Zumdahl • Zumdahl • DeCoste
World of
CHEMISTRY
Chapter 11
Modern Atomic
Theory
Chapter 11 Overview
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Describe Rutherford’s model of atom
Electromagnetic radiation
See how atoms emit light
Quantized nature of energy demonstrated by
emission spectrum of hydrogen
Bohr’s model of hydrogen atom
Wave mechanical model of electron position
Shapes of s, p, and d orbitals
Electron spin
Electrons filling principle energy levels
Valence and core electrons
Electron configurations
Periodic table trends
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11-3
Figure 11.1: The Rutherford atom.
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Rutherford Atom Review
• Alpha particle/Gold foil experiment
• Nuclear Atom
• Nucleus composed of protons & neutrons
• Nucleus small compared to atomic size
• Electrons account for rest of atom
• Unanswered questions:
• What are electrons doing? – How are they
arranged & how do they move?
• Thought electrons revolved around nucleus like
planets orbit the sun
• Couldn’t explain why electrons aren’t attracted to
protons causing atom to collapse
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11-5
Energy and Light
• Electromagnetic Radiation
• Transmits energy
• Heat from light bulb
• Solar energy (energy from sun)
• Warmth from fireplace
• Many kinds: X-rays, microwaves, etc.
• Differ in their wave characteristics
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Figure 11.2: A seagull floating on the ocean moves
up and down as waves pass.
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Wave Properties
• Wavelength (λ): distance between two
consecutive wave peaks
• Frequency (ν): how many waves pass
a certain point per given time period
• Speed: how fast a given peak travels
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Figure 11.3: The wavelength of a wave.


Crest
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Trough
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Electromagnetic Radiation
• Travel as waves
• Have different wavelengths
• See page 325
• Gamma rays – shortest, radio waves – longest
• Important means of energy transfer
• Solar energy – visible & ultraviolet radiation
• Heat from fireplace – infrared radiation
• “Light” – wave that carries energy through space
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Electromagnetic radiation has particle
characteristics
• Photons: tiny packets of energy that
travel in a stream
• Wave-particle nature of light: consists
of both waves and particles of energy
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Figure 11.5: Electromagnetic radiation.
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Figure 11.6: Photons of red and blue light.
Different wavelengths of electromagnetic radiation carry
different amounts of energy
In general – the longer the wavelength, the lower the energy
of the photons (red less energy than blue)
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11-13
Emission of Energy by Atoms
• Recall Flame Test Laboratory (different
elements gave off different colors)
• Color resulted from atoms releasing energy
by emitting visible light of specific
wavelengths (specific colors)
• Atoms became excited: absorbed heat
energy from flame
• Some of excess energy released as light –
carried away by photon
• Energy of photon = energy change of atom
• Short wavelength = high-energy photons
• Long wavelength = low-energy photons
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11-14
Figure 11.8: An excited lithium atom emitting
a photon of red light to drop to a lower
energy state.
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11-15
The Energy Levels
of Hydrogen
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Figure 11.9: A sample of H atoms receives
energy from an external source.
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Figure 11.9: The excited atoms release
energy by emitting photons.
Excited atom can release some or all of its
energy by emitting a photon (electromagnetic
radiation “particle”)
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11-18
Figure 11.10: An excited H atom returns to a
lower energy level.
Energy contained in photon = change in energy of atom
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11-19
Figure 11.11: Colors and wavelengths of
photons in the visible region.
Visible light photons emitted by Hydrogen – always the same
Because only certain photons are emitted, only certain energy
changes are occurring
Hydrogen atom has certain discrete energy levels
Energy levels of Hydrogen are quantized – only certain
values allowed
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Figure 11.12: The color of the photon emitted
depends on the energy change that produces it.
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Figure 11.13: Each photon emitted corresponds to
a particular energy change.
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Figure 11.14:
Continuous (a) and
discrete (b) energy
levels.
