Download Chapter 5

Introductory Chemistry:
Concepts & Connections
4th Edition by Charles H. Corwin
Chapter 5
Models of
The Atom
Christopher G. Hamaker, Illinois State University, Normal IL
© 2005, Prentice Hall
Dalton Model of the Atom
• John Dalton proposed that all matter is made up
of tiny particles.
• These particles are molecules or atoms.
• Molecules can be broken down into atoms by
chemical processes.
• Atoms cannot be broken down by chemical or
physical processes.
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Dalton’s Model
• According to the law of definite composition, the
mass ratio of carbon to oxygen in carbon dioxide
was always the same. Carbon dioxide was
composed 1 carbon atom and 2 oxygen atoms.
• Similarly, 2 atoms of hydrogen and 1 atom of
oxygen combine to give water.
• Dalton proposed that 2 hydrogen atoms could
substitute for each oxygen atom in carbon dioxide
to make methane with 1 carbon atom and 4
hydrogen atoms. Indeed, methane is CH4!
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Dalton’s Theory
• A Summary of Dalton’s Atomic Theory:
1) An element is composed of tiny, indivisible,
indestructible particles called atoms.
2) All atoms of an element are identical and have the
same properties.
3) Atoms of different elements combine to form
compounds.
4) Compounds contain atoms in small whole number
ratios.
5) Atoms can combine in more than one ratio to form
different compounds.
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Dalton’s Atomic Theory
• The first two parts of Dalton’s theory were later
proven incorrect.
• Proposals 3, 4, and 5 are still accepted today.
• Dalton’s theory was an important step in the
further development of atomic theory.
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Subatomic Particles
• About 50 years after Dalton’s proposal, evidence
was seen that atoms were divisible.
• Two subatomic particles were discovered.
– Negatively charged electrons, e–.
– Positively charge protons, p+.
• An electron has a relative charge of -1 and a
proton has a relative charge of +1.
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Thomson Model of the Atom
• J.J. Thomson proposed a subatomic model of the
atom in 1903.
• Thomson proposed that
the electrons were
distributed evenly
throughout a homogeneous
sphere of positive charge.
• This was called the “Plum
Pudding” Model of the
atom.
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Mass of Subatomic Particles
• Originally, Thomson could only calculate the
mass-to-charge ratio of a proton and an electron.
• Robert Millikan determined the charge of an
electron in 1911.
• Thomson calculated the masses of a proton and
electron:
– An electron has a mass of 9.11 × 10-28 g
– A proton has a mass of 1.67 × 10-24 g
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Types of Radiation
• There are three types of radiation:
– Alpha (a), Beta (b), & Gamma (g)
• Alpha rays are composed of helium atoms
stripped of their electrons (helium nuclei).
• Beta rays are composed of electrons.
• Gamma rays are high energy electromagnetic
radiation.
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Rutherford’s Gold Foil Experiment
• Rutherford fired alpha particles at thin gold foils.
If the “plum pudding” model of the atom was
correct, most a-particles should pass through
undeflected.
• However, some of the alpha
particles were deflected
backwards.
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Explanation of Scattering
• Most of the alpha particles passed through the foil
because an atom is largely empty space.
• At the center of an atom is the atomic nucleus
which contains the atom’s protons.
• The a-particles that
bounced backwards
did so after striking
the dense nucleus.
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Rutherford’s Model of the Atom
• Rutherford proposed a new model of the atom:
– The negatively charges electrons are distributed
around a positively charged nucleus.
• An atom has a diameter of
about 1 × 10-8 cm and the
nucleus has a diameter of
about1 × 10-13 cm.
• If an atom were the size
of the Astrodome, the
nucleus would be a
marble.
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Subatomic Particles Revisited
• Based on the heaviness of the nucleus, Rutherford
predicted that it must contain neutral particles in
addition to protons.
• Neutrons, n0, were discovered about 30 years
later. A neutron is about the size of a proton
without any charge.
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Atomic Notation
• Each element has a characteristic number of
protons in the nucleus. This is the atomic
number, Z.
• The total number of protons and neutrons in the
nucleus of an atom is the mass number, A.
