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ARRANGEMENT OF ELECTRONS IN
ATOMS
Chapter 4
Visible Light
We are all familiar with light but what is “visible” is just a very, very small portion of the electromagnetic spectrum
What colors make up the rainbow?
Red, Orange, Yellow, Green, Blue, Indigo, Violet (ROYGBIV)
The E-M Spectrum
Gamma Rays(Very Harmful / Cancerous)
X-rays(Cancerous in large doses; small doses medical scanning)
Ultraviolet ( not as harmful - sunburn; black lights)
Infrared(Heat, communication)
Microwaves(Cooking, communications)
Radio Rays(TV, Radio, other communications)
The Development of a New Atomic Model
Wavelength () - length of one complete wave
Frequency () - # of waves that pass a point during a certain time period hertz (Hz) = 1/s
Amplitude (A) - distance from the origin to the trough or crest
Waves
Agreater
amplitude(intensity)
greater frequency
(color)
crest
origin
trough
A
The Electromagnetic Spectrum
AM radio
Short waveradio
Television channels
FM radio
RadarMicrowave
Radio Waves Gamma Rays
X- Raysinfrared
Increasing photon energy
Increasing frequency
Decreasing wavelength
Red Orange Yellow Green Blue Indigo Violet
UV Rays
R O Y G B I V
Visible
Light
Electromagnetic Spectrum
Frequency & wavelength are inversely proportional
c = c: speed of light (3.00 108 m/s): wavelength (m, nm, etc.): frequency (Hz)
Electromagnetic Spectrum
GIVEN: = ? = 434 nm = 4.34 10-7 mc = 3.00 108 m/s
WORK: = c
= 3.00 108 m/s 4.34 10-7 m
= 6.91 1014 Hz
EX: Find the frequency of a photon with a wavelength of 434 nm.
So why is the electromagnetic spectrum so important to chemistry?
• Why is the steel emitting light when it is heated?
• We take it for granted that when things get hot they turn red then orange and finally white; but that isn’t good enough any more
Black Body Radiation Colors
Black Body Radiation Colors
So why is the electromagnetic spectrum so important to chemistry?
Incandescence is heat made visible – the process of turning heat energy into light energy.
Our usage of terms like "red hot," "white hot," and so on, is part of the color sequence black, red, orange, yellow, white, and bluish white, seen as an object is heated to successively higher temperatures.
So why is the electromagnetic spectrum so important to chemistry?
The light produced consists of photons emitted when atoms and molecules release part of their thermal vibration energy.
For increasing temperatures, the sequence of radiated colors is: black, red, orange, yellow-white, bluish-white.
Heat and Light
Planck (1900)
Observed - emission of light from hot objects
Concluded - energy is emitted in small, specific amounts (quanta)
Quantum - minimum amount of energy change
Energy and Light
E: energy (J, joules)h: Planck’s constant (6.6262 10-34 J·s): frequency (Hz)
E = h
The energy of a photon is proportional to its frequency.
Energy and Light
GIVEN:E = ? = 4.57 1014 Hzh = 6.6262 10-34 J·s
WORK:E = h
E = (6.6262 10-34 J·s)(4.57 1014 Hz)
E = 3.03 10-19 J
EX: Find the energy of a red photon with a frequency of 4.57 1014 Hz.
Niels Bohr and the Bohr model of the atomBohr hypothesized that instead of haphazardly orbiting the nucleus, electrons had clearly defined orbits – very similar to the planetary orbits circling our sun
His model is (cleverly) named the Planetary Model
Niels Bohr
Bohr Model (1913)
Bohr’s Proof
Bohr said this: If you assume that the electrons have clearly defined orbits that are congruent to the energy levels…
Bohr’s Proof
… then when an electron gets “excited” it jumps to a higher energy level. When it “relaxes” it emits a certain wavelength of light.
• Bohr showed the energy of an electron in an atom is quantized, which means it has a particular numerical value, not some arbitrary number.
