Atomic Structure
Subatomic Particles
Atoms are made up of even smaller particles called subatomic particles. These subatomic
particles are: electrons, protons, and neutrons. [There are other subatomic particles, but
from a chemical perspective, they are rather unimportant.]
Symbol, charge, and approximate mass of subatomic particles
Particle
Symbol
Electrical Charge
Approximate Mass
+
Proton
P
+1
1 amu
Neutron
N
0
1 amu
Electron
e
-1
0 amu
-24
amu=atomic mass unit
1 amu=1.66 x 10 g
Arrangement of subatomic particles
The nucleus contains the protons and neutrons. The nucleus is very small in comparison to
the size of the atom, yet it is very heavy. Most of the mass of the atom is found inside
the nucleus. If the nucleus was the size of a marble, the atom would be about the size of
a football field.
The electrons are found outside the nucleus in certain energy levels.
Atomic Number (Z) and Mass Number (M)
The atomic number is what determines the atom’s identity.
Atomic number = number of protons in an atom
For atoms that are electrically neutral, the number of protons = the number of electrons
For ions, the number of protons ≠ the number of electrons
Cations have less electrons than an electrically neutral atom
Anions have more electrons than an electrically neutral atom
Mass number = number of protons + number of neutrons
So, number of neutrons = mass number – number of protons
Isotopes
Isotopes are atoms of the same element that have the same number of protons but
different number of neutrons. Because isotopes of an element have different numbers of
neutrons, they also have different mass numbers.
To distinguish between isotopes of the same element, we write an isotope symbol to
indicate the mass number and the atomic number of the atom. Example, the isotope
symbol for one isotope of argon is:
Examples:
1. Write an isotope symbol for chlorine-35 and chlorine-37
2.
Which of the following is an isotope of boron-12,
or
?
Atomic Mass Scale
The atomic mass scale is based on the mass of the carbon-12 isotope. All masses are
determined relative to the mass of carbon-12, which by definition is exactly 12 amu.
The atomic mass of an element is the weighted average mass of the atoms in a naturally
occurring sample of the element. A weighted average mass reflects both the mass and the
relative abundance of the isotopes as they occur in nature.
Examples:
1. Calculate the average atomic mass of copper. Given the following information:
Isotope
% Abundance
63
Cu
69.1
65
Cu
30.9
2. Antimony has two common isotopes. If one of the isotopes is antimony-121 and has
an abundance in nature of 57.25%, what is the atomic mass (to 3 sig figs) of the
other isotope?
From the periodic table we know that the average atomic mass of antimony is
121.75 amu.
3.
The atomic mass of fictitious element X is 251.7 amu. If element X consists of
two isotopes that have mass numbers of 250 and 253, what is the approximate %
natural abundance of each isotope?