CHAPTER 5CHAPTER 5ELECTRONS IN ATOMSELECTRONS IN ATOMS
P. 126P. 126
ERNEST RUTHERFORD’S MODEL
Discovered dense + nucleus
•e-s move like planets around sun
•Mostly empty space
Didn’t explain chemical properties of elements
THE BOHR MODEL
Neils Bohr (1885-1962)
•Danish physicist
•Student of Rutherford “W
hy d
on’t
e-s
fall
into
nu
cleu
s”?Why don’t Why don’t
electrons electrons fall into fall into nucleus?nucleus?
Niels Bohr
THE BOHR MODEL
Niels Bohr
I pictured the electrons found in specific circular paths around the nucleus, and can jump from one level to another.
Furthermore, each level has a fixed amount of energy different from other levels
BOHR’S MODEL
fixed energy e- have called Energy levels
Like rungs of laddere- can’t exist btwn energy levels
energy levels not evenly spaced
• High levels closer (less energy needed to jump)
Bohr’s model of the atom 5:17
THE QUANTUM MECHANICAL MODEL• e-’s don’t move like big objects
• Rutherford & Bohr model• Energy - “quantized” (in chunks)
• exact energy needed to move e- 1 energy level called a quantum
• energy never “in btwn” • quantum leap in energy must exist
• Erwin Schrodinger (1926) mathematically described energy & position of e- in atom
SchrodingerSchrodinger
Quantum Leap TV intro
THE QUANTUM MECHANICAL MODEL
energy levels for e-
Orbits not circular Based on probability of
finding e- certain distance from nucleus electron cloud
ATOMIC ORBITALS• Principal Quantum # (n) - energy
level of e- (1, 2, 3 etc.)• atomic orbitals - regions of space regions of space w/w/ high high
probability probability of finding e- (not a true “orbit”)(not a true “orbit”)• within each energy level• Sublevels like rooms in a hotel• s, p, d, and f• Different shapes
Max # of e- that fit in energy level is: 2n2
How many eHow many e-- in level 2? in level 2? level 3?level 3?
s and p orbitals 1:20
d orbitals 3:40
atomic orbitals review (14:28)
ATOMIC ORBITALS
sspherical
pdumbell
dclover leaf
fcomplicated
# of orbitals (regions of
space)
Maximum electrons
First possible
energy level
1 2 1st
3 6 2nd
5 10 3rd
7 14 4th
Summary of Principal Energy Levels, Sublevels, and Summary of Principal Energy Levels, Sublevels, and OrbitalsOrbitals
Principal Principal energy energy levellevel
Number Number of of
sublevelssublevels
Type of Type of sublevelsublevel
Max # of Max # of electronselectrons
Electron Electron configurationconfiguration
n = 1n = 1 11 1s (1 orbital)1s (1 orbital) 22 1s1s22
n = 2n = 2 22 2s (1 orbital2s (1 orbital2p (3 orbitals)2p (3 orbitals)
88 2s2s22
2p2p66
n = 3n = 3 33 3s (1 orbital)3s (1 orbital)3p (3 orbitals)3p (3 orbitals)3d (5 orbitals)3d (5 orbitals)
1818 3s3s22
3p3p66
3d3d1010
n = 4n = 4 44 4s (1 orbital)4s (1 orbital)4p (3 orbitals)4p (3 orbitals)4d (5 orbitals) 4d (5 orbitals) 4f (7 orbitals)4f (7 orbitals)
3232 4s4s22
4p4p66
4d4d1010
4f4f1414
ORDER OF ELECTRON SUBSHELL FILLING:
NOT “IN ORDER”
1s2
2s2 2p6
3p6
4p6
5p6
6p6
7p6
3s2
4s2
5s2
6s2
7s2
3d10
4d10
5d10
6d10
4f14
5f14
1s2 2s2 2p6 3p63s2 4s2 4p65s23d10 5p6 6s24d10 6p6 7s25d104f14 7p66d105f14
energy levels overlap
Lowest energy fill first
Increa
sing e
nerg
y
ELECTRON CONFIGURATION
1s1
Principal energy level
row #
1-7
7 rows sublevel
s, p, d, or f
4 sublevels
group #
# valence e-
s: 1 or 2
p: 1-6
d: 1-10
f: 1-14
Total e- = Atomic #
1
2
3
4
5
6
7
6
7
per
iod
# =
# e
- en
erg
y le
vels 1A
2A
3B 4B5B 6B 7B
8B
1B 2B
3A 4A 5A6A 7A8A
group # = # valence e-
d
f
3d
4d
5d6d
4f
5f
SUBLEVELS D AND F ARE “SPECIAL”
Incr
easi
ng e
nerg
y
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
aufbau diagram - page 133
Aufbau - German for “building up”
SECTION 5.2 ELECTRON SECTION 5.2 ELECTRON ARRANGEMENT IN ATOMS P. 133ARRANGEMENT IN ATOMS P. 133
ELECTRON CONFIGURATIONS…….3 rules explain how e-’s fill their
orbitals:
1) Aufbau principle – e-’s enter lowest energy level first.
