Transcript The Mole-Ch 11-Sections 11.1

The Mole
Chapter 11 – Honors Chemistry
LSM High School
Section 11.1:
Measuring Matter

Objectives:
Describe how a mole is used in chemistry
 Relate a mole to common counting units
 Covert moles to number of representative
particles and number of representative particles
to moles.

How do Chemists measure
how much of a substance?
Chemists can measure mass or volume or
they can count pieces.
 Chemists can measure mass in grams.
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Chemists can measure volume in liters.
No, not that kind of mole!!!
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Chemists can count pieces in MOLES.
What are MOLES?

Moles are defined as the number of carbon
atoms in exactly 12 grams of the carbon-12
isotope.

1 mole is 6.02 x 1023 particles.
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Treat it like a very large dozen
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6.02 x 1023 is called Avogadro's number.
A Little History

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Amedeo Avogadro was born in
1776 in Turin, Italy.
He went on to study molecular
theory and helped other
scientists distinguish between
atoms and molecules.
Because of his
accomplishments in this field,
the variable that tells the
number of molecules in one
mole was named after him
What about units on
Avogadro’s number?

The units of Avogadro’s number can be
whatever particle you are counting.

Examples: atoms, molecules, ions, etc…

In chemistry these are called Representative
Particles
What are Representative
Particles?
These particles are the smallest pieces of a
substance.
• The types of representative particles that chemists
generally work with are:

• atoms – the smallest particle of an element
• ions – atoms with positive or negative charges
• molecules – two or more covalently bonded atoms
• formula units – the simplest ratio of ions that make up
an ionic compound
How Do We Use Moles?
Moles are used as conversion factors.
 This means they are used to change units.
 Remember, when solving using conversion
factors there are 3 questions you want to ask
yourself:

What unit do you want to get rid of?
 Where does it go to cancel out?
 What can you change it into?

Converting Moles to Particles
and Particles to Moles
Using Avogadro’s Number as a Conversion Factor
How do we write Avogadro’s
number as a conversion factor?
6.02 x 1023 particles or
1 mole
1 mole
6.02 x 1023 particles
Practice Problem 1
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How many atoms are in 2.50 mol of zinc?
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K:
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Answer: 1.51 x 1024 atoms Zn
UK:
Practice Problem 2
How many molecules of CO2 are the in 4.56
moles of CO2 ?
 K:
UK:
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Answer: 2.75 x 1024 molecules of CO2
Practice Problem 3
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How many moles of water is 5.87 x 1022
molecules of water?
K:
UK:
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ANSWER: 0.0975 moles of water
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Practice Problem 4
Given 3.25 mol AgNO3, determine the
number of formula units.
 K:
UK:
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ANSWER: 1.96 x 1024 formual units AgNO3
Section 11.2:
Mass and the Mole
• Relate the mass of an atom to the mass of a
mole of atoms.
• Calculate the number of moles in a given mass
of an element and the mass of a given number
of moles of an element.
• Calculate the number of moles of an element
when given the number of atoms of an element.
• Calculate the number of atoms of an element
when given the number of moles of the
element.
Let’s Look at the Periodic
Table!

What are some patterns that you see on the
chart?
Atomic Numbers
Always increase across a row.
 The atomic number is the number of protons
in an atom of that element.
 This number identifies it as an atom of a
particular element.

Atomic mass
Usually increase across a row – but not
always.
 Ex:
 Why do they have decimal values?
 The atomic mass (sometimes called average
atomic mass) is the weighted average mass
of the isotopes of that element.
 You will be doing a series of activities to
better understand how atomic masses and
Avogadro’s number was determined.
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How Atomic Masses and the
Mass of a Mole are Related:
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Atomic masses are a relative scale.

Use isotope carbon-12 as the standard
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Each atom of carbon-12 has a mass of exactly 12 amu (atomic
mass units)
Ex: One atom of hydrogen-1 has a mass of 1 amu, that
means that 1 atom of hydrogen-1 is one-twelfth the
mass of one atom of carbon-12
Atomic masses are on the periodic table but they
are not whole numbers b/c the values are weighted
averages of the masses of all the naturally
occurring isotopes of each element
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Since mole is the number of representative particles,
or atoms, in exactly 12 g of pure carbon-12, then the
mass of one mole of carbon-12 atoms is 12 g.
The mass in grams of one mole of ANY pure
substance is its molar mass.
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Same value of atomic mass but has units of g/mol
12.01 grams of carbon has the same number of
particles as 1.01 grams of hydrogen and 55.85 grams
of iron.
The number of particles is 6.02 x 1023 atoms.
Therefore, we can count things by weighing them.
Using Molar Mass:
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The molar mass can be used as a conversion factor.
It relates the mass of a substance to the number of
moles of that substance.
To calculate the mass from the number of moles, you
would use the molar mass as:
# of grams
1 mole
To calculate the moles from the mass, you would use
the molar mass as:
1mol
# of grams
Example Problems:
 
2.34molesC  
 
What is the mass, in grams, of 2.34 moles
of carbon?
grams C
K: 2.34 moles C
UK:
1.
12.01 g/moles C
28.1 g carbon
2. How many moles of magnesium are in
4.61g of Mg?
K:
UK:
0.190 mol Mg
Conversions from mass to
atoms and atoms to mass
There is no direct conversion between the
mass of a substance to the number of
representative particles of that substance.
 You must first convert to moles and then
convert to the desired unit either using molar
mass or Avogadro’s number.

Example Problems:
1. How many atoms are in 45.6 g Si?
K:
UK:
9.77 x 1023 atoms Si
2. How many atoms are in 0.120 kg Ti?
K:
UK:
1.51 x 1024 atoms Ti
3. What is the mass, in grams, of 1.50 x 1015
atoms N?
K:
UK:
3.49 x 10-8 g N
4. What is the mass, in grams, of 1.50 x 1015
atoms uranium?
K:
UK:
5.93 x 10-7 g U