Energy Levels and Orbitals

1
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.
• Write electron configurations for elements or
  ions (incl. noble gas config.)
• Draw orbital energy diagrams for elements or ions.

2
Emission Spectroscopy

• The spectra that were shown through emission
  spectroscopy led Niels Bohr to question the
  structure of the atom.

3
Electromagnetic Spectrum

• With white light, all of the colors of the
  visible spectrum are shown.

4
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.

5
Wavelengths and Energy

• E hc
• ?
• Energy hc two constants
  wavelength
• (Plancks and speed of light)
• 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.

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The Bohr Model

• According to Niels Bohr, an electron can circle
  the nucleus in orbits of only certain distances
  from the nucleus. Bohr called these orbits, or
  energy levels.
• An electron cannot be in-between energy levels
  (i.e. it is either on the first level or the
  second).
• Therefore, energy is quantized.

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Exciting electrons...

• Niels 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 (more on that in a
  moment)

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The Bohr Model (excitation)
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.
4th Energy Level
1st Energy Level
Energy Coming In!
Nucleus
3rd Energy Level
2nd Energy Level
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The Bohr Model (emission)
The Bohr Model When the electron releases the
energy, the electron returns to lower energy
levels. Other forms of electromagnetic radiation,
besides visible light, can be emitted.
4th Energy Level
1st Energy Level
Energy emitted (ultraviolet light)
Nucleus
Energy emitted (red light)
Energy emitted (infrared)
3rd Energy Level
2nd Energy Level
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The Bohr Model (alternate emission)
The Bohr Model When the electron returns to its
ground state, it has the option of jumping down
multiple energy levels, rather than one at a time.
4th Energy Level
1st Energy Level
Energy emitted (ultraviolet light)
Nucleus
Energy emitted (blue/green light)
2nd Energy Level
3rd Energy Level
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Types of Radiation

• The following are types of electromagnetic
  radiation, listed from highest energy to the
  lowest
• Gamma rays cosmic radiation,
• very high energy
• Ultraviolet rays (UV) solar radiation,
• high energy
• Infrared rays (IR) thermal radiation, remote
  controls, low energy
• Visible Light (more to follow)
• Microwave rays microwave oven, very low energy
• Radio lowest energy waves

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Types of Radiation

• 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

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Energy Levels and Spectra
Electrons release certain types of
electromagnetic radiation as they fall to
specific energy levels
Energy Level Change
Spectra Emission

• 2 --gt 1 Ultraviolet
• 3 --gt 1 Ultraviolet
• 4 --gt 1 Ultraviolet
• 3 --gt 2 Visible Red
• 4 --gt 2 Visible Blue/Green
• 5 --gt 2 Visible Blue
• 4 --gt 3 Infrared

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Quantum Theory

• Energy emission and absorption from elements like
  hydrogen led to scientists attempting to explain
  why

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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

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Using the Quantum Mechanical Model

• Quantum mechanics is a mathematical way of
  describing where electrons are located.
• It is based on the probability of finding an
  electron in the space outside the nucleus.

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Why Quantum Numbers?

• The quantum numbers are like an address
• State
• City
• Street
• House Number
• Each piece of information is needed to describe
  the location, and each one tells more specific
  information about where the electron is located.

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First Quantum NumberEnergy level (n)

• Each energy level is farther away from the
  nucleus.
• Electrons are attracted to the nucleus, so they
  will fill the lower energy levels first!

E5
E4
E3
E2
E1
nucleus
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Second Quantum Number Subshell (l)
As the energy levels increase, so do the number
of subshells that are needed to cover all the
space around the atom. The first energy level
(n1) has 1 subshell (s) The second energy level
(n2) has 2 subshells (s p) The third energy
level (n3) has 3 subshells (s, p, d) The
fourth energy level (n4) has 4 subshells (s, p,
d, f)
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Extension

• How many subshells would be present in energy
  level 5?
• Answer 5!
• s, p, d, f, and g
• How many subshells would be present in energy
  level 6?
• Answer 6!
• s, p, d, f, g, and h

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Subshells

• s orbital sphere
• p orbital peanut
• d orbital double peanut
• f orbital flower

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Quantum Mechanical Model

• To recap
• Energy level 1 1 subshell (s)
• Energy level 2 2 subshells (s and p)
• Energy level 3 3 subshells (s, p, and d)
• Energy level 4 4 subshells (s, p, d, and f)
• etc.
• Why are more subshells present?
• Each energy level is larger than the previous. As
  a result, there are more possible locations for
  where an electron could reside.

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Nucleus
1s subshell
2s subshell
2p subshell
3s subshell
3p subshell
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3d subshell
4s subshell
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Third Quantum NumberAtomic Orbitals ( ml )

• The atomic orbital essentially describes how many
  of that shape of subshell are needed to cover all
  the space around the nucleus.
• The more complicated the shape, the more orbitals
  are needed to cover all the space.

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Third Quantum NumberAtomic Orbitals ( ml )

• s has 1 orbital (just 1 type of s)

• p has 3 orbitals (px, py, pz)

• d has 5 orbitals (dxy, dxz, dyz, dz2, dx2-y2)

• f has 7 orbitals (etc., etc.,)

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There is 1 s orbital
There are 3 p orbitals
There are 5 d orbitals
There are 7 f orbitals
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Fourth Quantum Number Electron Spin ( ms )
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.
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Energy Level Possible Subshells Atomic Orbitals Number of Electrons in Each Subshell Maximum Possible Electrons in Energy Level
1 s 1 2 2
2 s p 1 3 2 6 8
3 s p d 1 3 5 2 6 10 18
4 s p d f 1 3 5 7 2 6 10 14 32
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Aufbau Principle / Hunds Rule
Aufbau Fill from the ground up Hunds
Rule When choosing between equivalent orbitals,
fill the empty orbitals first