10/28/2016

Gold Foil Experiment
◦ Small, Dense and Positive Nucleus
◦ Electrons in space
◦ Suggested electrons dispersed around nucleus
Honors Chemistry
Dr. Kevin Moore

Energy in the form of a wave
◦ Peaks and Troughs
◦ Where do we see waves?

BUT HOW?
Frequency (ν) - # of wave peaks that pass a point
per unit time (typically seconds)
Wavelength (λ) – distance between two peaks on a
wave
λ
# per second = ν

Particle?
◦ Photon – packet of energy
Speed of Light (c) – 3.00 x 108 m/sec
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1600

1400
Directly Proportional to the Frequency
Wavelength(nm)
1200
1000
Planck’s Constant: h = 6.626 x 10-34 Jsec
800
600

400
200
What is the energy of yellow light having a
wavelength of 625 nm?
0
0
1E+15
2E+15
3E+15
⁄
.
Frequency
6.626
Electromagnetic Energy (Radiation)
◦
◦
◦
◦

X-Rays
Visible Light
Infrared
Microwaves
·
4.80
10
Atoms absorb energy
◦ Released at some random point later
◦ Specific frequencies of light released
3.18
10
We see Light
Energy Released
Excited State
Infrared
Visible
UV
X-Ray
Gamma
380-780 nm
Radio/TV
Increasing λ
Microwaves
Increasing ν & E
1 mm
1m
Energy

10
=4.80 x1014 sec-1
New Excited State
Energy
Ground State
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White Light is composed of all frequencies
Excited atoms give off discreet frequencies of
light
◦
◦
◦
◦
◦
◦

Sodium – Yellow
Hydrogen – Blue
Neon – Red
Mercury – Blue
Lithium – Red
Potassium - Purple

1 electron system
n=4
◦ Available energy levels
 Quantized (Discreet)
◦ Principle Energy Levels (n)
 Correspond to the distance
from the nucleus to the electron

Balmer Series
Excitation

Energy

◦ Return to 2nd Energy Level
Atomic Line Spectra
n=2
◦ Lines of specific frequencies of light given off by an
atom

Explained the Hydrogen atom
◦ Used a solar system model
◦ Energy levels represented as planetary orbits
n=3
n=1

Mathematically predicted the entire spectrum
of Hydrogen
◦ Failed for all other atoms
◦ Unable to handle correlation between electrons
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
Louis De Broglie

◦ Suggested that the electron could be treated as a
wave

◦ n – Principal Quantum #
 Region of space in which a group of electrons will be
found
Erwin Schroedinger
◦ l – Angular Momentum Quantum #
◦ Used the wave mechanical model to predict the
structure of the atom

 Sublevel in which a group of electrons will be found
◦ ml – Magnetic Quantum #
Werner Heisenberg
 Orientation of individual orbitals in a sublevel which
may contain a pair of electrons
◦ Showed that you cannot determine the velocity and
the position of the electron
◦ Heisenberg Uncertainty Principle

n = Integer Values
l (range of possible values)

ml (range of possible values)

◦ 0…n-1
◦ -l … 0 … +l

ms (spin states of electron)
◦ +1/2 (spin up)
◦ -1/2 (spin down)
Uniquely define the location of every electron

Each electron in an atom must have a unique
set
n=1
l=0
ml=0
ms=+½, –½
2 possible electrons
n=2
l=0
ml=0
ms=+½, –½
2 possible electrons
n=2
l=1
ml=-1, 0, 1
ms=+½, –½
6 possible electrons
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
l=0

◦ s sublevel

l=1

l=2
◦ Energy increases with complexity
 s<p<d<f
◦ Contained in subshells (subLevels)
◦ p sublevel



s, p, d & f orbitals

Shells (Principal Energy Levels)
◦ d sublevel
◦ Distance from nucleus
◦ f sublevel
◦ Row on Periodic Table
◦ Each shell contains the same # of sublevels as its #
 n=1, 2, …
l=3
Pictures are the most probable
location of finding electron
Electron density maps are called
orbitals
◦
◦
◦
◦
s – simplest
p
d
f


Electrons occupy orbitals of equal energy with
parallel spins
Pauli Exclusion Principle
◦ An atomic orbital can hold no more than two
electrons with opposite spins
 2 e- in each orbital
 Opposite spins

Aufbau Principle
◦ Electrons fill orbitals which have the lowest energy
state

Helium
◦ 1s2

↿⇂
1s
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
Shell (n)
1
2
2
3
3

Subshell (l)
s
s
p
s
Orbital (ml)
s
s
px py pz
s
Sublevels
◦
◦
◦
◦
d
px py pz dxy dxz dyz dx2-y2 dz2

s: 1 orbital
p: 3 orbitals
d: 5 orbitals
f: 7 orbitals
Orbitals
◦ Each contain a maximum of 2 e-
How do electrons decide which shells/subshells
to occupy?
◦
◦
◦
◦
ENERGY!!!
Lowest energy is favorable state (ground state)
Periodic Table
Energy Level (shell) = Row on Periodic Table
 1st shell  1 sublevel (s)  2 electrons
 2nd shell  2 sublevel (s, p)  2 + 6 electrons
 3rd shell  3 sublevel (s, p, d)  2 + 6 + 10 electrons
4d
5s
Energy

p
Each shell has additional sublevel
◦ Important sublevels (s, p, d, f)
3
4p
3d
4s
3p
3s
2p
2s
1s
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
◦ s of a new shell is always filled before d of the previous
shell
1s
2s
2p
3s
3p
3d
4s
4p
4d
4f
5s
5p
5d
5f
6s
6p
6d
7s
7p






