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CHAPTER 1
2
ELECTR
ONS IN A
TOMS
BRODERSEN
HONORS CHEM 2
013/1
4
BOHR’S MODEL
Why don’t the electrons fall into the nucleus?
Move like planets around the sun.In circular orbits at different
levels.Amounts of energy separate one
level from another.
BOHR’S MODEL
Nucleus
Electron
Orbit
Energy Levels
Nucleus
Electron
Orbit
Energy Levels
BOHR POSTULATED THAT:
Fixed energy related to the orbit
Electrons cannot exist between orbits
The higher the energy level, the further it is away from the nucleus
An atom with maximum number of electrons in the outermost orbital energy level is stable (unreactive)
ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATIONELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION
ELECTROMAGNETIC RADIATION.
ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATIONELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATIONMost subatomic particles behave as PARTICLES and obey Most subatomic particles behave as PARTICLES and obey the physics of waves.the physics of waves.
wavelengthVisible light
wavelength
Ultaviolet radiation
Amplitude
Node
ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATIONELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION
Waves have a frequencyWaves have a frequency
Use the Greek letter “nu”, Use the Greek letter “nu”, , for , for frequency, and units are “cycles per frequency, and units are “cycles per sec”sec”
All radiation: All radiation: • • = c = cwhere c = velocity of light = 3.00 x 10where c = velocity of light = 3.00 x 1088 m/secm/sec
ELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATIONELECTROMAGNETIC ELECTROMAGNETIC RADIATIONRADIATION
HOW DID HE DEVELOP HIS THEORY?
•He used mathematics to explain the visible spectrum of hydrogen gas•http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/linesp16.swf
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low energy
High energy
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
THE LINE SPECTRUM
•electricity passed through a gaseous element emits light at a certain wavelength•Can be seen when passed through a prism•Every gas has a unique pattern (color)
LINE SPECTRUM OF VARIOUS ELEMENTS
BOHR’S TRIUMPH
His theory helped to explain periodic law
Halogens are so reactive because it has one e- less than a full outer orbital
Alkali metals are also reactive because they have only one e- in outer orbital
DRAWBACK
•Bohr’s theory did not explain or show the shape or the path traveled by the electrons.•His theory could only explain hydrogen and not the more complex atoms
ATOMIC LINE EMISSION ATOMIC LINE EMISSION SPECTRA AND NIELS BOHRSPECTRA AND NIELS BOHRATOMIC LINE EMISSION ATOMIC LINE EMISSION SPECTRA AND NIELS BOHRSPECTRA AND NIELS BOHR
Bohr’s greatest contribution Bohr’s greatest contribution to science was in building a to science was in building a simple model of the atom. simple model of the atom. It was based on an It was based on an understanding of theunderstanding of the LINE LINE EMISSION SPECTRAEMISSION SPECTRA of of excited atoms.excited atoms.
Problem is that the model Problem is that the model only works for Honly works for HNiels BohrNiels Bohr
(1885-1962)(1885-1962)
ATOMIC SPECTRAATOMIC SPECTRAATOMIC SPECTRAATOMIC SPECTRA
+Electronorbit
One view of atomic structure in early 20th century was that an One view of atomic structure in early 20th century was that an electron (e-) traveled about the nucleus in an orbit.electron (e-) traveled about the nucleus in an orbit.
ATOMIC SPECTRA AND ATOMIC SPECTRA AND BOHRBOHRATOMIC SPECTRA AND ATOMIC SPECTRA AND BOHRBOHR
Bohr said classical view is Bohr said classical view is wrong. wrong.
Need a new theory — now called Need a new theory — now called QUANTUMQUANTUM or or WAVE WAVE MECHANICSMECHANICS..
e- can only exist in certain e- can only exist in certain discrete orbitsdiscrete orbits
e- is restricted to e- is restricted to QUANTIZEDQUANTIZED energy state (quanta = bundles energy state (quanta = bundles of energy)of energy)
Schrodinger applied idea of e- Schrodinger applied idea of e- behaving as a wave to the behaving as a wave to the problem of electrons in atoms.problem of electrons in atoms.
