO2 (g) represents the macroscopic representation? Cite evidence from the sentence reading that led you to that conclusion.
2. Which reading of the chemical equation C (s) + O2 (g) → CO2 (g) represents the microscopic representation? Cite evidence from the sentence reading that led you to that conclusion.
3. When chemists perform chemical equations, are they making products in microscopic or macroscopic quantities? Why?
Your instructor will provide you with a handout so you can practice converting between grams and moles.