Transcript Section 6.1 Atoms and Moles C. The Mole

Section 6.1
Atoms and Moles
Objectives
1. Students will be able to describe the concept of
average mass.
2. Students will be able to demonstrate how counting
can be done by weighing.
3. Students will be able to define atomic mass and use
it to convert between mass and number of atoms.
4. Students will be able to define a mole and
Avogadro’s number.
5. Students will be able to complete calculations
converting between mass, moles, and number of
atoms.
Section 6.1
Atoms and Moles
Class Warmup
Section 6.1
Atoms and Moles
Counting by Weighing
How many pennies are in
the jar?
Count or Weigh?
To determine by weighing what
information is necessary?
1. Mass of one penny
2. Total mass of the jar of pennies
3. Mass of the jar
Section 6.1
Atoms and Moles
Counting by Weighing
Is there a flaw in this method?
• What if all of the pennies do not have the same mass?
• Find the avg. mass.
How is this relevant to counting the number of atoms?
• Atoms of the same element can have different masses
(isotopes)
Section 6.1
Atoms and Moles
Counting by Weighing
Let’s say you wanted to place 1000 pennies into a
container. You can do this without actually counting
them if you have a balance. If you know the avg. mass
of a penny, you can determine the total mass of 1000
pennies.
Section 6.1
Atoms and Moles
B. Atomic Masses: Counting Atoms by Weighing
• Atoms have very tiny masses so scientists made a unit
to avoid using very small numbers.
1 atomic mass unit (amu) = 1.66 10-24 g
• Scientists defined the atomic mass unit:
1 Carbon-12 atom = 12 amu
• Since carbon-12 has a sum of 12 protons and
neutrons and the protons and neutrons have very
similar masses:
1 proton ~ 1 amu
1 neutron ~ 1 amu
Section 6.1
Atoms and Moles
B. Atomic Masses: Counting Atoms by Weighing
The mass of an atom is approximately equal to its mass
number
65Cu ~ 65 amu
35Cl ~ 35 amu
For carbon 12, the atomic mass = 12 amu (exactly)
For all other isotopes, the exact mass is slightly
different from the mass #
35Cl = 34.969 amu
37 Cl = 36.966 amu
Section 6.1
Atoms and Moles
B. Atomic Masses: Counting Atoms by Weighing
• The atomic mass listed on the periodic table is the
weighted average of all naturally occurring isotopes
of that element
Section 6.1
Atoms and Moles
C. The Mole
Objectives:
1. Define a mole
2. Use the mole to calculate atoms, molecules, and
mass.
Section 6.1
Atoms and Moles
C. The Mole
The Mole
• A counting unit used to count atoms, molecules,
formula units, electrons, etc.
• 1 mole = 6.02 x 1023 (Avogadro’s Number)
How was the mole determined?
• 1 mole equals the number of carbon-12 atoms is
exactly 12 grams of carbon 12.
Section 6.1
Atoms and Moles
C. The Mole
• Because of the way atomic mass units and the mole are
defined, there is a relationship between grams and amu
So…
1 atom of sodium (Na) = 22.98977 amu
1 mole of sodium (Na) = 22.98977 grams
Section 6.1
Atoms and Moles
C. The Mole
• One mole of anything contains 6.022 x 1023 units of that
substance.
– Avogadro’s number is 6.022 x 1023.
• A sample of an element with a mass equal to that
element’s average atomic mass (expressed in g)
contains one mole of atoms.
Section 6.1
Atoms and Moles
C. The Mole