He developed the He developed the WAVE WAVE EQUATIONEQUATION
Solution gives set of math Solution gives set of math expressions called expressions called WAVE WAVE
FUNCTIONS, FUNCTIONS, Each describes an allowed energy Each describes an allowed energy
state of an e-state of an e-
E. SchrodingerE. Schrodinger1887-19611887-1961
QUANTUM OR WAVE QUANTUM OR WAVE MECHANICSMECHANICSQUANTUM OR WAVE QUANTUM OR WAVE MECHANICSMECHANICS
HEISENBERG UNCERTAINTY HEISENBERG UNCERTAINTY PRINCIPLEPRINCIPLE
Problem of defining nature Problem of defining nature of electrons in atoms of electrons in atoms solved by W. Heisenberg.solved by W. Heisenberg.
Cannot simultaneously Cannot simultaneously define the position and define the position and momentum (= m•v) of an momentum (= m•v) of an electron.electron.
We define e- energy exactly We define e- energy exactly but accept limitation that but accept limitation that we do not know exact we do not know exact position.position.
Problem of defining nature Problem of defining nature of electrons in atoms of electrons in atoms solved by W. Heisenberg.solved by W. Heisenberg.
Cannot simultaneously Cannot simultaneously define the position and define the position and momentum (= m•v) of an momentum (= m•v) of an electron.electron.
We define e- energy exactly We define e- energy exactly but accept limitation that but accept limitation that we do not know exact we do not know exact position.position.
W. HeisenbergW. Heisenberg1901-19761901-1976
ARRANGEMENT OF ARRANGEMENT OF ELECTRONS IN ATOMSELECTRONS IN ATOMSARRANGEMENT OF ARRANGEMENT OF ELECTRONS IN ATOMSELECTRONS IN ATOMSElectrons in atoms are arranged asElectrons in atoms are arranged as
LEVELSLEVELS (n) (n)
SUBLEVELSSUBLEVELS (l) (l)
ORBITALSORBITALS (m (mll))
QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSThe The shape, size, and energyshape, size, and energy of each orbital is a function of 3 quantum of each orbital is a function of 3 quantum
numbers which describe the location of an electron within an atom numbers which describe the location of an electron within an atom or ionor ion
n n (principal)(principal) ---> energy level---> energy level
ll (orbital) (orbital) ---> shape of orbital---> shape of orbital
mmll (magnetic)(magnetic) ---> designates a particular ---> designates a particular suborbitalsuborbital
The fourth quantum number is not derived from the wave functionThe fourth quantum number is not derived from the wave function
ss (spin)(spin) ---> spin of the electron ---> spin of the electron (clockwise or counterclockwise: ½ or – ½)(clockwise or counterclockwise: ½ or – ½)
QUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERSQUANTUM NUMBERS
So… if two electrons are in the same So… if two electrons are in the same place at the same time, they must be place at the same time, they must be repelling, so at least the spin repelling, so at least the spin quantum number is different!quantum number is different!
The The Pauli Exclusion PrinciplePauli Exclusion Principle says that says that no two electrons within an atom (or no two electrons within an atom (or ion) can have the same four quantum ion) can have the same four quantum numbers.numbers.
If two electrons are in the same energy If two electrons are in the same energy level, the same sublevel, and the level, the same sublevel, and the same orbital, they must repel.same orbital, they must repel.
Think of the 4 quantum numbers as the Think of the 4 quantum numbers as the address of an electron… Country > address of an electron… Country > State > City > StreetState > City > Street
ENERGY LEVELS
Further away from the nucleus means more energy.
There is no “in between” energy
Energy Levels
First
Second
Third
Fourth
Fifth
Incr
easi
ng e
nerg
y }
THE QUANTUM MECHANICAL MODELEnergy is quantized. It comes in chunks.
A quanta is the amount of energy needed to move from one energy level to another.
