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Energy Levels and Orbitals

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Title: 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.

6
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.

7
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)

8
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
9
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
10
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
11
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

12
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

13
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

14
Quantum Theory
  • Energy emission and absorption from elements like
    hydrogen led to scientists attempting to explain
    why

15
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

16
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.

17
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.

18
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
19
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)
20
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

21
Subshells
  • s orbital sphere
  • p orbital peanut
  • d orbital double peanut
  • f orbital flower

22
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.

23
Nucleus
1s subshell
2s subshell
2p subshell
3s subshell
3p subshell
24
3d subshell
4s subshell
25
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.

26
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.,)

27
There is 1 s orbital
There are 3 p orbitals
There are 5 d orbitals
There are 7 f orbitals
28
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.
29
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
30
Aufbau Principle / Hunds Rule
Aufbau Fill from the ground up Hunds
Rule When choosing between equivalent orbitals,
fill the empty orbitals first
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