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```Chemical Composition
Learning Targets
 I can…
 calculate formula weight, molecular weight, and molar
mass.
 calculate percent composition given a chemical formula.
 convert between the mass, moles, and number of particles
of a substance.
Atomic Masses: Counting Atoms
by Weighing
 Atoms are too tiny to be weighed in grams and
kilograms.
 Instead, we use the atomic mass unit, or amu
 The atomic mass unit =1.66x10-24g
 For example, the average atomic mass for carbon
atoms is 12.01 amu
Calculating Mass Using Atomic
Mass Units (amu)
 Calculate the mass, in amu, of a sample of aluminum
that contains 75 atoms.
 1 Al atom= 26.98 amu
 75 Al atoms x 26.98amu/Al atom=
 2024 amu
Calculating Mass Using Atomic
Mass Units (amu)
 Calculate the mass, in amu, of a sample of nitrogen that
contains 23 atoms.
 1 N atom = 14.01 amu
 23 N atoms x 14.01 amu/N atom=
 322.2 amu
Calculating the Number of Atoms
from the Mass
 Calculate the number of sodium atoms present in a
sample that has a mass of 1172.49 amu
 1 Na atom = 22.99 amu
 1172.49 amu x 1 Na atom/ 22.99 amu=
 51.00 Na atoms
Calculating the Number of Atoms
from the Mass
 Calculate the number of oxygen atoms in a sample that
has a mass of 288 amu.
 1 O atom = 16.00 amu
 288 amu x 1 O atom/ 16.00 amu=
 18.0 O atoms
The Mole
 We have used atomic mass units for mass, but these are
extremely small units.
 In the laboratory a much larger unit, the gram, is the
convenient unit for mass.
 We will learn to count atoms in samples with masses
given in atoms.
The Mole
 We have a sample of aluminum that has a mass of
26.98g.
 What mass of copper contains exactly the same number
of atoms as this sample of aluminum?
 26.98g sample Al = ? Grams of Cu
 We need to know the average atomic masses for
aluminum and copper:
 Al= 26.98 amu
 Cu= 63.55 amu
 Which atom has the greater atomic mass, aluminum or
copper?
 Copper
 If we have 26.98 g of aluminum, do we need more or
less than 26.98 g of copper to have the same number of
copper atoms as aluminum atoms?
 More
 We need more than 26.98 g of copper because each copper
atom has a greater mass than each aluminum atom.
 What does this mean?
 A given number of copper atoms will weigh more than an
equal number of aluminum atoms.
 How much copper do we need?
 aluminum atoms are 26.98 amu
 Copper atoms are 63.55 amu
 26.98g of aluminum and 63.55g of copper contain
exactly the same number of atoms
 Therefore:
 26.98g of aluminum contains the same number of
aluminum atoms as 63.55g of copper contains copper
atoms.
Another example:
 Now compare carbon and helium
 Does a 12.01g sample of carbon contain more or less or the
same number of atoms as 4.003g of helium?
 Carbon = 12.01 amu
 Helium = 4.003 amu
 Both samples contain the exact same number of atoms
The Mole
 The number of atoms present in all of these samples
assumes special importance in chemistry.
 The mole (abbreviated mol) can be defined as:
 The number equal to the number of carbon atoms in 12.01
grams of carbon.
 That is to say, one atomic mass unit is defined to be 1/12
of the mass of a carbon-12 atom.
The Mole
 Techniques for counting atoms very precisely have been
used to determine this number to be 6.022x1023
 One mole of something consists of 6.022x1023 units of
that substance is called:
 For example,
 Just as a dozen is 12 eggs…
 One mole of eggs is 6.022x1023 eggs
How do We Use the Mole in
Chemical Equations?
 12.01g sample of Carbon contains 6.022x1023 atoms
 The atomic mass of hydrogen is 1.008 amu
 A 1.008 g sample of hydrogen contains 6.022x1023 atoms
 The atomic mass of aluminum is 26.98 amu
 A 26.98 g sample of aluminum contains 6.022x1023 atoms
The point is:
 A sample of any element that weighs a number of grams
equal to the average atomic mass of that element
contains 6.022x1023 atoms
 Or one mol of that element
Comparison of 1-Mol Samples of
Various Elements
Element
# of Atoms Present
Mass of Sample (g)
Aluminum
6.022x1023
26.98
Gold
6.022x1023
196.97
Iron
6.022x1023
55.85
Sulfur
6.022x1023
32.07
Boron
6.022x1023
10.81
Xenon
6.022x1023
131.3
 Aluminum (Al), a metal with a high strength-to-weight
ratio and a high resistance to corrosion, is often used
for structures such as high-quality bicycle frames.
Compute both the number of moles of atoms and the
number of atoms in a 10.0g sample of aluminum.
 In this case we want to change from mass to moles of
atoms.
 10.0g of Al= ? Moles of Al atoms
 The mass of one mol (6.022x1023 atoms) of aluminum is:
 26.98g
 The sample we are considering has a mass of:
 10.0g
 Its mass is less than 26.98g, so this sample contains:
 Less than 1 mol of aluminum atoms.
 1 mol Al= 26.98 g Al
 10.0 g Al x 1 mol Al/ 26.98 g Al=
 0.371 mol Al
 Next we convert from moles of atoms to the number of
atoms.




6.022x1023 Al atoms=
1 mol Al atoms
0.371 mol Al x 6.022x1023 Al atoms/ 1 mol Al=
2.23x1023 Al atoms
Molar Mass
 A chemical compound is:
 A collection of atoms
 For example,
 Methane consists of molecules each containing one carbon
atom and four hydrogen atoms (CH4)
 How can we calculate the mass of 1 mol of methane?
 What is the mass of 6.022x1023 CH4 molecules?
Molar Mass
 Each CH4 molecule contains:
 1 carbon atom
 4 hydrogen atoms
 1 mol of CH4 molecules consists of:
 1 mol of carbon atoms
 4 mol of hydrogen atoms
How can we find the mass of 1
mol CH4?
 Mass of 1 mol of C= 1 x 12.01g = 12.01g
 Mass of 4 mol of H= 4 x 1.008g = 4.032g
 Mass of 1 mol of CH4 =
 16.04 g
Calculating Molar Mass
 Calculate the molar mass of sulfur dioxide, a gas
produced when sulfur containing fuels are burned.
 The formula for Sulfur dioxide is:
 SO2
 1 mol SO2 molecules =
 1 mol S atoms
 2 mol O atoms
Calculating Molar Mass
 Mass of 1 mol of S=
 1 x 32.07 =
 32.07 g
 Mass of 2 mol of O=
 2 x 16.00 =
 32.00 g
 Mass of 1 mol of SO2 =
 64.07 g = molar mass
```
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