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Ch. 6
Electronic Structure
and the Periodic Table
Part 1: Light, Photon
Energies, and Emission
Spectra
• Compare the wave and particle natures of light.
• Define a quantum of energy, and explain how it is
related to an energy change of matter.
• Contrast continuous electromagnetic spectra and
atomic emission spectra.
radiation: the rays and particles —alpha particles,
beta particles, and gamma rays—that are emitted
by radioactive material
electromagnetic radiation
quantum
wavelength
Planck's constant
frequency
photoelectric effect
amplitude
photon
electromagnetic spectrum
atomic emission spectrum
Light, a form of electronic radiation,
has characteristics of both a wave and
a particle.
Review

We left off with Rutherford and Chadwick
discovering nucleus and neutrons
This proved that J.J. Thomson’s “plum
pudding” model of the atom was incorrect.


So where do we go next?
Rutherford’s model of the atom
Rutherford Model: Problems

Nucleus surrounded by electrons.
 How
did e- fill up space surrounding a (+)
nucleus?

What prevented the electrons from being
pulled right into the nucleus?

To answer this, must understand
relationship of light and electrons
The Atom and Unanswered Questions
(cont.)
• In the early 1900s, scientists observed
certain elements emitted visible light when
heated in a flame.
• Analysis of the emitted light revealed that an
element’s chemical behavior is related to the
arrangement of the electrons in its atoms.
At Rutherford Model

Electrons pictured as
particles

Light pictured as waves

Discovered electrons
have wave-like
properties, and light has
particle-like properties
What do we do?

Describe electrons as having
dual wave-particle nature (or properties)
 Sometimes
it acts like a particle
 Sometimes it acts like a wave

Has stood up against many experiments to
prove it wrong

Explains how electron isn’t pulled into
nucleus.
Electromagnetic (EM) radiation

Light as a wave
 Form
of Energy that exhibits wavelike
behavior as it travels
 Speed

= 3.00 x 108 m/s
(speed of light through air)
The Wave Nature of Light (cont.)
• The wavelength (λ) is the shortest
distance between equivalent points on a
continuous wave.
• The frequency (ν) is the number of waves
that pass a given point per second.
• The amplitude is the wave’s height from the
origin to a crest.
The Wave Nature of Light (cont.)
Wavelength
Wavelength : distance between corresponding
points on a wave
λ
= wavelength
λ is the Greek letter lambda
Wavelength is usually in nanometers (nm) or meters
Wavelength
http://en.wikipedia.org/wiki/Wavelength
Frequency

(ν), the Greek letter nu (not Vee)

Number of waves that pass a given point
in a specific amount of time

Frequency units are in Hertz (Hz) or 1/
seconds ( /s )
Relationship between wavelength
and frequency

c = λν
 Where

c = speed of light
Correlation?
 As
wavelength decreases, frequency increases
 As
wavelength increases, frequency decreases
 This
is an inverse relationship
The Wave Nature of Light (cont.)
• The speed of light (3.00  108 m/s) is the
product of it’s wavelength and frequency
c = λν.
What is the frequency of a wave
with wavelength of 100 nm?
The Particle Nature of Light
• The wave model of light cannot explain all
of light’s characteristics.
• Matter can gain or lose energy only in small,
specific amounts called quanta.
• A quantum is the minimum amount of energy
that can be gained or lost by an atom.
Max Planck (1900)

Described light as having particle-like
properties

When hot object loses energy, it doesn’t
do it continuously as it would if it were a
wave

Loses energy in form of a quanta
Quanta?

Quantum – finite quantity of energy that
can be gained or lost by an atom
 Specific:
if it costs $1.25 to get a soda from
machine, and you give it $1.00, do you get a
a soda?

Photon – individual quantum of light
The Particle Nature of Light (cont.)
• The photoelectric effect is when electrons
are emitted from a metal’s surface when
light of a certain frequency shines on it.
Einstein


In 1905, said Planck’s work applied to all EM.
Explains photoelectric effect –
 must
absorb photon with specific energy to dislodge
an electron
 When
electron is dislodged, it must be in the form of a
particle
 But
as it moves, (we see it as color), it is in the form of
a wave

Shows dual nature of light (wave and particle)
The Particle Nature of Light (cont.)
• Albert Einstein proposed in 1905 that light
has a dual nature.
• A beam of light has wavelike and particle-like
properties.
• A photon is a particle of electromagnetic
radiation with no mass that carries a quantum
of energy.
Ephoton = h
Ephoton represents energy.
h is Planck's constant.
 represents frequency.
Energy of a Photon (or any wave of
energy)

E=hν



E = energy ( in joules, j)
v = frequency
h = Planck's constant


6.63 x 10 -34 J * s (Joule Seconds)
As frequency goes up, what happens to the energy?
What is the energy with a wave of
frequency 1 x 1016 Hz?
Light through a prism

Continuous spectrum
 All
wavelengths in a given range are included
 Why we see rainbows
 Separated by wavelength

Electromagnetic spectrum
 Consists
of all electromagnetic radiation,
arranged by increasing wavelengths
Light through a prism
http://spaceplace.nasa.gov/en/kids/misrsky/misr_sky.shtml
Electromagnetic Spectrum
Hydrogen Atom Spectrum

Pass high voltage through Hydrogen gas

Gas glows, and you can pass this light through
prism

Creates a bright line spectrum or atomic
emission spectra
Atomic Emission Spectra (cont.)
Atomic Emission Spectra (cont.)
• The atomic emission spectrum of an
element is the set of frequencies of the
electromagnetic waves emitted by the
atoms of the element.
• Each element’s atomic emission spectrum is
unique.
Hydrogen’s atomic emission
spectra
 Each
line caused by light of a
different wavelength
 What
causes the light?
Electrons, man

Electrons get boosted by voltage from ground
state (or normal state) to excited state.

When they relax back down to ground (almost
immediately), they give off certain amounts of
energy

Line spectrum: produced when an electron
drops from a higher energy orbit to a lower
energy orbit
Ground State vs. Excited State

Ground State
 The

state of lowest energy of an atom
Excited State
A
state in which an atom has a higher
potential energy than its ground state
What causes the lines?

Each line = energy from electron as it
drops from excited state to ground state

Energy of photon = difference in energy
between ground and excited state.