Electrons in Atoms
Chapter 4
The New Atomic Model
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Investigations  relationship between light
and atom’s electrons
How are electrons arranged? Why don’t
they fall into the nucleus?
Light a wave or particle?
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Wave Description:
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Electromagnetic Radiation: energy that
acts like a wave in space
All forms create Electromagnetic Spectrum
Electromagnetic Spectrum
Electromagnetic Spectrum
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All waves move at speed of light, c, 3.00x108 m/s
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Waves identified by:
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wavelength, , the distance b/
corresponding points on adjacent waves.
Units: nm, cm, or m
Frequency, , # of waves that pass a given
point in a specific time, 1 sec. Unit: 1/s =
Hertz, Hz
Wavelength and Frequency
Wavelength and Frequency
c = 
speed of
light, m/s

wavelength,
m
Inverse proportion equation!!
Frequency,
1/s
Calculation

Calculate the wavelength of a radio wave with
a frequency of 102.7 x 106s-1

Determine the frequency of light whose
wavelength is 5.267 nm.
Particle Nature of Light

Photoelectric
Effect: emission of
electrons from a
metal when light
shines on the metal
Photoelectric Effect
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Light had to be certain frequency to knock eloose
Light must also be a particle!
Max Planck(1900) explanation: objects emit
energy in small packets called quanta.
A photon is a single quantum of
(visible) light as well as a single quantum of
all other forms of electromagnetic radiation,
and can be referred to as a "light quantum".
Max Planck

Quantum of energy is the smallest amount of
energy that can be lost or gained by an atom
E = h
Energy of
quantum,
in joules, J
Frequency,
s-1
Planck’s
constant,
6.626x10-34 Js
Energy Calculation

What is the energy of green light, with a
wavelength of 500. nm?
1 nm =
10-9
m
500. nm = 5.00 x 10-7 m
Electromagnetic Spectrum
What is the color of an electromagnetic radiation with an
energy of 3.4 x10-19 J?
Photoelectric Effect
Albert Einstein
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Light is both wave and particle!
Particle of light = photon, having zero mass and
a quantum of energy
Photons hit metal and knock e- out, but photon
has to have enough energy
H-atom Emission Spectrum
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Pass a current through gas at low pressure
it excites the atoms
Ground state: lowest energy state of an
atom
Excited state: atom has higher potential
energy than it has in ground state
H – Atom Spectrum

When atom jumps from excited state to
ground state it gives off energy  LIGHT!
E2
Ephoton = E2 – E1 = hv
E1
Bohr’s Model of the H Atom (and only H!)
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Applied quantization of energy transfer to the atomic model
Studied atomic spectrum of H to come up with atomic
model.
Atomic emission spectra:
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Most sources produce light that contains many wavelengths
at once.
However, light emitted from pure substances may contain
only a few specific wavelengths of light called a line
spectrum (as opposed to a continuous spectrum).
Atomic emission spectra are inverses of atomic
absorption spectra.
18
Atomic Emission Spectra of C and H
Hydrogen: contains 1 red, 1 green, 1 blue and 1 violet.
Carbon: Contains many more emission lines as compared to H. Why?
19
H-atom Line Emission Spectrum
Element Emission Spectras
Helium – 23 lines
Neon – 75 lines
Argon - 159 lines
Xenon – 139 lines
Mercury – 40 lines
H-atom Line Emission Spectrum
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More lines in UV (Lyman series) and
IR(Paschen series)
Why did H-atom only emit certain colors
of light?
Bohr Model of H-atom
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1913 – Niels Bohr
e- circles nucleus in certain paths, orbits or
atomic energy levels
e- is higher in energy the farther away from
nucleus
e- cannot be between orbits
Video - 23
Bohr Model of H-atom
Bohr Model of H-atom
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From wavelengths of emission spectrum
Bohr calculated energy levels of H-atom
Model worked ONLY for H-atom
End Part 1