Transcript Chapter 4: Arrangement of Electrons in Atoms

Chapter 4:
Arrangement of Electrons in Atoms
Coach Kelsoe
Chemistry
Pages 97-122
Section 4-1:
The Development of a New
Atomic Model
Coach Kelsoe
Chemistry
Pages 97 – 103
The New Atomic Model
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The previous models we
studied did not explain
where electrons were
located.
A new atomic model
revealed a relationship
between light and an
atom’s electrons.
We used to think of light
as a wave. Light acts like
a wave, but has particlelike characteristics as well.
The Wave Description of Light
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Visible light is a type of
electromagnetic radiation, a form
of energy that exhibits wavelike
behavior as it travels through space.
Gamma rays, x-rays, ultraviolet,
visible light, infrared, microwaves,
and radio waves make up the
electromagnetic spectrum.
The Wave Description of Light
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Wave motion involves wavelength and
frequency.
Wavelength is the distance between
corresponding points on adjacent waves.
Frequency is the number of waves that pass
a given point in a specific time, usually one
second. Frequency is measured in hertz.
HIGH FREQUENCY
LOW FREQUENCY
Frequency and Wavelength
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Frequency and
wavelength are related
mathematically:
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c= λv, where c= speed of
light, λ= wavelength, and
v= frequency of the
electromagnetic wave
Since c is a constant,
the product of λv is
constant too.
Photoelectric Effect
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The photoelectric effect refers to the
emission of electrons from a metal when light
shines on the metal.
For a given metal, no electrons were emitted if
the light’s frequency was below a certain
minimum, no matter how long the light was
shone. If light was a wave, it would be able to
knock loose an electron from the metal. We
couldn’t explain why there must be a
minimum frequency in order for this effect to
take place.
Planck’s Explanation
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German physicist Max
Planck suggested that a
hot object emits energy in
small, specific amounts
called quanta (plural for
quantum).
A quantum is the
minimum quantity of
energy that can be lost or
gained by an atom.
The Particle Description of Light
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Planck proposed the following
relationship between a quantum of
energy and the frequency of radiation.
Where E = energy (in Joules), v =
frequency of the radiation emitted, and
h is Planck’s constant, which is
always 6.626 x 10-34 J·s.
The Particle Description of Light
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Einstein expanded this theory by saying
that electromagnetic radiation has a
dual wave-particle nature.
Light exhibits many wavelike features,
but also acts like a stream of particles.
Each particle, or “photon,” of light
carries a quantum of energy.
A photon is a particle of
electromagnetic radiation having 0 mass
and carrying a quantum of energy.
Einstein’s Photons
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According to this equation, the
minimum energy for an electron to be
ejected from a metal surface depends
on the minimum frequency.
The Hydrogen-Atom Line-Emission Spectrum
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When current is passed through a gas
at low pressure, the potential energy of
some of the gas atoms increases.
The lowest energy state of an atom is
its ground state.
A state in which an atom has a higher
potential energy than it has in its
ground state is an excited state.
When an excited atom returns to
ground state, it gives off light.
What a Bright Idea!
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Excited neon
atoms produce
light when
they fall back
to a lower
energy excited
state or the
ground state.
Colors of Neon Signs
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When
electric
current
passes
through
noble
gases,
they emit
different
colors.
Helium
Pink
Neon
Orange-Red
Argon
Lavender
Krypton
Xenon
White
Blue
Line-Emission Spectrum
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When a narrow beam of the emitted
light was shined through a prism, it was
separated into a series of specific
frequencies (and therefore specific
wavelengths, λ = c/v) of visible light.
The line-emission spectrum of an
element is a series of wavelengths of
emitted light created when the visible
portion of light from excited atoms is
shined through a prism.
Line Spectra for Hydrogen, Mercury, and Neon
H
Hg
Ne
Hydrogen’s Line-Emission Spectrum
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Scientists expected to observe a
continuous spectrum when they
observed hydrogen through the
spectrum, but they only saw a few
select colors.
A continuous spectrum is the
emission of a continuous range of
frequencies of electromagnetic
radiation.
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An example of this is a rainbow.
The Perplexing Electron of Hydrogen
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Scientists were confused at why
hydrogen only emits specific frequencies
when current passes through it.
This suggested that the electron of a
hydrogen atom exists only in very
specific energy states.
If an atom at excited state (E2) falls
back to ground state (E1), it releases a
photon that has energy:
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E2-E1=Ephoton= hv
Bohr to the Rescue!
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Danish physicist
Niels Bohr
proposed a model
of the hydrogen
atom that linked
the atom’s
electron with
photon emission.
Atoms Aren’t Bohr-ing!
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Bohr’s model of the atom says:
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The electron can circle the nucleus only in
allowed paths, or orbits.
When the electron is in one of these orbits,
the atom has a definite, fixed energy.
The electron, and the atom as well, is in its
lowest energy state when the electron is in
the orbit closest to the nucleus.
The orbit is separated from the nucleus by a
large empty space where the electron can
not exist.
Here’s a Real-World Example
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Let’s say I was climbing
a ladder…
 On the first rung of
that ladder, I don’t
have much potential
energy. It won’t hurt
if I fall off.
 On the 20th rung of the
ladder, I do have a lot
of potential energy.
Welcome to brokenneck city!
And That’s Not Where It Ends!
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And just as I can
not stand where
there is not a
rung, an electron
can not occupy
where there is
not an orbit.
How Does This Explain the Lines?
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While an electron is in an
orbit, it can neither lose or
gain energy, but it can move
to a higher energy orbit by
gaining an amount of energy
equal to the difference in
energy between the higherenergy orbit and the initial
lower-energy orbit.
When hydrogen is in an
excited state, its electron is
in a higher-energy orbit.
How Does This Explain the Lines?
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When the atom falls back from the
excited state, the electron drops down
to a lower-energy orbit.
In the process, a photon is emitted that
has an energy equal to the energy
difference between E2-E1.
Photons are absorbed when the electron
moves to a higher energy orbit.
Photons are emitted when electrons
move to a lower-energy orbit.