Excitation of Hydrogen Atoms
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 328
Return to Ground State
Bohr’s Proof
n=1
n=2
n=6n=5n=4
n=3
n=7
1= 1.097373 x 107 m-1
λ nr2 ne
2
1 1-
Lyman Series (uv)
Balmer Series (vis and uv)
Paschen Series (ir)
Emission Spectrum of an Element
1 nm = 1 x 10-9 m = “a billionth of a meter”
410 nm 434 nm 486 nm 656 nm
1 nm = 1 x 10-9 m = “a billionth of a meter”
Continuous and Line Spectra
Hydrogen to Steel
If Hydrogen emits 4 distinct wavelengths of light when its one electron is excited what can we extrapolate to that of steel which is made mostly of iron?
http://jersey.uoregon.edu/vlab/elements/Elements.html
Flame Emission Spectra
Photographs of flame tests of burning wooden splints soaked in different salts.Photographs of flame tests of burning wooden splints soaked in different salts.
Include link to web page
http://www.unit5.org/christjs/flame%20tests.htm
methane gas wooden splint strontium ioncopper ionsodium ion calcium ion
Fireworks
Composition of Fireworks
Gunpowder Sulfur, charcoal, potassium nitrate (saltpeter)
Salts (to give color) Red = lithium Green = copper
Good News Bad News
Good News
Bohr’s Model works and moves us along in the development of the Atomic Theory
End of this little unit
Bad (Frustrating) News
Lots of Math
Everything I taught you only works for Hydrogen and therefore is completely wrong and obsolete.
Check for Understanding
c= λν E=hν c=3.0 x108 m/s h=6.626 x 10-34 J s
What is the frequency of a radar photon with an energy of 7.2 x 10-24 J?
What is the frequency of light having a wavelength of 6.20x10-7m?
Models of the Atom
Dalton’s model (1803)
Thomson’s plum-pudding model (1897)
Rutherford’s model (1909)
Bohr’s model (1913)
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
Greek model(400 B.C.)
+--
--
-e
e
e
+
+ +
+
++
++
e
ee
e
e
ee
Your Current View of the Atom
electrons
nucleus
Again… so why is it so important to chemistry?
Einstein (1905) Observed - photoelectric effect
Again… so why is it so important to chemistry?
Einstein (1905) Concluded - light has properties of both waves
and particles
“wave-particle duality”
Photon - particle of light that carries a quantum of energy
Quantum Mechanical Model
Modern atomic theory describes the electronic structure of the atom as the probability of finding electrons within certain regions of space (orbitals).
Modern View
The atom is mostly empty space
Two regions Nucleus
protons and neutrons Electron cloud
region where you might find an electron
Also called the electron cloud model
Modern View of Atom
e-e- Ground state
Excited state
Electrons can only be atspecific energy levels,NOT between levels.
Models of the Atom
Dalton’s model (1803)
Thomson’s plum-pudding model (1897)
Rutherford’s model (1909)
Bohr’s model (1913)
Charge-cloud model (present)
Dorin, Demmin, Gabel, Chemistry The Study of Matter , 3rd Edition, 1990, page 125
Greek model(400 B.C.)
+--
--
-e
e
e
+
+ +
+
++
++
e
ee
e
e
ee
C. Johannesson
Electrons as Waves
Louis de Broglie (1924) Applied wave-particle theory to e-
e- exhibit wave properties
QUANTIZED WAVELENGTHS
C. Johannesson
Quantum Mechanics
Heisenberg Uncertainty Principle Impossible to know both the velocity and position of
an electron at the same time
C. Johannesson
Quantum Mechanics
σ3/2 Zπ
11s 0
eΨ a
Schrödinger Wave Equation (1926) finite # of solutions quantized energy levels
defines probability of finding an e-
Quantum Theory
quantum theory- Describes mathematically the wave
properties of electrons and other small particles
orbital- a region of an atom in which there is a high probability of finding electrons
Today’s atomic model predicts quantized, or particular energy levels for electrons.
does not describe the exact path or location electrons take or can be found around the nucleus
concerned with the probability, or likelihood, of finding an electron in a certain position
Two electrons can occupy each orbital, also called an electron cloud.
