Transcript atomic emission spectrum

Particle Nature of Light
• New Experiments showed
that the wave model of light
needed to be revised.
• The wave model of light cannot explain
why heated objects emit only certain
frequencies of light at a given
temperature
• Or why some metals emit electrons
when colored light of a specific
frequency shines on them.
Quantum Concept

Max Planck (1900)
 Observed
objects
- emission of light from hot
 Concluded
- energy is
emitted in small, specific
amounts (quanta)
 Quantum
- minimum amount of
energy that can be gained or lost by an
atom
Quantum
Concept
 Planck (1900)
vs.
Classical Theory
Quantum Theory
Because these steps are so small, the
energy/temperature seems to rise in a
continuous, rather than a stepwise, manner.
The quantum concept
Remember: Energy is related to
frequency (and also wavelength!)
• Planck’s constant has a value of
6.626 x 10–34 J · s, where J is the symbol
for the joule, the SI unit of energy.
• Looking at the equation, you can see that
the energy of radiation increases as the
radiation’s frequency, v (of “f”),
increases.
The Quantum Concept
• According to Planck’s theory, for a
given frequency, n, (or f) matter can
emit or absorb energy only in wholenumber multiples of hn; that is, 1hn,
2hn, 3hn, and so on.
• Matter can have only certain amounts
of energy—quantities of energy
between these values do not exist.
Light as a Particle

Einstein (1905)
 Observed
- photoelectric effect
The Photoelectric Effect
• In the photoelectric effect, electrons,
called photoelectrons, are emitted from a
metal’s surface when light of a certain
frequency shines on the surface.
Light as a Particle

Einstein (1905)
 Concluded
- light has properties of
both waves and particles
“wave-particle duality”
 Photon
- particle of light that
carries a quantum of energy
The Photoelectric Effect
• Further, Einstein proposed that the
energy of a photon of light must have a
certain minimum, or threshold, value to
cause the ejection of a photoelectron.
• According to this theory, even small
numbers of photons with energy above
the threshold value will cause the
photoelectric effect.
Bohr’s Model of the Atom

A combination of Plank’s theory and
Einstein’s theories helped Bohr develop
and explain his Energy Level Model.
•Bohr studied the
Atomic Emission
Spectrum of the
Hydrogen Atom to
calculate the energy
of each of the energy
levels in the atom.
Atomic Emission Spectra of Hydrogen
•Hydrogen’s atomic emission spectrum
consists of several individual lines of color,
not a continuous range of colors (like a
rainbow) as seen in the visible spectrum.
•The fact that only certain colors appear in
hydrogen’s atomic emission spectrum means
that only certain specific frequencies of light
are emitted.
Atomic Emission Spectra
• The atomic emission spectrum of an
element is the set of frequencies of the
electromagnetic waves emitted by atoms
of the element.
•An atomic emission spectrum is characteristic
of the element being examined and can be
used to identify that element.
Bohr Model


Bohr proposed that electrons exist only in
orbits with specific amounts of energy
called energy levels
Therefore…
 e-
can only gain or lose certain
amounts of energy
 only
certain photons are produced
 These
photons have different colors
because they have different energy
values
Line-Emission Spectrum
•Bohr
said the different colors of light emitted were
caused by an electron that had been excited away
from their ground state near the nucleus, returning
to its ground state.
excited state
ENERGY IN
PHOTON OUT
ground state



The lowest allowable energy state of
an electron in an atom is called its
ground state.
Electrons must absorb a specific
“quantum of energy” to be excited to
a higher energy level. This is said to
be it’s “excited state”
When the excited electron drops back
down to its “ground state” it emits a
photon of light with the same energy
as the difference between the energy
levels it drops.
Energy Level Model



Bohr measured the wavelengths and
frequencies of the light emitted and
looked for patterns.
He came up with an equation to
determine the energy of the light that will
be emitted or absorbed when an electron
changes energy levels.
This is how he calculated the energy of
each of the Energy Levels in his model!!!
Bohr Energy Level Model
65
4
 Energy
3
2
1
of photon
depends on the
difference in energy
levels
 Bohr’s calculated
energies matched
the IR, visible, and
UV lines for the H
atom
•The smaller the electron’s orbit, the lower the
atom’s energy state, or energy level.
•Conversely, the larger the electron’s orbit, the
higher the atom’s energy state, or energy level.
An explanation of hydrogen’s line spectrum
• The four electron
transitions that
account for visible
lines in hydrogen’s
atomic emission
spectrum are shown.
Other Elements

Each element has a unique bright-line
emission spectrum.

“Atomic Fingerprint”
Helium
Bohr’s calculations only worked for
hydrogen! 
Energy states of hydrogen
• And although a hydrogen atom contains only a
single electron, it is capable of having many
different excited states.