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Chapter 2
The composition
and Structure of
the Atom
Denniston
Topping
Caret
4th Edition
2.1 Matter and Structure
• Understanding the structure of the atom will
help to understand the properties of the
elements.
• Keep in mind that these, as all theories, are
subject to constant refinement. The picture
of the atom isn’t final.
1
2.2 Composition of the Atom
• Atom - the basic structural unit of an element.
– The smallest unit of an element that retains the
chemical properties of that element.
• Radioactive decay - certain kinds of atoms can
“split” into smaller particles and release large
amounts of energy.
• Atoms consist of three primary particles.
– electrons
– protons
– neutrons
• Nucleus - small, dense, positively charged region in
the center of the atom. Contains:
- protons - positively charged particles
- neutrons - uncharged particles
• Surrounding the nucleus is a diffuse region of
negative charge populated by:
- electrons - negatively charged particles
Table 2.1 Selected Properties of the Subatomic Particles
Name
Charge Mass (amu)
Electrons (e)
-1
5.4 x 10-4
9.1095 x 10-28
Protons (p)
+1
1.00
1.6725 X 10-24
0
1.00
1.6750 x 10-24
Neutrons (n)
Mass (grams)
Charge of
particle
Mass
Number
A
Z
Atomic
Number
X
C
3
Symbol of
the atom
atomic number (Z) - the number of protons
in the atom
mass number (A) - sum of the number of
protons and neutrons
• Isotopes - atoms of the same element
having different masses.
– contain same number of protons
– contain different numbers of neutrons
Isotopes of Hydrogen
Hydrogen
(Hydrogen - 1)
Deuterium
(Hydrogen - 2)
Tritium
(Hydrogen - 3)
• Isotopes of the same element have identical
chemical properties.
• Some isotopes are radioactive.
• Find chlorine on the periodic table.
• What is the atomic number? 17
• What is the mass given? 35.45
• This is not the mass number of an isotope.
• It is called the atomic mass - the weighed
average of the masses of the isotopes that make
up chlorine.
• Chlorine consists of chlorine-35 and chlorine37 in a 3:1 ratio.
• The weighted average is an average corrected
by the relative amounts of each isotope present
in nature.
• Calculate the atomic mass of naturally
occurring chlorine if 75.77% of chlorine
atoms are chlorine-35 and 24.23% of
chlorine atoms are chlorine-37.
Step 1: Convert the percentage to a decimal fraction.
0.7577 chlorine-35
0.2423 chlorine-37
Step 2: Multiply the decimal fraction by the mass of that
isotope to obtain the isotope contribution to the atomic
mass.
For chlorine-35:
0.7577 x 35.00 amu = 26.52 amu
For chlorine-37
0.2423 x 37.00 amu = 8.965 amu
Step 3: sum to get the weighted average
atomic mass of chlorine =
26.52 amu + 8.965 amu = 35.49 amu
• Ions - electrically charged particles that result
from a gain or loss of one or more electrons by
the parent atom.
• Cation - positively charged.
– result from the loss of electrons
– 23Na  23Na+ + 1e-
• Anion - negatively charged.
– results from the gain of electrons
– 19F + 1 e-  19F-
4
How many protons, neutrons and
electrons are in the following ions?
39
19
K

32
16
S
2-
24
12
Mg
2
2.3 Development of the Atomic
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Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom.
– John Dalton
– early 1800’s
• Much of Dalton’s Theory is still regarded as
correct today. *See starred items.*
Postulates of Dalton’s Atomic Theory
1.
All matter consists of tiny particles called atoms.*
2.
An atom cannot be created, divided, destroyed, or converted
to any other type of atom.
3.
Atoms of a particular element have identical properties.
4.
Atoms of different elements have different properties.*
5.
Atoms of different elements combine in simple wholenumber ratios to produce compounds (stable aggregates of
atoms.)*
6.
Chemical change involves joining, separating, or rearranging
atoms.*
*
These postulates are still regarded as true.
Subatomic Particles:
Electrons, Protons and Neutrons
• Electrons were the first subatomic particles to
be discovered using the cathode ray tube.
• Protons were the next particle to be discovered.
–
Protons have the same size charge but opposite in
sign.
–
Proton is 1837 times as heavy as electron.
• Neutrons
–
Postulated to exist in 1920’s but not
demonstrated to exist until 1932.
–
Almost the same mass as the proton.
The Nucleus
• The initial ideas of the atom did not have
a “nucleus.”
• “Plum Pudding Model”
• Earnest Rutherford’s “Gold Foil
Experiment” lead to the understanding of
the nucleus.
• Most of the atom is empty space.
• Most of the mass is located in a small,
dense region.
6
2.4 The Relationship between
Light and Atomic Structure
• Spectroscopy - absorption or emission of light by
atoms.
– Used to understand the electronic structure.
• To understand the electronic structure, we must
first understand light. Electromagnetic radiation
– travels in waves from a source
– speed of 3.0 x 108 m/s
Electromagnetic Spectrum
high energy
short wavelength
low energy
long wavelength
• emission spectrum - light emitted when a
substance is excited by an energy source.
•The emission spectrum of hydrogen lead to the
modern understanding of the electronic structure of
the atom.
2.5 The Bohr Atom
8
• Initial understanding of the atom by Niels Bohr
Electrons exist in fixed
energy levels
surrounding the nucleus.
Promotion of
electron occurs as
it absorbs energy
Energy is released as
the electron travels
back to lower levels.
Quantization of energy
Excited State
Relaxation
n=3
n=2
n=1
• Orbit - what Bohr called the fixed energy levels.
• Ground state - the lowest possible energy state.
• The orbits are also identified using “quantum
numbers”: n = 1, 2, 3, …
• When the electron relaxes (c) the energy released is
observed as a single wavelength of light.
2.6 Modern Atomic Theory
9
• Bohr’s model of the atom when applied to atoms
with more than one electron failed to explain their
line spectra.
• One major change from Bohr’s model is that
electrons do not move in orbits.
• Atomic orbitals - regions in space with a high
probability of finding an electron.
• Electrons move rapidly within the orbital giving a
high electron density. (More in Ch 3)