H: 1s1
He: 1s2
Li: 1s22s1
Be: 1s22s2
B: 1s22s22p1
O
◦

1s22s22p4
Ne
◦ 1s22s22p6
Competing energies
◦ 1<2<3<4<
◦ s<p<d<f
◦ 4s fills before the 3d

# of electrons in each sublevel
◦
◦
◦
◦
8s

Each new Energy Level begins with an s sublevel

s – 1 orbital : 2 electrons
p – 3 orbitals: 6 electrons
d – 5 orbitals: 10 electrons
f – 7 orbitals: 14 electrons
Oxygen : 1s22s22p4
Electron
n
l
ml
1 (1s)
1
0
0
+1/2
ms
2 (1s)
1
0
0
-1/2
3 (2s)
2
0
0
+1/2
4 (2s)
2
0
0
-1/2
5 (2p)
2
1
-1
+1/2
6 (2p)
2
1
0
+1/2
7 (2p)
2
1
1
+1/2
8 (2p)
2
1
-1
-1/2
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

Magnesium : 1s22s22p63s2
Electron
n
l
ml
1 (1s)
1
0
0
+1/2
ms
2 (1s)
1
0
0
-1/2
3 (2s)
2
0
0
+1/2
4 (2s)
2
0
0
-1/2
5 (2p)
2
1
-1
+1/2
6 (2p)
2
1
0
+1/2
7 (2p)
2
1
1
+1/2
8 (2p)
2
1
-1
-1/2
9 (2p)
2
1
0
-1/2
10 (2p)
2
1
1
-1/2
11 (3s)
3
0
0
+1/2
12 (3s)
3
0
0
-1/2
Noble Gas contains a core set of electrons

◦ Completed sublevels (shells & subshells)
◦ All Atoms beyond include the core set of electrons




Na: [Ne]3s1
Mg: [Ne]3s2
K: [Ar]4s1
Each row begins in the s sublevel
From 4th Row, all atoms have the d-block
◦ Previous Shell
◦ 10 electrons (max)

Sc: 1s22s22p63s23p64s23d1
◦ [Ar]4s23d1

Fe: 1s22s22p63s23p64s23d6
◦ [Ar]4s23d6

As:1s22s22p63s23p64s23d104p3
◦ [Ar]4s23d104p3
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
6th Row

◦ 4f fills after the 6s
Electrons in outermost shell
◦ Chemical Reactivity
◦ Typically beyond Previous Noble Gas
◦ Show in orbital diagrams
 In most cases, a single e- fills the 5d first

Oxygen: [He]2s22p4
↿⇂
2s



Potassium: [Ar]4s1
Iron (Fe)
↿j
Selenium
4s
4s
◦ [Ar]4s23d6

↿⇂
◦ [Ar]4s23d104p4
↿⇂
4s








s & d subshells have very small energy
difference
Cr: [Ar]4s23d4
↿⇂ ↿ ↿ ↿ ↿ jj
4s
3d
1
5
Cr: [Ar]4s 3d
↿
↿ ↿ ↿ ↿ ↿ jj
4s
3d
1
10
Cu:[Ar]4s 3d
↿
↿⇂ ↿⇂ ↿⇂ ↿⇂ ↿⇂jj
4s
3d

↿⇂ ↿ ↿ j
2p
↿⇂ ↿ ↿ ↿ ↿ j
3d
↿⇂ ↿⇂ ↿⇂ ↿⇂ ↿⇂
3d
↿⇂ ↿ ↿ j
4p
Electron Configuration
◦ Electrons are always removed from the outer shell
 even if inner shell filled last
 outer shell s electrons are removed before d electrons

Fe: [Ar]4s23d6

 Fe+2:
[Ar]3d6


Fe+3: [Ar]3d5

↿⇂ ↿⇂ ↿ ↿ ↿ ↿ jj
4s
4s
4s
3d
↿⇂ ↿ ↿ ↿ ↿ jj
3d
↿ ↿ ↿ ↿ ↿ j
3djj
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
p subshell typically gains electrons
◦ O: [He]2s22p4

◦ O-2: [He]2s22p6

◦ C: [He]2s22p2

◦
C+2:

[He]2s2
↿⇂
2s
↿⇂ ↿
2p

↿ j
◦ Relax back to a lower energy state
◦ Releases Energy in the form of light
↿⇂ ↿⇂ ↿⇂ ↿⇂ j
2s
2p
2s
2p
2s
2p
↿⇂ ↿
↿⇂
↿
j
Atoms (electrons) absorb energy and move to
an excited state

Atom is more complicated than solar system
Electrons have a specific arrangement in the
atom with specific energy levels

Properties are explained by configuration

j
◦ Quantum Numbers define arrangement
◦ Atomic Radius
◦ Ionization Energy
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