Notes: Section 5.3
Atomic Number:
The number of protons in the nucleus.
Note: Henry Mosley
Developed a technique using X-Rays for determining the number of protons an element contains.
Mass Number:
(Protons + Neutrons).
Atomic Mass:
The atomic mass shown in the periodic table is an average of all the naturally occurring isotopes of
the element.
Relative Mass:
Comparing the mass to a standard.
Atomic Mass Unit “amu” or “u”
1/12 the Mass of Carbon-12.
Isotope:
An element that has the same number of protons but a different number of
neutrons.
(Uranium-235 and Uranium-238)
(Carbon-12 and Carbon-14)
Three Isotopes of Hydrogen:
Protium
Deuterium
1
1H
2
1H
0 neutrons
1 proton
Mass = 1
99.985%
Relative Abundance
Tritium
3
1H
1 neutron
1proton
Mass = 2
2 neutrons
1 proton
Mass = 3
0.015%
Relative Abundance
Negligible
Relative Abundance
Relative Abundance:
 Describes the relative amounts of the naturally occurring isotopes and represents them
as a percent.
Calculating Atomic Mass
when given the
Relative Abundance:
Example:
Copper-63 = 69.2% with 62.93 amu
Copper-65 = 30.8% with 64.93 amu
Notes Section 5.3-5.4
Class Activity:
Complete the Table:
Subatomic
Particle
Symbol
Proton
Charge
Location
Relative mass
(amu)
nucleus
1 amu
p+
Electron
1-
Neutron
Complete the table for each isotope.
Atomic #
Isotope
Symbol
Z=(protons)
neutrons
electrons
Mass
Number
(p+ + n0)
Isotope
Mass in
(amu)
Relative
abundance
Carbon-12
12
6
C
6
12 - 6 = 6
6-
6 + 6 = 12
12.000
98.89%
Carbon-13
13
6
C
6
13 – 6 = 7
6-
6 + 7 = 13
13.003
1.11%
Nitrogen-14
N
14.003
99.63%
Nitrogen-15
N
15.000
0.37%
Chlorine-35
Cl
34.969
75.77%
Chlorine-37
Cl
36.966
24.23%
Use the information from the table to calculate Average Atomic Mass for nitrogen and chlorine.
(Use example as a guide)
Isotope
Example: Carbon-12
Carbon-13
Mass (amu)
12.000
13.003
x
x
x
Relative abundance (%/100)
98.89/100
=
1.11/100
=
11.8668
amu
+ 0.1443333 amu
12.0111333 amu
Average atomic mass = 12.011 amu