Chapter 4
Atomic Structure
I.
Defining the Atom
An atom is the smallest particle of an element that retains its identity in a
chemical reaction.
A.
Early philosophers believed that atoms were indivisible and
indestructible.
B.
Dalton’s Atomic theory. Dalton used experimental methods, to
transform Democritus’s ideas on atoms into scientific theory.
1. All elements are composed of tiny particles called atoms.
2. Atoms of the same element are identical. The atoms of any one
element are different from those of any other element.
3. Atoms of different elements can physically mix together or can
chemically combine in simple whole number ratios to form
compounds.
4. Chemical reactions occur when atoms are separated, joined, or
rearranged. Atoms of one element, however, are never changed
into atoms of another element as a result of a chemical reaction.
C.
Sizing up the Atom
Despite their small size, individual atoms are observable with instruments
such as a scanning tunneling microscope.
II.
Structure of the Nuclear Atom
Atoms are divisible and are composed of subatomic particles.
A.
Electrons are negatively charged particles. Electrons are located
outside the nucleus of the atom. Electrons have very little mass when
compared to the mass of protons and neutrons. Electrons are
represented by e-. Electrons have a mass of 9.11 x 10-28 g.
B.
Protons are positively charged particles and are located inside the
nucleus. Protons have a mass of 1.67 x 10-24 g. This means that a
proton is about 1840 times the mass of an electron. Protons are
represented by p+.
C.
Neutrons are neutral particles and are located inside the nucleus.
Neutrons have no effective charge and are therefore considered to be
neutral. Neutrons have a mass of 1.67 x 10-24 g. Neutrons are
represented by n0.
D.
Atomic Nucleus
1. Most of the volume in an atom is occupied by electrons.
Rutherford’s Gold foil experiment proved that the nucleus is very
small relative to the total volume of the atom.
2. The nucleus is the central core of an atom and it is composed of
protons and neutrons.
III.
Distinguishing Among Atoms
A.
The number of protons present in an atom determines which atom you
have. For this reason the number of protons is also called the atomic
number.
B.
The mass of an atom is determined by the number of protons and
neutrons. Example: if you have Carbon with six protons and six
neutrons the mass number would be 12. The number of neutrons in an
atom is the difference between the mass number and the atomic
number.
C.
Not all elements are the same. Since the number of protons controls
which element you have, what happens when you have a different
number of neutrons? Isotopes are what we say you have when you
have the same element with a different number of neutrons.
Example:
Protons
Neutrons
Mass
Carbon
6
6
12
Carbon
6
7
13
Carbon
6
8
14
In this case all three are Carbon atoms, but they have different
masses and are called isotopes.
D.
The mass system for atoms is called an atomic mass unit (amu). This
system of mass based on the mass of Carbon-12. An amu is equal to
exactly on twelfth the mass of Carbon-12.
1. A proton has a mass of 1 amu.
2. A neutron has a mass of 1 amu.
3. An electron has a mass of 0 amu.
E.
The mass for the elements listed on the periodic table takes account all
of the naturally occurring isotopes for an element and their relative
abundances. This is called the weighted mass average.
Example:
Carbon has two main isotopes
Abundance
12
98.89%
6C
13
1.11%
6C
To solve the weighted take the mass times its abundance
(12).9889+(13)0.0111 = 12.0111
1. Chlorine
35
17Cl
37
17Cl
2.
6
3Li
7
3Li
Isotope Problems
Abundance
75.77%
24.23%
Solve for the weighted mass average.
7.5%
92.5%
Solve for the weighted mass average.
3. Boron
10
5B
11
5B
19.6%
80.4%
Solve for the weighted mass average.
4. Magnesium
3 isotopes
24
Mg
78.70%
12
25
10.13%
12Mg
26
Mg
11.17%
12
5.
28
14Si
29
14Si
30
14Si
Solve for the weighted mass average.
92.21%
4.70%
3.09%