1. Counting by Weighing

2. Who wants to count 100 M&M’s?
Suppose you work at a candy store…would you want to count out M&M’s one by one?
If you think about it, it makes way more sense to use a scale and count the M&M’s by weighing them
What would you need to know?

3. Average Mass Obtain the mass of 5 different M&M’s.
How do we determine the average mass?
For counting purposes we assume all the items behave as though they were identical.

4. Now what if we have 2 kinds of candy?
I want one bag of gum drops and one bag of M&M’s but they both need to have EXACTLY the same number of items.
How would I figure this out?

5. MAIN IDEA!!!
Items can have different masses yet represent the same number of items.

6. Atomic Masses
Counting by Weighing

7. Keep in mind the gum drop/M&M example
Atoms are extremely TINY so normal units like grams and kilograms are way to LARGE.
For example, the mass of a single carbon atom is 1.66 x grams

8. Atomic Mass Units (amu)
To avoid using exponents like , scientists defined a much smaller unit of mass called atomic mass units (amu)
1 amu = 1.66 x grams

9. Using Atomic Mass Units
Let’s think about
Average Atomic Mass = amu
What mass of carbon atoms must we have to have 1000 carbon atoms present?

10. Using Atomic Mass Units Continued…
We weigh a pile of carbon atoms and the result is
3.00 x 1020 amu. How many carbon atoms are present?
Recall 1 carbon atom = amu

11. Using Atomic Mass Units Continued…
These principles and calculations apply to all of the other atoms
Atomic mass on the PTE refers to amu
Do the following:
1. What is the mass in amu of a sample containing 75 aluminum atoms?
2. Calculate the number of sodium atoms present in amu.

12. THE MOLE!!!!! AMU’s are extremely small units
In lab, we commonly use grams. How do we count atoms in samples with masses given in grams?

13. Visual representations
If we weigh out samples of all the elements such that each sample has a mass equal to that element’s average atomic mass in grams, these samples all contain the same number of atoms

14. The Mole
This number (the number of atoms presents in all the samples) is called the mole.
Mole = the number equal to the number of carbon atoms in grams of carbon
This number has been determined to be x 1023 (Avogadro’s number)

15. The Mole
One mole of something always contains x 1023 units of that substance.
Think about the concept of 1 dozen
A sample of an element with a mass equal to that element’s average atomic mass expressed in grams represents 1 mol of atoms
grams of Carbon= 1 mole of carbon

16. Using the mole in calculations
A sample of hydrogen weighs grams. How many moles of hydrogen are present?
What is the mass of 1 mole of hydrogen?
1 mole of hydrogen = g

17. Calculations Continued…
We know the mass of 1 mol of H atoms so we can determine the number of moles of H atoms in any other sample by comparing its mass with the with the mass of 1 mole of H atoms.
We can follow this process for any element

18. Calculations Continued…
Once we know how many moles of something we have, we can figure out how many individual units are present
1 mole = ? Units
1 mole = x 1023 units
Recall our example…0.496 moles of H. How many atoms of H are present?
DIMENSIONAL ANALYSIS

19. Now you give it a try
Compute the number of moles and the number of atoms present in 10.0g of aluminum.

20. A more complicated example…
How many silicon atoms are present in a 5.68 mg sample of silicon.