• We use atomic notation to display the number of
protons and neutrons in the nucleus of an atom:
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Using Atomic Notation
• An example:
• The element is sodium (symbol Na).
• The atomic number is 11 – sodium has 11 protons.
• The mass number is 23 – the atom of sodium has
23 protons + neutrons.
• The number is neutrons is: A – Z = 23 – 11 =
12 neutrons.
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Isotopes
• All atoms of the same element have the same
number of protons.
• Most elements occur naturally with varying
numbers of neutrons.
• Atoms of the same element that have a different
number of neutrons in the nucleus are called
isotopes.
• Isotopes have the same atomic number but
different mass numbers.
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Isotopes Continued
• We often refer to an isotope by stating the name of
the element followed by the mass number.
– Cobalt-60 is
60
37
– Carbon-14 is
14
6
Co
C
• How many protons and neutrons does an atom of
mercury-202 have?
– The atomic number of Hg is 80, so it has 80 protons
– Hg-202 has 202 – 80 = 122 neutrons
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Simple & Weighted Averages
• A simple average assumes the same number of
each object.
• A weighted average takes into account the fact
that we do not have equal numbers of all the
objects.
• A weighted average is calculated by multiplying
the percentage of the object (as a decimal number)
by its mass for each object and adding the
numbers together.
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Average Atomic Mass
• Since not all isotopes of an atom are present in
equal proportions, we must use the weighted
average.
• Copper has two isotopes:
– 63Cu with a mass of 62.930 amu and 69.09% abundance
– 65Cu with a mass of 64.928 amu and 30.91% abundance
• The average atomic mass of copper is:
– (62.930 amu)(0.6909) + (64.928 amu)(0.3091)
= 63.55 amu
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Periodic Table
• We can use the periodic table to obtain the atomic
number and atomic mass of an element.
• The periodic table shows the atomic number,
symbol, and atomic mass for each element.
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Wave Nature of Light
• Light travels through space as a wave, similar to
an ocean wave.
• Wavelength is the distance light travels in one
cycle.
• Frequency is the number of wave cycles
completed each second.
• Light has a constant speed: 3.00 × 108 m/s.
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Wavelength vs. Frequency
• The longer the wavelength of light, the lower the
frequency.
• The shorter the wavelength of light, the higher the
frequency.
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Visible Spectrum
• Light usually refers to radiant energy that is
visible to the human eye.
• The visible spectrum is the range of wavelengths
between 400 and 700 nm.
• Radiant energy that has a wavelength lower than
400 nm and greater than 700 nm cannot be seen
by the human eye.
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Radiant Energy Spectrum
• The complete radiant energy spectrum is an
uninterrupted band, or continuous spectrum.
• The radiant
energy
spectrum
includes most
types of
radiation,
most of which
are invisible
to the human
eye.
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The Wave/Particle Nature of Light
• In 1900, Max Planck proposed that radiant energy
is not continuous, but is emitted in small bundles.
This is the quantum concept.
• Radiant energy has both a wave nature and a
particle nature.
• An individual
unit of light
energy is a
photon.
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The Quantum Concept
• The quantum concept states that energy is present
in small, discrete bundles.
• For example:
– A tennis ball that rolls down a ramp loses potential
energy continuously.
– A tennis ball that rolls down a staircase loses potential
energy in small bundles. The loss is quantized.
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Bohr Model of the Atom
• Niels Bohr speculated that electrons orbit about
the nucleus in fixed energy levels.
• Electrons are found only in specific energy levels,
and nowhere else.
• The electron
energy levels
are quantized.
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Emission Line Spectra
• When an electrical voltage is passed across a gas
in a sealed tube, a series of narrow lines is seen.
• These lines are the emission line spectrum. The
emission line spectrum for hydrogen gas shows
three lines: 434 nm, 486 nm, and 656 nm.
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Evidence for Energy Levels
• Bohr realized that this was the evidence he needed
to prove his theory.
• The electric charge temporarily excites an electron
to a higher orbit. When the electron drops back
down, a photon is given off.
• The red line is the
least energetic and
corresponds to an
electron dropping
from energy level 3
to energy level 2.
