Atomic Theory
Objectives:
 1. Name and describe the three subatomic particles in
an atom.
 2. Determine the number of protons, neutrons, and
electrons in an atom or ion.
 3. Define isotope and atomic mass.
Key Terms:
 Proton, neutron, atomic mass unit, atomic number,
ion, mass number, isotope, atomic mass
Subatomic Particles
Atoms are made of even smaller subatomic particles
called protons, neutrons, and electrons.
 protons - have a positive charge, found in the nucleus and have
an atomic weight of 1amu
 electrons - have a negative charge, move in the space around the
nucleus, the number is equal to protons in a neutral atom and
they have no appreciable mass
 neutrons - have no net charge, found in the nucleus, and have an
atomic weight of 1amu
Atomic Number
A. Atomic number – number of protons in the nucleus.
 The number of protons determines the positive charge
of the atom
 The number of protons determines the atom’s identity
 The electrons and neutrons of an atom can change but it
cannot change the amount of protons and still be of the
same element
 In neutral atoms the # of protons = # of electrons
B. Mass Number = protons + neutrons
 32amu – 16(+) = 16(+)
C. & D. – Chemical symbol and name
Ions
Whenever an atom gains or loses electrons it becomes
an ion. You can find the net charge of the atom by
subtracting the number of electrons from the atomic
number.
 Loses an e-… the result is positive (atomic # >
electrons) called a cation - overall positive charge
 example: Ca… 20(+) - 18(-) = 2+ or Ca2+
 Gains an e-… the result is negative (atomic # <
electrons) called an anion - overall negative charge
 example: O… 8(+) - 10(-) = 2- or O2-
Isotopes
Most elements in the first two rows of the periodic
table have at least two known isotopes.
 Isotopes have the same number of protons and electrons but
differ in the number of neutrons. (mass - protons = neutrons)
 Atoms in an element normally contain mixtures of isotopes in
specific ratios
 Isotopes are named after their masses (neutrons + protons)
 Example: hydrogen-1 (H-1), hydrogen-2 (H-2), hydrogen
(H-3)
Determining Average Atomic Mass
from the Isotopes Found in Nature
 Look at the three common isotopes of silicon. Multiply the
masses of the isotopes by their fractional abundances
(percent found in nature) and add the products together.
Element Symbol
Mass amu
Fractional Abundance
Contributing Mass
Si-28
Si-29
Si-30
27.977
92.21%
25.80
28.976 4.70%
1.36
29.974 3.09%
0.93
Average Atomic Mass 28.09amu
 Since most atoms exist in isotopes with known ratios, the
mass number expressed on the periodic table will usually
not agree exactly with the average amu number.
Relative Atomic Mass
The relative atomic mass of an atom is expressed in
atomic mass units (amu).
 This unit is derived from the carbon atom and is
measured to be 1/12 of the mass of the carbon-12 atom.
 The amu expresses the most common isotope of the
atom
 C-12 is more common than C-13 or C-14
Chemical Notation
 Chemical notation
 Displays atomic mass, atomic number and charge
 If a charge is not listed then the number of electrons is equal
to the number of protons
 In the example to the left
 11(+), 23-11 = (+), 10 (-)
 1+ charge tells you that there is 1 more proton than electrons