Lesson 4. Lesson 4: Grams vs. Moles

Carole Namowicz
Chemistry
50 min
High School Honors Chemistry
v2

Overview

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.

Standards

Next Generation Science Standards
  • Physical Science
  • NGSS Practice
    • Using Mathematics
  • NGSS Crosscutting Concept
    • Energy
Computational Thinking in STEM
  • Data Practices
    • Analyzing Data
  • Modeling and Simulation Practices
    • Assessing Computational Models
    • Using Computational Models to Understand a Concept
  • Systems Thinking Practices
    • Defining Systems and Managing Complexity
    • Thinking in Levels
    • Understanding the Relationships within a System

Credits

Unit designed by Carole Namowicz a teacher at Lindblom.

Activities

  • 1. Measuring a Mole
  • 2. So How Much is a Mole?
  • 3. So How Much is a Mole?
  • 4. So How Much is a Mole?
  • 5. So How Much is a Mole?
  • 6. Chemistry Units of Measure
  • 7. Chemical Equations at the Microscopic and Macroscopic Scale
  • 8. Converting Between Grams and Moles

Student Directions and Resources


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.

1. Measuring a Mole


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.

 

 

 

 

 

2. So How Much is a Mole?


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.


Question 2.1

1. How many cans were required to make a mole of aluminum?



Question 2.2

2. Now weigh a single pop can and record its mass.



Question 2.3

3. Calculate how many moles are in a single pop can.

3. So How Much is a Mole?


Create a 1 mole sample of iron using paperclips. It should be close to a mole, but it need not be exact.


Question 3.1

1. Estimate the total number of paperclips used (don't waste time counting them out one by one).



Question 3.2

2. Now weigh 5 paperclips and record their mass.



Question 3.3

3. Calculate how many moles are in 5 paperclips.

4. So How Much is a Mole?


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).

 

 

 

 


Question 4.1

1. Record the mass of one mole of water.

5. So How Much is a Mole?


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.

 

 

 

 


Question 5.1

1. Approximately how many scoops of salt were needed to make a mole of salt?



Question 5.2

2. Now weigh one teaspoon of salt and record its mass.



Question 5.3

3. Calculate how many moles of salt are in one teaspoon of salt.

6. Chemistry Units of Measure


In the conservation of mass labs we performed in class, all measurements were taken in grams.


Question 6.1

Brainstorm other units chemists use to measure quantities in. Write as many as you can think of below.



Question 6.2

Enter all of your units in the table below and categorize each of the units as macroscopic or microscopic.



Question 6.3

Which of those units do you think chemists use in the lab (you may include more than one unit in your answer)? Explain why.

7. Chemical Equations at the Microscopic and Macroscopic Scale


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. If, or when, solid carbon reacts with gaseous oxygen, one atom of carbon and one molecule of oxygen are consumed for every one molecule of gaseous carbon dioxide produced.
  2. If, or when, solid carbon reacts with gaseous oxygen, one mole of carbon and one mole of oxygen are consumed for every one mole of gaseous carbon dioxide produced.

 

 

Question 7.1

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.



Question 7.2

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.



Question 7.3

3. When chemists perform chemical equations, are they making products in microscopic or macroscopic quantities? Why?

8. Converting Between Grams and Moles


Your instructor will provide you with a handout so you can practice converting between grams and moles.