BOHR 2.0 – THE ELCTRON CLOUD
RULES OF THE ROAD –PART 1
▪Aufbau Principle: sublevels must be filled in order from the lowest to the highest energy levels and you may not move on to
the next energy level unless the previous energy level is full. Use the Periodic Table as a guide (read left to right):
1s
2s
2p
3s
3p
4s
3d
4p
5s
4d
5p
5d
6p
Octet Rule: Need to remember no more than 8
valence electrons. 2 is okay for 1st energy
level
s–sublevel = 2 e–
p–sublevel = 6 e–
s– + p–sublevel = 2 e– + 6 e– = 8 e–
• s– and p–sublevels e– do occupy the energy
levels represent by the period #.
• s– and p–sublevels e– compose the valence
7s
5f
6d
7p
shell.
• d– and f–sublevels cannot match the period #, variable n, as they would make the valence energy level have more than 8
e–. d-sublevels gain e– in energy levels just below the valence. i.e n – 1.
• Again the variable n represents the period #. f-sublevels have the same problem but are gained two energy levels below
the valence, .e n– 2.
6s
4f
RULES OF THE ROAD –PART 2
SUBLEVEL CONFIGURATIONS
Tells the # of
in each sublevel (specific orbitals are not shown). Transcribes energy level, ▪ Block’s and Orbitals: 1 orbital = 2 e- = 1
sublevel, and # of e– in the sublevel as indicated below:
n# (s/p/d/f) # of electrons
i.e. 1 s 2
In essence the P.T. is “read” L→R, top to bottom, until you “read” to the element that is wished to
be described. Follows Aufbau Principle / Octet Rule. Remember d- is being gained in n–1
energy levels, f- to n–2 energy levels. Note: each element box represent the gain of one e–.
e–
Procedure: Read P.T. like a sentence (by energy levels = period #, sublevels = block letter, total ein sublevel = superscript) to "punctuation" (element box) on the P.T.
Examples:
He atom is Atomic #2 so read and transcribe to 2nd box
= 1s2
Ti atom is Atomic #22 so read and transcribe to 22nd box.
= 1s2 2s2 2p6 3s2 3p6 4s2 3d2
ABBREVIATED SUBLEVEL CONFIGURATIONS
▪Pauli Exclusion Principle: when
e– do share an orbital, they must be of different “spins.” One “spinning” clockwise ↑ (+1/2)
Uses a noble gas (the far right column) to represent or
thecounter-clockwise
kernel, or core, e↓– (-1/2).
and just shows the outer
▪ Hund’s Rule: all orbitals are must be occupied (1 e-) before any
energy level of e–- using the same method as sublevel configurations
↑↓ NOT ↑↑
NOT ↓↓
orbital within a sublevel (“block”) may be full (2 e-)
Procedure: Identify your “punctuation” (element box) on the P.T. Write the symbol in brackets ([X])
for the noble gas located at the far right of the preceding horizontal row on the table. i.e. For zinc, ▪ Pauli Exclusion Principle: when e– do share an orbital, they must
be of different “spins.” Otherwise being of like charge they would
move up to the third period and across to Ar. To describe the first 18 e– of a zinc atom, write [Ar].
repulse from one another. Thus on ise “spinning” clockwise ↑
Move back down a row (to the row containing the element you wish to describe and transcribe
(+1/2) and the other counter-clockwise ↓ (-1/2).
that row until you reach your “punctuation” (element box). Note: Cannot use a Noble Gas symbol
↑↓ NOT ↑↑
NOT ↓↓
to represent entire configuration.
So if when filling a orbital diagram, i.e. 2p4
i.e. [Ne] incorrect for Ne correct is [He] 2s2 2p6
Following Hund’s Rule and Pauli Exclusion Principle =
Examples:
ground
state
He abbreviated configuration would be = 1s2 i.e. cannot use [NG] to abbreviate
2p
↑↓ ↑ ↑
each got 1 e– before any got 2 e–
Ti abbreviated configuration would be = [Ar] 4s2 3d2
NOT Hund’s Rule and Pauli Exclusion Principle =
ORBITAL DIAGRAMS
excited state
Tells the # and spin of the e– in each orbital. Transcribes energy level, sublevel, # of e– in each
2p ↑↓ ↑↓
should have been in 3rd orbital
orbital (represented by a ), and e– spin (↑ or ↓) of each e– located in each orbital as indicated
Each box represent an orbital:
below:
empty
↑ occupied
or
↑↓ full
n# (s/p/d/f)
for each orbital in sublevel ↑or↓ for each e–
i.e. 1s2 is now 1 s ↑↓
ELECTRON CONFIGURATIONS – THE “CUTE” WAY
Procedure: Read P.T. like a sentence (by energy levels = period #, sublevels = block letter), draw In our analogy elements are universities, streets are energy levels, houses
a box for each orbital in each sublevel (#e-In a sublevel ÷ 2 = # of orbitals). Then fill boxes with are sublevels, rooms are orbitals, and students are e–. All living under
arrows representing e- following Hund’s Rule and Pauli Exclusion Principle to "punctuation". the dictatorial rules of nature.
