Transcript electrons - Solon City Schools

Quantum Theory & Electron clouds
The Great Niels Bohr
(1885 - 1962)
The Hydrogen Atom and Quantum Theory
Niels Bohr (Danish) tried to explain the spectrum of
hydrogen atoms.
•Energy is transferred in photon units (quanta),
therefore specific amounts of energy are absorbed
or emitted
•Because the energy of an electron is quantized
(discreet), there are only certain energy levels
(orbits) for electrons
•Therefore, this e-m radiation can behave as waves
or particles = Wave-Particle Duality Theory
•When an electron gains a certain amount of energy
(absorbs a certain number of photons) it becomes
excited and moves to a higher energy level
Spectroscopic analysis of the visible spectrum…
…produces all of the colors in a continuous spectrum
Spectroscopic analysis of the hydrogen
spectrum…
…produces a “bright line” spectrum
Electron transitions
involve jumps of
definite amounts of
energy.
This produces bands
of light with definite
wavelengths.
Bohr Model Energy Levels
Schrodinger Wave Equation

d

V 
8  m dx
h
2
2
2
2
 E
Equation for probability of a
single electron being found
along a single axis (x-axis)
Erwin Schrodinger
Heisenberg Uncertainty Principle
“One cannot simultaneously
determine both the position
and momentum of an electron.”
You can find out where the
electron is, but not where it
is going.
Werner
Heisenberg
OR…
You can find out where the
electron is going, but not
where it is!
•Heisenberg Uncertainty Principle: shows that we
can only predict or estimate the position and momentum of
the electron
•Because we record position by measuring radiant energy
from a particle, we can never know the exact position and
the exact momentum of an electron!!!
Electrons are so small and so easily affected by energy (of
any form) that "lighting up" an electron to see it causes it to
change momentum (position and/or direction)
The de Broglie Hypothesis
Basic Premise of Quantum Mechanics = small particles
in small regions of space (e-) do NOT behave like large
visible objects.
•Therefore… we cannot tell exactly where an e- is at
any given time or how it got there.
Location
__Probability__
•We CAN predict the position of the e
Hypothetical example for comparison:
If I want to find you on a Saturday
afternoon…
•Probable places for finding
electrons = ORBITALS
Friend's house
Shopping
McDonald's
Home in bed
Cruising
Work
Working in Chem Lab
15%
10%
15%
30%
15%
15%
0%
QUANTUM THEORY: shows how the electron
determines an atom's behavior and properties
•4 Quantum Numbers must be used to
describe the position of the electrons in
an atom
•n, l, m, s are the "numbers"
Each electron has a different set (just
like each locker out in the hall has a
different combination)
•Principle Quantum Number -
n
•Identifies the major energy levels that electrons
can occupy
•Shows the distance from the nucleus
•Numbered 1,2,3,4,5,6,7,8 (any integer from 1-∞
•Equation: 2n2 - shows how many electrons can be in each
energy level (e.g. level 3: 2(3)2 = 18 e-'s)
•Angular momentum Quantum Number - l
•Identifies the shape of the sublevels of the main
energy levels
•s, p, d, f - used to identify the shape = "sphere,
peanut, double peanut, flower"
•1st energy level has 1 subshell (s)
•2nd energy level has 2 subshells (s,p)
•3rd energy level has 3 subshells (s,p,d)
•4th energy level has 4 subshells (s,p,d,f)
•At higher energies, these orbitals overlap
•l can be from 0 to n-1
•s=0
•p=1
•d=2
•f=3
Shapes of the Charge Clouds
s for "Sphere": simplest shape, or shape of the
simplest atoms like hydrogen and helium
•Electrons don't interfere with, or block, each other
from the pull of the nucleus - ball shape
•Each energy level has an "s" orbital at the lowest
energy within that level
Shapes of the Charge Clouds
•p for "Peanut": more complex shape
that occurs at energy levels 2 and
above
Shapes of the Charge Clouds
d for "Double Peanut": complex shape
occurring at energy levels 3 and above
The arrangement of these orbitals allows for
"s" and "p" orbitals to fit closer to the
middle/nucleus
Shapes of the Charge Clouds
f for "Flower": 7 bizarre-shaped orbitals for
electrons of very large atoms
•electrons filling these orbitals are weakly
attached to the atom because they are so far
away from the pull of the nucleus
Magnetic Quantum Number
m
•Is indicated by the orientation of these orbitals in
each dimension.
•"p" orbitals line up on the x, y, z, axes in space
•"d" and "f" orbitals can line up on the axes as well as
in between them
•Equation: there are n2 orbitals (of various shapes)
per energy level
•e.g. level 3 has 32 = 9 orbitals (one s, 3 p's, 5 d's)
•m can be from –l to +l
•s subshell can have only one orientation (orbital)
•m can be 0 only, meaning 1 possible orientation in space
•p subshell can have 3 different orientations (orbitals)
•m can be -1, 0, +1, meaning 3 possible orientations in space
•d subshell can have 5 different orientations (orbitals)
•m can be -2, -1, 0, +1, +2 meaning 5 possible orientations in
space
•f subshell can have 7 different orientations (orbitals)
•m can be -3, -2, -1, 0, +1, +2, +3 meaning 7 possible
orientations in space
•Spin Quantum Number (s): indicates the
clockwise or counterclockwise spin of the
electron
•Designated by -1/2 or +1/2
•Needed because NO 2 ELECTRONS CAN
HAVE THE SAME SET OF QUANTUM
NUMBERS!
•Sometimes called left-handed or righthanded spin
Pauli Exclusion Principle
•No two electrons in an atom have the same
set of 4 quantum numbers!
•Therefore, only 2 electrons can fit in any one
orbital
•This works because spinning electrons act like
tiny electromagnets and magnetically attract
each other when they have opposite spin
Aufbau Principle
•Electrons fill the lowest energy levels first
(always)
•As principle quantum number increases, spacing
between the shells decreases
•Therefore, the third subshell and beyond start to
overlap energy levels
Hund's Rule = electrons entering a
subshell containing more than one
orbital will spread out over the
available orbitals with their spins in
the same direction until all orbitals
have one electron in them
Order of Fill for sublevels
•Three ways to describe the
electron structure of atoms in
the ground state: orbital notation,
electron configuration notation,
electron dot notation
20%
20%
20%
5
20%
3
This is how many orbitals the
p-subshell has.
20%
1. 1
2. 2
3. 3
4. 5
7
2
1
5. 7
How many electrons can be in the 3rd
level (n=3)
1. 2
20%
20%
20%
20%
20%
2. 3
3. 10
4. 18
36
18
10
3
2
5. 36
This quantum number indicates the
shape of the orbital.
1. Angular momentum quantum
number
20% 20% 20% 20% 20%
2. Magnetic quantum number
3. Principle quantum number
nt
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5. Spin quantum number
nt
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4. Shape quantum number
This is subshell shown below.
1. n
2. m
17%
17%
17%
17%
17%
d
4. p
p
3. s
17%
5. d
f
s
m
n
6. f
This says that electrons entering a subshell will
spread out over the available orbitals
1. Aufau principle
20% 20% 20% 20% 20%
2. Hund’s rule
3. Heisenberg
uncertainty principle
4. Pauli exclusion
principle
5. De Broglie
hypothesis
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