an investigation into electrons and their location and behavior within the atom learning targets: ...
TRANSCRIPT
Energy Levels and Orbitals
An investigation into electrons and their location and behavior within the atom
Learning Targets: Describe the process of excitation and emission
of energy by an electron.
Emission Spectroscopy
The spectra that were shown through emission spectroscopy led Niels Bohr to question the structure of the atom.
Emission Spectroscopy
With white light, all of the colors of the visible spectrum are shown.
Emission Spectroscopy
Since that was NOT what the spectra of elements looked like, Bohr began to look at why only certain wavelengths of color appeared.
Emission Spectroscopy
E = hc λ
Energy h = 6.63 x 10-34 Js wavelength
in metersc = speed of light = 3x108 m/s
This equation shows that larger wavelengths indicate lower amounts of energy and smaller wavelengths indicate higher amounts of energy... an inverse relationship.
Bohr realized that the specific wavelengths revealed specific amounts of energy.
Emission Spectroscopy
Specific amounts of energy!!
That inferred that energy within the atom existed at specific amounts. Bohr called these orbits, or energy levels.
An electron cannot be in-between energy levels, it can only be within an energy level.
Therefore, energy is quantized.
The Bohr Model
Emission Spectroscopy
Bohr realized that the spectra were being created as electrons moved between these energy levels:
If an electron absorbs energy, it may jump to a higher energy level.
When an electron is at a higher energy level we say that the electron is in its “excited” state.
When the electron releases energy in the form of radiation, we say that the electron has returned to its “ground” state.
The type of radiation that is emitted depends on the amount of energy released.
Emission Spectroscopy
Nucleus
1st Energy Level
3rd Energy Level2nd Energy Level
4th Energy Level
Energy Coming In!
The Bohr Model
When energy enters the atom, an electron
(shown in red) can absorb the energy becoming
excited, AND jumping to
higher energy levels.
Emission Spectroscopy
Nucleus
1st Energy Level
3rd Energy Level2nd Energy Level
4th Energy Level
Energy emitted (infrared)
Energy emitted (red light)
Energy emitted (ultraviolet light)
The Bohr ModelWhen the electron
releases the energy, the
electron returns to
lower energy levels. Other
forms of electromagnet
ic radiation, besides visible light, can be emitted.
Emission Spectroscopy
Nucleus
1st Energy Level
3rd Energy Level2nd Energy Level
4th Energy Level
Energy emitted (blue/green light)
Energy emitted (ultraviolet light)
The Bohr ModelWhen the electron
returns to its ground state,
it has the option of
jumping down multiple
energy levels, rather than
one at a time.
Emission Spectroscopy
Nucleus
1st Energy Level
3rd Energy Level2nd Energy Level
4th Energy Level
Energy emitted (blue/green light)
The Bohr ModelSince a
sample of gas has many
atoms, there are many
electrons. This is why Bohr saw multiple
colors.But there were
other electromagnetic waves, too.
Energy emitted (red light)
Emission Spectroscopy
This is the full electromagnetic spectrum.
Emission Spectroscopy
Bohr saw Visible Light: wavelength is in the
range of 400 to 700 nanometers (4 x 10-7 meters)
ROY G. BIV White light is made of all
the colors of light
Electromagnetic Waves
Emission Spectroscopy
Gamma rays: cosmic radiation, very high energy
Ultraviolet rays (UV): solar radiation, high energy
Infrared rays (IR): thermal radiation, remote controls, low energy
Microwave rays: microwave oven, very low energy
Electromagnetic Waves
Emission Spectroscopy
2 --> 1Ultraviolet3 --> 1Ultraviolet4 --> 1Ultraviolet3 --> 2 Visible Red4 --> 2 Visible Blue/Green5 --> 2 Visible Blue4 --> 3 Infrared
Energy LevelChange
Spectra Emission
Electrons release certain types of electromagnetic radiation as they fall to specific lower energy levels.
Quantum Mechanical Model
In addition to knowing that there were energy levels in the atom, three scientists began to notice other things... Heisenberg – impossible to know the exact
position and exact speed of an electron at the same time
De Broglie – electrons have wave-like properties, as in they move in wave patterns
Schroedinger – developed probability of finding each electron in a given location
Quantum Mechanical Model
Heisenberg Bohr suggested that the
electrons move in perfect circles around the nucleus.
Heisenberg showed that, instead, the electron moves in a three dimensional cloud of probability that is smeared out over the orbit – Heisenberg uncertainty principle
Quantum Mechanical Model
DeBroglie Bohr suggested that
the electrons move in perfect circles around the nucleus.
DeBroglie showed that there were other shapes because the electrons moved like waves – wave-particle duality.
Quantum Mechanical Model
Schrodinger Schroedinger realized
how to put the theories of Bohr, Heisenberg, and DeBroglie together by creating a mathematical equation to find the most likely location for each electron within an atom – wave equation.
Watch this YouTube video.
Quantum Mechanical Model
Every electron within an atom has “coordinates”. Schrodinger gave these coordinates numerical values, known as quantum numbers. Each quantum number describes part of the coordinates that determine the energy and probable location of any electron for any atom.
Energy levels begin at the number 1.
Each level is higher in energy than the next.
The higher in energy, the farther away from the nucleus.
Quantum Mechanical Model
First Quantum Number Energy Level
Quantum Mechanical Model
Second Quantum Number Subshell Atoms are three
dimensional. Within the energy
levels exist different shapes, or subshells.
The shapes are determined by how much energy is required to create them.
Quantum Mechanical ModelThird Quantum Number Orbitals
Did you notice that there were different positions of some of the subshells?
The different positions, or orientations, are called orbitals, not orbits.
The orbitals are determined by which subshell they are in and in which positions they are.
The s orbital does NOT have a different position.
The p orbital has THREE different orientations – x,
y, and z.
Each electron can be spin up (+1/2) or spin down (-1/2)
No two electrons in the same orbital orientation can have the same spin.
With only one spin up and one spin down, the maximum number of electrons that can fit into any given orbital orientation is two.
This is called the Pauli Exclusion Principle.
Quantum Mechanical Model
Fourth Quantum Number Electron Spin