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How Atoms Differ
Objectives Covered in this Presentation:
7. Locate the three fundamental particles in the atom; indicate the
relative mass and charge of these particles.
8. Indicate what the atomic number and mass number of an atom
represent.
9. Calculate the mass number when given the number of protons and
neutrons.
10. Indicate which isotope is used as the standard for atomic mass.
11. Write the nuclear symbol for an atom.
12. Determine the number of protons, neutrons, and electrons when
given the atomic number and mass number.
13. Calculate average atomic mass from relative abundance.
14. Define isotope and identify which nuclides are isotopes of the
same element when given the atomic number and mass number of
these nuclides.
Atomic structure
Protons: positive particles in nucleus
Neutrons: particle with no charge in the
nucleus
Electron: particle with negative charge orbiting
the nucleus
Subatomic particles
Quarks: make up protons and neutrons
Six types: up, down, strange, charm, truth
(top) and beauty (bottom)
(I am not making up those names, nuclear
physicists can be weird.)
Subatomic Particles
(Cont)
Other subatomic particles include
leptons, neutrino, tau, muon etc.
Atomic Symbols
SI symbols used internationally.
Based on Greek, Latin, a person’s name, or
location discovered.
Examples: Iron <Latin Ferrum>: Fe
Silver <Latin Argentum>: Ag
Atomic Symbols
1st letter capitalized
2nd and 3rd letter lower case
MUST be printed in traditional letters
NO cursive
Aluminum: Al not Al
Cobalt: Co not CO or co
Atomic Number
Equal to the number of protons
The number of protons for an element
never changes
Atomic Number
In a stable atom, the number of
protons = electrons
The positive and negative charges are
equal and the overall charge is neutral
Nuclear Charge
If electrons are gained, the number of
electrons is greater than the number
of protons and the charge is
__________
Electric Charge
If electrons are lost, the number of
electrons is less than the number of
protons and the charge is
__________
Atomic Mass
May be noted as mass weight, or
atomic weight
Equals: protons + neutrons
Atomic Mass
Each proton and each neutron have
an atomic mass of 1 amu. (Atomic
mass units).
Electrons are so tiny (.00005 amu)
that their mass is negligible.
Atomic Mass is an
Average
Remember: the number of protons for
any element does not change
The number of neutrons does vary.
Example: All carbon atoms have 6
protons. Most carbon atoms has 6
neutrons, but very few have 8.
6P + 6N = 12 amu
6P + 8N = 14 amu
If you have 9 carbon atoms with 12
amu and 1 carbon atom with 14 amu,
what is the average amu?
If you have 9 carbon atoms with 12
amu and 1 with 14 amu, what is the
average amu?
9 atoms X 12 amu = 108
1 atom X 14 amu = 14
108 + 14 = 122
122/10 =12.2 amu
Isotopes: atoms of the same element
with different numbers of neutrons
and slightly different amu.
The greater the number of protons for
an element, the greater the number of
isotopes.
Periodic Table History: 1869
Mendeleev created 1st Periodic Table.
He arranged it by atomic mass and
successfully predicted unknown
elements.
Modern Periodic Table:
Arranged by atomic number
List element name
Element symbol
Atomic number
Atomic mass
Periods: (rows) all elements in the row
have the same number of energy
levels
Families: (column) have the same
number of valence electrons and
oxidation state.
Valence electrons: number of
electrons in the outer most energy
level
Oxidation state: number of electrons
needed to be gained or lost to have 8.
Using the periodic table
Protons = ?
Using the periodic table
Protons = atomic number
Electrons = ?
Electrons = number of protons
Neutron = ?
Neutrons = atomic mass – atomic
number