Who wants to count 100 M&M’s?
0 Suppose you work at a candy store…would you want
to count out M&M’s one by one?
0 If you think about it, it makes way more sense to use a
scale and count the M&M’s by weighing them
0 What would you need to know?
Average Mass
0 Obtain the mass of 5 different M&M’s.
0 How do we determine the average mass?
0 For counting purposes we assume all the items
behave as though they were identical.
Now what if we have 2 kinds of
candy?
0 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.
0 How would I figure this out?
MAIN IDEA!!!
0 Items can have different masses yet represent the
same number of items.
Keep in mind the gum
drop/M&M example
0 Atoms are extremely TINY so normal units like grams
and kilograms are way to LARGE.
0 For example, the mass of a single carbon atom is 1.66
x 10-24 grams
Atomic Mass Units (amu)
0 To avoid using exponents like 10-24 , scientists defined
a much smaller unit of mass called atomic mass units
(amu)
0 1 amu = 1.66 x 10-24 grams
Using Atomic Mass Units
0 Let’s think about
0 Average Atomic Mass = 12.01 amu
0 What mass of carbon atoms must we have to have
1000 carbon atoms present?
Using Atomic Mass Units
Continued…
0 We weigh a pile of carbon atoms and the result is
3.00 x 1020 amu. How many carbon atoms are present?
0 Recall 1 carbon atom = 12.01 amu
Using Atomic Mass Units
Continued…
0 These principles and calculations apply to all of the other
atoms
0 Atomic mass on the PTE refers to amu
0 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 1172.49
amu.
THE MOLE!!!!!
0 AMU’s are extremely small units
0 In lab, we commonly use grams. How do we count
atoms in samples with masses given in grams?
Visual representations
0 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
The Mole
0 This number (the number of atoms presents in all the
samples) is called the mole.
0 Mole = the number equal to the number of carbon
atoms in 12.01 grams of carbon
0 This number has been determined to be 6.022 x 1023
(Avogadro’s number)
The Mole
0 One mole of something always contains 6.022 x 1023
units of that substance.
0 Think about the concept of 1 dozen
0 A sample of an element with a mass equal to that
element’s average atomic mass expressed in grams
represents 1 mol of atoms
12.011 grams of Carbon= 1 mole of carbon
Using the mole in calculations
0 A sample of hydrogen weighs 0.500 grams. How
many moles of hydrogen are present?
0 What is the mass of 1 mole of hydrogen?
0 1 mole of hydrogen = 1.008 g
Calculations Continued…
0 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.
0 We can follow this process for any element
Calculations Continued…
0 Once we know how many moles of something we
have, we can figure out how many individual units are
present
0 1 mole = ? Units
0 1 mole = 6.022 x 1023 units
0 Recall our example…0.496 moles of H. How many
atoms of H are present?
0 DIMENSIONAL ANALYSIS
Now you give it a try
0 Compute the number of moles and the number of
atoms present in 10.0g of aluminum.
A more complicated example…
0 How many silicon atoms are present in a 5.68 mg
sample of silicon.