Quantized nature of energy
surprised scientists (b)
Previously assumed atom
could exist at any energy
level (a)
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Figure 11.15: The difference between continuous
(a) and quantized (b) energy levels.
Ramp – can be
at any elevation
Staircase – can move
from one step to another
or even skip, but must
be on a step
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11-24
Figure 11.17: The Bohr model of the hydrogen
atom.
Electrons moved in
circular orbits like
planets
Electrons could jump
from one orbit to another
by emitting/absorbing a
photon
Didn’t work for other
atoms
Showed
experimentally to be
incorrect
Paved way for other
theories
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11-25
We do not know exactly
how the electrons move in
an atom!
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11-26
The Wave Mechanical Model of the Atom
• Louis Victor de Broglie & Erwin
Schrödinger: since light has both wave
and particle characteristics, an electron
might also exhibit these characteristics
• New hydrogen model applied to other
atoms (Bohr’s did not)
• Electron states are described by orbitals
(which are nothing like orbits)
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11-27
Figure 11.18: A representation of the photo of the
firefly experiment (lightning bugs).
Shows probability (or
likelihood) of where
firefly will be found
Usually near the
center, but can be
found in any of the
shaded areas at any
time
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11-28
Figure 11.19:
The orbital that
describes the
hydrogen
electron
in its lowest
possible energy
state.
Darker pink =
greater
probability
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11-29
Drawbacks of wave mechanical model:
• Gives no information about when the
electron occupies a certain point in
space or how it moves
• We will probably never know the details
of electron motion
• Confident that Bohr model is incorrect
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11-30
Figure 11.20: The hydrogen 1s orbital. (Lowest
Energy State or Ground State)
1s orbital – spherical
Probability map (more accurate)
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11-31
representation
The Hydrogen Orbitals
• Size defined as the sphere that contains 90%
of the total electron probability
• Spends 90% of its time somewhere within the
sphere
• Spends 10% of its time somewhere outside of the
sphere
• Electron can absorb energy & move to higher
energy state
• Bohr model – orbit with larger radius
• Wave mechanical model – different kinds of
orbitals with different shapes
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Figure 11.21: The first four principle energy levels in
the hydrogen atom.
Further from nucleus
Closer to nucleus
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Figure 11.22: How principal levels can be
divided into sublevels.
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Figure 11.23: Principal level 2 shown divided into
the 2s and 2p sublevels.
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Figure 11.24: The relative sizes of the 1s and 2s
orbitals of hydrogen.
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Figure 11.25: The three 2p orbitals.
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Figure 11.26: Diagram of principal energy
levels 1 and 2.
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Orbital Labels
• Number tells principal energy level
• Letter indicates shape
• s = spherical
• p = two-lobed (x, y, & z indicates axis)
Principle
energy
level 2
2s orbital
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Spherical shape
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Hydrogen Orbitals
• Why does hydrogen have more than
one orbital if it only has 1 electron?
• Orbital is potential space for an electron
• Can only occupy 1 orbital at a time, but can
be transferred to another by adding energy
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11-40
Figure 11.27: Relative sizes of the spherical 1s, 2s,
and 3s orbitals of hydrogen.
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11-41
Orbital summary
• As principle energy level increases, the
number of sublevels increases
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n = 1 – 1 sublevel
n = 2 – 2 sublevels
n = 3 – 3 sublevels , etc.
Further from nucleus = more space = more room
for orbitals
• Sublevels are s (1 orbital), p (3 orbitals), d (5
orbitals), and f (7 orbitals)
• Orbitals keep same shape, but get larger as n
increases
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Figure 11.28: The shapes and labels of the
five 3d orbitals.