Quantum Numbers
Four Quantum Numbers: Specify the “address” or “seat” of each
electron in an atom
UPPER LEVEL
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
1. Principal Quantum NumberPrincipal Quantum Number ( nn )
Energy level (ladder rungs)
Size of the orbital
Positive integer
1s
2s
3s
Quantum Numbers
1. Principal Quantum NumberPrincipal Quantum Number > number, further away from the
nucleus 1- right next to the nucleus 3- further away from nucleus
> number, higher the energy level n = 2 greater energy level than n = 1 these electrons have more energy
than electrons in the n = 1 level
1s
2s
3s
Quantum Numbers
2. Angular Momentum Quantum #Angular Momentum Quantum # ( ll ) Energy sublevel (orbital) Shape of the orbital Often represented by letters than numbers
s p d fCourtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
px pz py
x
y
z
x
y
z
x
y
z
d-orbitals
Zumdahl, Zumdahl, DeCoste, World of Chemistry 2002, page 336
Quantum Numbers
Orbitals combine to form a spherical shape.
2s
2pz2py
2px
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Quantum Numbers
3. Magnetic Quantum NumberMagnetic Quantum Number ( mmll ) Orientation of orbital
Specifies the exact orbital within each sublevel
Shapes of s, p, and d-Orbitals
Quantum Numbers
4. Spin Quantum NumberSpin Quantum Number ( ms ) Electron spin +½ or -½
An orbital can hold 2 electrons that spin in opposite directions.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Maximum Capacities of Subshells and Principal Shells
n 1 2 3 4 ...n
l 0 0 1 0 1 2 0 1 2 3
SubshellSubshelldesignationdesignation s s p s p d s p d f
Orbitals inOrbitals insubshell subshell 1 1 3 1 3 5 1 3 5 7
SubshellSubshellcapacity capacity 2 2 6 2 6 10 2 6 10 14
Principal shellPrincipal shellcapacity capacity 2 8 18 32 ...2n2
Hill, Petrucci, General Chemistry An Integrated Approach1999, page 320
Filling Rules for Electron Orbitals
Aufbau Principle: Electrons are added one at a time to the lowest energy orbitals available until all the electrons of the atom have been accounted for.
Pauli Exclusion Principle: An orbital can hold a maximum of two electrons.To occupy the same orbital, two electrons must spin in opposite directions.
Hund’s Rule: Electrons occupy equal-energy orbitals so that a maximum number of unpaired electrons results.
*Aufbau is German for “building up”
Diagonal Rule
1s
2s
3s
4s
5s
6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
6d
4f
5f
General Rules
Pauli Exclusion PrinciplePauli Exclusion Principle Each orbital can hold TWO electrons with opposite
spins.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
RIGHTWRONG
General Rules
Hund’s RuleHund’s Rule Within a sublevel, place one electron per
orbital before pairing them.
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Orbital Diagrams and Electron Configurations
Orbital diagrams Show how electrons are distributed within sublevels Electrons represented by an arrow Orbital is represented by a box Direction of spin represented by direction of arrow
Electron configuration Abbreviated form of orbital diagram
Orbital Diagrams and Electron Configurations
H 1 e-
Orbital diagram
1s
Electron configuration1s1
↑
O
8e-
Orbital Diagrams and Electron Configurations
Orbital Diagram
• Electron Configuration
1s1s22 2s2s22 2p2p44
1s 2s2p
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
Ne
10e-
Orbital Diagrams and Electron Configurations
Orbital Diagram
• Electron Configuration
1s1s22 2s2s22 2p2p66
1s 2s 2p
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
neon's electron configuration (1s22s22p6)
Noble Gas Configuration
[Ne] 3s1
third energy level
one electron in the s orbital
orbital shape
Na = [1s22s22p6] 3s1 electron configuration
AA
BB
CC
DD
• Shorthand Configuration
S 16e-
Valence ElectronsValence ElectronsCore ElectronsCore Electrons
S 16e- [Ne] 3s2 3p4
1s2 2s2 2p6 3s2 3p4
Electron Configuration
Longhand Configuration
Courtesy Christy Johannesson www.nisd.net/communicationsarts/pages/chem
S32.066
16