Download The atom

Chemistry Unit 2
Part 1 Atoms:
The Building
Blocks of Matter
(atomic structure,
isotopes, and
average atomic
mass)
The atom
• Idea was first supported as
early as 400 b.c by
Democritus
• Atom is based on the Greek
word for “indivisible”
1700’s:
Law of Conservation of Mass
Mass is neither
created nor destroyed
during chemical or
physical reactions.
Total mass of reactants
=
Total mass of products
Antoine Lavoisier
Atomic Particles
Particle Charge
Mass (kg)
Location
Electron
-1
9.109 x 10-31
Electron
cloud
Proton
+1
1.673 x 10-27
Nucleus
0
1.675 x 10-27
Nucleus
Neutron
The Atomic
Scale
Most of the mass
of the atom is in
the nucleus (protons
and neutrons)
 Electrons are
found outside of
the nucleus (the
electron cloud)
 Most of the
“q”
volume of the atom
is empty space

is a particle called a “quark”
Atomic Number
Atomic number (Z- the whole number on the
PT) of an element is the number of protons in
the nucleus of each atom of that element.
Element
Carbon
# of protons Atomic # (Z)
6
6
Phosphorus
15
15
Gold
79
79
Isotopes
Elements occur in
nature as
mixtures of
isotopes.
Isotopes are
atoms of the
same element
that differ in
the number of
neutrons
2 Ways to Notate an isotope
1.) Hyphen notation
Element name – mass number
Example 1: Nitrogen-15
How many protons?
______
7
7
How many electrons? ______
8
How many neutrons? ______
Mass # = p+ + n0
2 Ways to Notate an isotope
2.) Nuclear Symbol
Example 2:
238
92
U
How many protons?
How many electrons?
How many neutrons?
92
92
______
146
______
______
Mass Number:
the number of protons and neutrons in
the nucleus of an isotope (this is not on
the pT).
Nuclide
p+
n0
e-
Oxygen - 18
8
10
8
18
Arsenic - 75
33
42
33
75
Phosphorus - 31
15
16
15
31
Mass # = p+ + n0
Mass #
Isotopes…Again
(must be on the test)
Isotopes are atoms of the same element having
different masses due to varying numbers of neutrons.
Isotope
Protons
Electrons
Neutrons
Hydrogen–1
(protium)
1
1
0
Hydrogen-2
(deuterium)
1
1
1
Hydrogen-3
(tritium)
1
1
2
Nucleus
 The
average atomic mass is a
WEIGHTED average of ALL of the
isotopes for an element.
 The
average atomic mass of a
sample of an element can be
found on the periodic table
◦Ex.) Zinc = 65.39 amu
Average Atomic Mass Ex:
Carbon = 12.011
Isotope
Symbol
Composition of
the nucleus
% in nature
Carbon-12
12C
6 protons
6 neutrons
98.89%
Carbon-13
13C
6 protons
7 neutrons
1.10%
Carbon-14
14C
6 protons
8 neutrons
<0.01%
+
+
+
=
(% abundance x mass number of isotope 1)
(% abundance x mass number of isotope 2)
(% abundance x mass number of isotope 3)
etc.
average atomic mass for the element
DO NOT DIVIDE BY ANYTHING! This is not
a regular average. Don’t forget to change
the percentage to a decimal first!!! (move it
to the left)
Copper has two naturally occurring
isotopes: copper-63 (69.17%) and
copper-65 (30.83%).
Calculate the average atomic mass of
copper if the relative masses of the
isotope are copper 63 (62.93 amu)
and copper-65 (64.93 amu).
Cu-63 (0.6917 x 62.93 amu)
Cu-65 + (0.3083 x 64.93 amu)
1)
2)
3)
Change % to a decimal
Line up your info.
Use parentheses on calculator if you
want to do it all at once without having
to write it down for each isotope!
Cu-63 (0.6917 x 62.93 amu)
Cu-65 + (0.3083 x 64.93 amu)
Avg. mass = 63.55 amu