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11-43
Further Development of the Wave Mechanical
Model
• Applies to all atoms
• Helps explain the periodic table
• Electrons spin like a top – can only spin in
one of two directions
• Use arrows to represent spin (↑or↓)
• Electrons must have opposite spins to occupy
the same orbital
• Pauli exclusion principle: an atomic orbital
can hold a maximum of two electrons and
those two electrons must have opposite spins
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Principle components of wave mechanical model
• Atoms have principal energy levels (n)
• Energy of level increases as n increases
• Each principal energy level contains one or
more types of orbitals, called sublevels
• Number of sublevels present = n (p. 338)
• Label with n and shape (ex.: 3p)
• Orbital can have 0 to 2 electrons, 2 electrons in
same orbital must have opposite spins
• Shape of orbital indicates probabilities, not
electron movement
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11-45
Electron Arrangements of First 18 Elements
Electrons will occupy orbitals closest to
nucleus first
As n increases, orbital becomes larger –
electron is further from nucleus
Electron configuration = electron
arrangement (Example: 1s22s1)
Abbreviate Na: 1s22s22p63s1 = [Ne]3s1
Orbital Diagram = Box Diagram: orbitals are
represented by boxes grouped by sublevel
with arrows indicating electrons
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Electron Configurations and Orbital Diagrams
Atom
Configuration
Diagram
Hydrogen
1s1
1s
Helium
1s2
1s
Beryllium
1s22s2
1s
2s
↑↓
↑↓
Carbon
1s22s22p2
1s
2s
↑↓
↑↓
Oxygen
1s22s22p4
1s
2s
↑↓
↑↓
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11-47
↑
↑↓
2p
↑
↑
2p
↑↓
↑
↑
Valence Electrons
• Electrons in the outermost (highest) principal
energy level of an atom
• Most important electrons to chemists –
electrons involved when atoms form bonds
(attach to each other)
• Atoms in same group on periodic table have
same number of valence electrons in outer
orbital (orbitals are at different principal
energy levels)
• Core electrons
• Inner electrons
• Not involved in bonding
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Figure 11.30: Partial electron configurations
for the elements potassium through krypton.
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Figure 11.31: Orbitals being filled for elements in
various parts of the periodic table.
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11-50
Orbital Filling
• If energy level has d orbitals, s orbitals
from next level will fill first
• After lanthanum – lanthanide series – fill
4f orbitals
• After actinium – actinide series – fill 5f
orbitals
• Except Helium – group number
indicates sum of electrons in outer s & p
orbitals (number of valence electrons)
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11-51
Figure 11.34:
Periodic table
with atomic
symbols,
atomic
numbers,
and partial
electron
configurations.
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Atomic Properties & the Periodic Table
• Chemistry is fundamentally based on
observed properties of substances
• Atomic theory is attempt to help us
understand why these things occur
• Theories may change
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Figure 11.35: Classification of elements as
metals, nonmetals, and metalloids.
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Metals, Nonmetals, & Metalloids
• Metals
• Lustrous appearance, change shape without
breaking (pulled into wire), excellent conductors
• Tend to lose electrons to form positive ions
• Nonmetals
• Lack properties of metals, some exceptions
• Tend to gain electrons to form negative ions
• Metalloids
• Have properties of metals and nonmetals
• Along stair step
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Not all metals/nonmetals behave exactly the
same way
• Metals
• As you go down group – more likely to lose
electrons (further from nucleus)
• Most chemically active – lower left corner
of periodic table
• Nonmetals
• Most chemically active in upper right
corner (not noble gases)
• Strongest attraction (closer to nucleus)
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Figure 11.36: Relative atomic sizes for selected
atoms.
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Atomic Size
• Increases as you go down group: more
electrons = larger atom
• Decreases as you go across period:
more protons in nucleus – greater pull
on electron
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Ionization Energy
• The energy required to remove an electron
from an individual atom in the gas phase
• Metals relatively low – easily lose electrons
(small amount of energy needed)
• Nonmetals relatively large – prefer to gain
electron, not lose
• Decreases going down a group, increases
going across a period
• Bottom left – lowest (most chemically active)
• Upper right – highest (most chemically active)
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