Modern View of Atomic Structure
Protons
 Positively charged subatomic particles
 Located in the nucleus of an atom
 Make up most of the atom’s mass
 Pos. charge of proton = Neg charge of electron
(e-)
 1 proton has as much mass as around 1840 e-‘s
 Periodic table is arranged in increasing number
of protons
Neutrons
 Neutral subatomic particle
 Located in the nucleus of the atom
 DOES contribute to the mass of the atom
 Most common type of hydrogen (called
protium) contains NO neutrons. However,
deuterium (heavy hydrogen) does contain 1
proton and 1 neutron in the nucleus.
Electrons
 Negatively charged subatomic particle
 Located in the space outside of the nucleus
 Mass of the electron is negligible
Subatomic Particles Summary
Subatomic
Mass in grams
Particle
Electron (e-)
9.11 x 10-28 g
Proton (p, +)
1.67 x 10-24g
Neutron (n,o)
1.67 x 10-24 g
Mass in amu
0.000549 amu
1.0073 amu
1.0087 amu
Distinguishing between Atoms
Atomic Number
 Is the number of protons in the nucleus of an atom
of an element
 Used to identify an element
 ALWAYS a whole number
Mass Number
 Most of the atom’s mass is concentrated in the
nucleus (protons and neutrons)
 Mass Number = # of protons + # of neutrons
 Electron mass is so small that it is NOT included
in the mass #
 Number of neutrons = Mass # - Atomic #
Isotopes
 All atoms of the same element have the same # of
protons but may have DIFFERENT number of
neutrons
 Isotope  atoms of the same element with different
# of neutrons
 Often identify an element by mass number
 Example: carbon-12, carbon-14
Atomic Mass
 Atomic masses shown on the PT (periodic table)
represent a weighted average based on the relative
abundance of each isotope of a particular atom.
 Atomic mass unit (amu) is a unit of mass equal to
1/12th the mass of a carbon-12 atom.
 amu’s are used instead of grams because the
masses of subatomic particles are small
 More useful to compare the relative masses of
atoms using a reference isotope as a standard
 Carbon-12 was chosen as the reference isotope.
Average Atomic Mass
The atomic mass on the periodic table is an average of
all the isotopes of that particular element. Every
element (except the synthetic ones) has one or more
isotopes
Σ (mass of isotope X relative abundance)
Practice:
1. Neon has two isotopes: Ne-20 (having a mass of 20
u) and Ne-22 (having a mass of 22 u). Given the
following abundances of these isotopes in nature,
what is the average atomic mass of neon?
Mass number
Abundance
Ne-20
90%
Ne-22
10%
2. What is the average atomic mass of hafnium given
the following abundance information on its
isotopes?
Mass number
Abundance
Hf-176
5%
Hf-177
19 %
Hf-178
27 %
Hf-179
14%
Hf-180
35%
3. Calculate the atomic mass of potassium if the
abundance atomic masses of the isotopes making up
its naturally occurring samples are as given below.
Isotope
Relative
abundance
Atomic Mass
potassium-39
93.12 %
38.964 amu
potassium-41
6.88 %
40.962 amu