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“Atomic Fingerprints”
• The emission line spectrum of each element is
unique.
• We can use the line spectrum for the identify of
elements, using their “atomic fingerprint”.
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Energy Levels and Sublevels
• It was later shown that electrons occupy energy
sublevels within each level.
• These sublevels are given the designations
s, p, d, and f.
– These designations are in reference to the sharp,
principal, diffuse, and fine lines in emission spectra.
• The number of sublevels in each level is the same
as the number of the main level.
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Energy Levels and Sublevels
• The first energy level has 1 sublevel:
– 1s
• The second energy level has 2 sublevels:
– 2s and 2p
• The third energy level has 3 sublevels:
– 3s, 3p, and 3d
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Electron Occupancy in Sublevels
• The maximum number of electrons in each of the
energy sublevels depends on the sublevel:
– The s sublevel holds a maximum of 2 electrons.
– The p sublevel holds a maximum of 6 electrons.
– The d sublevel holds a maximum of 10 electrons.
– The f sublevel holds a maximum of 14 electrons
• The maximum electrons per level is obtained by
adding the maximum number of electrons in each
sublevel.
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Electrons per Energy Level
Chapter 5
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Electron Configurations
• Electrons are arranged about the nucleus in a
regular manner. The first electrons fill the energy
sublevel closest to the nucleus.
• Electrons continue filling each sublevel until it is
full and start filling the next closest sublevel.
• A partial list of sublevels in order of increasing
energy is:
 1s < 2s < 2p < 3s < 3p < 4s < 3d < 4p < 5s < 4d …
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Filling Diagram for Sublevels
• The order does
not strictly follow
1, 2, 3, etc.
• For now, use
Figure 5.16 to
predict the order
of sublevel filling.
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Electron Configurations
• The electron configuration of an atom is a
shorthand method of writing the location of
electrons by sublevel.
• The sublevel is written followed by a superscript
with the number of electrons in the sublevel.
– If the 2p sublevel contains 2 electrons, it is written 2p2
• The electron sublevels are arranged according to
increasing energy.
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Writing Electron Configurations
• First, determine how many electrons are in the
atom. Iron has 26 electrons.
• Arrange the energy sublevels according to
increasing energy:
– 1s 2s 2p 3s 3p 4s 3d …
• Fill each sublevel with electrons until you have
used all the electrons in the atom:
– Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d 6
• The sum of the superscripts equals the atomic
number of iron (26)
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Chapter 5
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Elements 1-18
Chapter 5
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Elements 19-36
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Elements 37-84
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Quantum Mechanical Model
• An orbital is the region of space where there is a
high probability of finding an atom.
• In the quantum mechanical atom, orbitals are
arranged according to their size and shape.
• The higher the energy of an orbital, the larger its
size.
• s-orbitals have
a spherical
shape.
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Cross Section of
1s, 2s, 3s Orbitals
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Shapes of p-Orbitals
• Recall that there are three different p sublevels.
• p-orbitals have a dumbbell shape.
• Each of the p-orbitals has the same shape, but
each is oriented along a different axis in space.
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p Orbitals
Chapter 5
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Shapes of d-Orbitals
There are five different d sublevels.
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d Orbitals
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First of Seven f Orbitals
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Conclusions
• Atoms are composed of protons, neutrons, and
electrons.
• The protons and neutrons are located in the
nucleus and the electrons are outside the nucleus.
• Atoms are mostly empty space.
• The number of protons is referred to as the atomic
number for the atom.
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Conclusions Continued
• All atoms of the same element have the same
number of protons.
• Isotopes are atoms with the same number of
protons but differing numbers of neutrons.
• The mass number for an isotope is the total
number of protons plus neutrons.
• The atomic mass of an element is the weighted
average of the masses of all the naturally occuring
isotopes.
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Conclusions Continued
• Light has both the properties of waves and
particles.
• The particles of light are referred to as photons.
• The energy of photons is quantized.
• Electrons exist around the nucleus of atoms in
discrete, quantized energy levels.
• Electrons fill energy sublevels starting with the
lowest energy sublevel and filling each successive
level of higher energy.
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