Note: You must always draw all orbitals present in a sublevel whether occupied or not. Recall So pretend we are moving students into campus housing. The housing is on
1st, 2nd, 3rd, 4th, 5th and 7th street (the energy levels). There are houses on
#e-In a sublevel ÷ 2 = # of orbitals, as illustrated below:
Examples:
Draw the orbital diagram for He
Draw the orbital diagram for Ti
1 s ↑↓
1 s ↑↓
3s ↑↓
4s ↑↓
2s ↑↓
2p ↑↓ ↑↓ ↑↓
3p ↑↓ ↑↓ ↑↓
3d ↑ ↑
these streets.
The houses are called s, p, d and f houses and each must be fully occupied
before any larger house can be begun to be occupied on a street. The s
house has 1 bedroom, the p house has 3 bedrooms, and d house has 5
bedrooms, and the f house has 7 bedrooms.
In each bedroom there is a bunk bed, so two students can sleep in a
bedroom but according to campus policy each room in a house must be
occupied before a second student may be moved into a room.
Also campus
policyPrinciple:
states that when
each room
occupied
a male
▪Pauli
Exclusion
e– domust
sharebean
orbital,bythey
must(↑)
first then
a female
(↓),“spins.”
two males
and/or two
females ↑(↓↓)
are or
not
be of
different
One(↑↑)
“spinning”
clockwise
(+1/2)
allowed
to live together. ↓ (-1/2).
counter-clockwise
↑↓ NOT
↑↑
NOT ↓↓
PHOTOELECTRIC EFFECT
1) (charged –1 each, with a mass of 1/1836 amu each) surround the nucleus
of the atom in distinct energy levels. e– occupy the lowest possible energy
levels when the atom is in the ground state.
2) When e– are given enough energy (in the form of light, heat or electricity), e–
will rise in energy level by the same amount of energy that the e– were
given. The more energy e– absorb, the higher they rise. This is raised state is
called the excited state. This is in accordance with the Law of
Conservation of Energy, which states that energy cannot be created or
destroyed by physical or chemical change.
3) Since e– are negatively charged, and therefore attracted to the positively
charged nucleus, they will eventually fall back to the ground state. As the
electrons fall back to the ground state, they release the energy that caused
them to rise in the first place in the form of photons (light).
4) Photons are the smallest particles known, and are mass-less. They travel at
the fastest theoretical speed possible, c = 3.00 X 108 m/sec, otherwise
known as the speed of light. Photons are, in fact, particles of light.
5) The type/color of the light is determined by the amount of energy lost by the
e– when it dropped back to the ground state. Light particles travel in a wave
pattern. The length of each wave is called, strangely enough, a wavelength.
The more energy a photon has, the shorter its wavelength A photon’s
wavelength is used to classify itself within the electromagnetic spectrum;
shown below. Photons that make up a very small part of the
electromagnetic spectrum, known as visible light, are what we see as
color. From low energy to high energy, the colors of visible light are red,
orange, yellow, green, blue, indigo, and violet (ROY-G-BIV)
6) There are three properties of light waves: energy (E, measured in joules),
wavelength (λ, measured in meters) and frequency (v, measured in
wavelengths per second, also called Hz or Hertz). Since all photons travel
the same speed (3.00 X 108 m/sec) in a vacuum, photons with shorter
wavelength will pass a given point with greater frequency than photons with
longer wavelengths.
e–
PHOTOELECTRIC EFFECT - Excitations
REMEMBER, in order to transition from high
to low, (i.e. n=1 to n=5) energy of a
specific
wavelength
must
be
ABSORBED. This energy is different for
EVERY ATOM due to the interaction of
e−/e− repulsion and the effectiveness of
Nuclear Charge. NO LIGHT EMMISSION
OCCURS.
SPECTROSCOPY
Each element has a unique emission spectrum and absorption spectra because each
element has a different number of e– and each element has its own arrangement
of these e– in different energy levels. As a result, the different e– transitions give
rise to different emission spectra. This is according to Bohr's explanation of the
atom that each element has a unique emission spectrum and absorption spectra:
e–in the ground state have energy of fixed
values i.e. the energy of the e– in a
particular energy level is quantized. If
energy is provided to an atom in its
ground state, a specific amount of this
energy is absorbed and the e– jumps to a
higher energy level i.e. an excited state.
This excited state is unstable and the e–
will fall back to a lower energy level.
As e– fall back, the excess energy is released in the form of light of a definite amount of
energy. The frequency of the light depends on the difference in energy levels and
is given by the equation E2 - E1 = hv (h is Plancks constant, v is the frequency)
A line spectrum can then be observed both as emission and absorption spectrum.
PHOTOELECTRIC
EFFECT - Emissions
REMEMBER, in order to
transition from low to
high, (i.e. n=5 to n15)
energy of a specific
wavelength must be
EMITTED. This energy
is different for EVERY
ATOM due to the
interaction of e−/e−
repulsion
and
the
effectiveness of Nuclear
Charge. NO LIGHT
ABSORBTION
OCCURS.