Since the energy of an atom is never “in between” there must be a quantum leap in energy.
Schrödinger derived an equation that described the energy and position of the electrons in an atom
ATOMIC ORBITALS
Principal Quantum Number (n) = the energy level of the electron.
Within each energy level the complex math of Schrödinger's equation describes several shapes.
These are called atomic orbitalsRegions where there is a high
probability of finding an electron
S ORBITALS
1 s orbital for
every energy level
1s 2s 3s
Spherical shaped
Each s orbital can hold 2 electrons
Called the 1s, 2s, 3s, etc.. orbitals
P ORBITALS
Start at the second energy level
3 different directions
3 different shapes
Each orbital can hold 2 electrons
The p Sublevel has 3 p orbitals
THE D SUBLEVEL CONTAINS 5 D ORBITALS•The D sublevel starts in the 3rd energy level •5 different shapes (orbitals)•Each orbital can hold 2 electrons
THE F SUBLEVEL HAS 7 F ORBITALS
•The F sublevel starts in the fourth energy level
•The F sublevel has seven different shapes (orbitals)
•2 electrons per orbital
SUMMARY
s
p
d
f
# of shapes (orbitals)
Max # of electrons
1 2 1
3 6 2
5 10 3
7 14 4
Sublevel
Starts at energy level
ELECTRON CONFIGURATIONS
The way electrons are arranged in atoms.
Aufbau principle- electrons enter the lowest energy first.
This causes difficulties because of the overlap of orbitals of different energies.
Pauli Exclusion Principle- at most 2 electrons per orbital - different spins
ELECTRON CONFIGURATIONS
First Energy Levelonly s sublevel (1 s orbital)only 2 electrons1s2
Second Energy Levels and p sublevels (s and p orbitals are
available)2 in s, 6 in p2s22p6
8 total electrons
Third energy level
s, p, and d orbitals
2 in s, 6 in p, and 10 in d
3s23p63d10
18 total electrons
Fourth energy level
s,p,d, and f orbitals
2 in s, 6 in p, 10 in d, and 14 in f
4s24p64d104f14
32 total electrons
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
ELECTRON CONFIGURATIONHund’s Rule- When electrons occupy orbitals of equal energy
they don’t pair up until they have to .
The first to electrons go into the 1s orbital
Notice the opposite spins
only 13 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
The next electrons go into the 2s orbital
only 11 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 2p orbital
• only 5 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The next electrons go into the 3s orbital
• only 3 more
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
• The last three electrons go into the 3p orbitals.
• They each go into separate shapes
• 3 unpaired electrons
• 1s22s22p63s23p3
ORBITALS FILL IN ORDER Lowest energy to higher energy.
Adding electrons can change the energy of the orbital.
Half filled orbitals have a lower energy.
Makes them more stable.
Changes the filling order
WRITE THESE ELECTRON CONFIGURATIONS
Titanium - 22 electrons
1s2
2s2
2p6
3s2
3p6
4s2
3d2
Vanadium - 23 electrons 1s2
2s2
2p6
3s2
3p6
4s2
3d3
Chromium - 24 electrons
1s2
2s2
2p6
3s2
3p6
4s2
3d4 is expected
But this is wrong!!
CHROMIUM IS ACTUALLY
1s2
2s2
2p6
3s2
3p6
4s1
3d5
Why?
This gives us two half filled orbitals.
Slightly lower in energy.
The same principal applies to copper.
COPPER’S ELECTRON CONFIGURATIONCopper has 29 electrons so we expect
1s2
2s2
2p6
3s2
3p6
4s2
3d9
But the actual configuration is
1s2
2s2
2p6
3s2
3p6
4s1
3d10
This gives one filled orbital and one half filled orbital.
Remember these exceptions
GREAT S
ITE T
O PRACTI
CE
AND INSTA
NTLY
SEE
RESULTS F
OR
ELECTR
ON CONFIG
URATIO
N.
PRACTICE
Time to practice on your own filling up electron configurations.
Do electron configurations for the first 20 elements on the periodic table.