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Chapter 2
5 September 2013
2.1 Composition of the Atom
The Structure of the
Atom and the
Periodic Table
Selected Properties of the Three Basic
Subatomic Particles
Name
Charge
Mass(amu)
Mass (grams)
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)
1
2.1 Composition of the Atom
2.1 Composition of the Atom
Determining the Composition of an
Atom
Determine the number of protons, neutrons
and electrons in each of the following:
11
5
B
55
26
Fe
Isotopes
• Isotopes - atoms of the same element
having different masses
4
– contain same number of protons
– contain different numbers of neutrons
Isotopes of hydrogen
Hydrogen
(Hydrogen - 1)
Deuterium
(Hydrogen - 2)
Tritium
(Hydrogen - 3)
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2.1 Composition of the Atom
2.1 Composition of the Atom
Isotopic Calculations
• Isotopes of the same element have identical
chemical properties
• Some isotopes are radioactive
• Find chlorine on the periodic table
• What is the atomic number of chlorine?
17
• What is the mass given
35.45
• This is not the mass number of an isotope
Atomic Mass
• What is this number, 35.45?
• The atomic mass - the weighted average of
the masses of all the isotopes that make up
chlorine
• Chlorine consists of chlorine-35 and
chlorine-37 in about 3:1 ratio
• Weighted average is an average corrected
by the relative amounts of each isotope
present in nature
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2.1 Composition of the Atom
2.1 Composition of the Atom
Determining Atomic Mass
Calculate the atomic mass of naturally
occurring chlorine if 75.77% of chlorine
atoms are chlorine-35 and 24.22% 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 these partial weights to get the
weighted average atomic mass of chlorine:
26.52 amu + 8.965 amu = 35.48 amu
4
2.1 Composition of the Atom
2.1 Composition of the Atom
Atomic Mass Determination
• Nitrogen consists of two naturally occurring
isotopes
– 99.63% nitrogen-14 with a mass of 14.003 amu
– 0.37% nitrogen-15 with a mass of 15.000 amu
• What is the atomic mass of nitrogen?
Ions
• 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-
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2.2 Development of Atomic
Theory
• Dalton’s Atomic Theory - the first
experimentally based theory of atomic
structure of the atom.
2.2 Development of
Atomic Theory
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
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2.2 Development of
Atomic Theory
4. Atoms of different elements have
different properties
5. Atoms of different elements
combine in simple whole-number
ratios to produce compounds (stable
aggregates of atoms)
6. Chemical change involves joining,
separating, or rearranging atoms
2.2 Development of
Atomic Theory
Postulates 1, 4, 5 and 6 are still regarded
as true.
Evidence for Subatomic Particles:
Electrons, Protons and Neutrons
• Electrons were the first subatomic
particles to be discovered using the
cathode ray tube
7
2.2 Development of
Atomic Theory
Evidence for Protons and
Neutrons
• Protons were the next particle to be discovered, by
Goldstein
– 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
2.2 Development of
Atomic Theory
Evidence for the Nucleus
• Initial assumed protons and neutrons were
uniformly distributed throughout the atom
• Ernest Rutherford’s “Gold Foil
Experiment” lead to the understanding of
the nucleus
– Most alpha particles pass through the foil
without being deflected
– Some particles were deflected, a few even
directly back to the source
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2.3 Light, Atomic Structure,
and the Bohr Atom
2.2 Development of
Atomic Theory
Rutherford’s Gold Foil Experiment
• Most of the atom is empty space
• The majority of the mass is located in a
small, dense region
Models of the Atom
(a) Thomson
(b) Rutherford
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2.3 Light, Atomic Structure, and
the Bohr Atom
2.3 Light, Atomic Structure,
and the Bohr Atom
• Rutherford’s atom – tiny, dense, positively
charged nucleus of protons surrounded by
electrons
• How do we describe the relationship of the
electrons to each other and the nucleus?
• Use the measurement of particle energy
rather than position
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
10
2.3 Light, Atomic Structure,
and the Bohr Atom
Wavelengths
• Light is propagated (moves) as a collection
of sine waves
• Wavelength is the distance between identical
points on successive waves
• All wavelengths travel at the same velocity,
but have their own characteristic energy
Electromagnetic Spectrum
high energy
short wavelength
low energy
long wavelength
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2.3 Light, Atomic Structure,
and the Bohr Atom
2.3 Light, Atomic Structure,
and the Bohr Atom
Bohr Theory
• Atoms can absorb and emit energy via
promotion of electrons to higher energy
levels and relaxation to lower levels
• Energy that is emitted upon relaxation is
observed as a single wavelength of light
• Spectral lines are a result of electron
transitions between allowed levels in the
atoms
• The
emission
emission
spectrum
spectrum
- light
of hydrogen
emitted when
lead to
a the
substance
modernisunderstanding
excited by an of
energy
the
source.
electronic structure of the atom
12
The Bohr Atom
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
The Bohr Atom
13
Electronic Transitions
• Amount of energy absorbed in jumping from one
energy level to a higher energy level is a precise
quantity
• Energy of that jump is the energy difference
between the orbits involved
• Orbit - what Bohr called the fixed energy levels
• Ground state - the lowest possible energy state
Bohr Theory
•
•
•
•
Allowed levels are quantized energy levels, orbits
Electrons are found only in these energy levels
Highest-energy orbits are farthest from the nucleus
Atoms
– absorb energy by excitation of electrons to higher energy
levels
– release energy by relaxation of electrons to lower energy
levels
• Energy differences may be calculated from the
wavelength of light emitted
14
Modern Atomic Theory
• 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
The periodic law and the
periodic table
• Dmitri Mendeleev and Lothar Meyer - two
scientists working independently developed
the precursor to our modern periodic table.
• They noticed that elements have distinct
regular variation of their properties when
listed in order of atomic mass.
• Periodic law - the physical and chemical
properties of the elements are periodic
functions of their atomic numbers.
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