2) Pauli Exclusion Principle - 2 e-’s max/orbital (hotel room) - different spins
PAULI EXCLUSION PRINCIPLENo 2 electrons in an atom can have the same four quantum numbers.
Wolfgang Pauli
To show different direction of spin, a pair in the same orbital is written as:
QUANTUM NUMBERS
Each e- has unique set of 4 quantum #’s describing it
1) Principal quantum #2) Angular momentum quantum #3) Magnetic quantum #4) Spin quantum #
ELECTRON CONFIGURATIONS
3) Hund’s Rule- When e-’s occupy orbitals of same energy, they won’t pair up until they must
write e- configuration for Phosphorus
all 15 e-’s must be accounted for
The first 2 e-’s go into the 1s orbital
Notice opposite direction of spinsIn
crea
sing
ene
rgy
1s
2s
3s
4s
5s6s
7s
2p
3p
4p
5p
6p
3d
4d
5d
7p 6d
4f
5f
The next e-’s go in 2s orbital
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 e-’s go in 2p orbital
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 e-’s go in 3s orbital
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 3 e-’s go in 3p orbitals
They each go into separate shapes (Hund’s)
• 3 unpaired e-’s
= 1s22s22p63s23p3 Orbital notation
An internet program about electron configurations is:
Electron Configurations
I electron config (song) 3:24
FILLING ORBITALS
Lowest higher energy
Adding e-’s changes energy of orbital
•Full orbitals best situation•half filled orbitals next best
• more stable• Changes filling order
WRITE THE ELECTRON CONFIGURATIONS FOR THESE ELEMENTS: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 (expected)
But this is not what happens!!
CHROMIUM IS ACTUALLY:
1s22s22p63s23p64s13d5 Why?
2 half filled orbitals
•Half full slightly lower in energy
•Same applies to copper
COPPER’S E- CONFIGURATION
• Copper has 29 e-s so expect: 1s22s22p63s23p63d94s2
• actual configuration is:
1s22s22p63s23p63d104s1
• 1 more full orbital & 1 half filled
• Exceptions
• d4 • d9
IRREGULAR CONFIGURATIONS OF CHROMIUM AND COPPER
Chromium steals a 4s e- to make its 3d sublevel HALF FULL
Copper steals a 4s electron to FILL its 3d sublevel
SECTION 5.3SECTION 5.3 PHYSICS AND THE QUANTUM PHYSICS AND THE QUANTUM MECHANICAL MODELMECHANICAL MODEL
P. 138P. 138
• Study of light led to quantum mechanical model
• Light is electromagnetic radiation
• EM radiation: gamma rays, x-rays, radio waves, microwaves
• Speed of light = 2.998 x 108 m/s
• “c” - celeritas (Latin for speed)
• All EM radiation travels same in vacuum
Light
- Page 139
“R O Y G B I V”
Frequency Increases
Wavelength Longer
PARTS OF A WAVE
Wavelength
Amplitude
Crest
Trough
Equation:
c =
c = is a constant (2.998 x 108 m/s)
(nu) = frequency, in units of hertz (hz or sec-1) (lambda) = wavelength, in meters
ELECTROMAGNETIC RADIATION PROPAGATES THROUGH SPACE AS A WAVE MOVING AT THE SPEED OF LIGHT.
WAVELENGTH AND FREQUENCY
• inversely related
• one gets bigger, other smaller
• Different frequencies = different colors
• wide range of frequencies (spectrum)
- Page 140
Use Equation: c =
Radiowaves
Microwaves
Infrared .
Ultra-violet
X-Rays
GammaRays
Low Frequency
High Frequency
Long Wavelength
Short WavelengthVisible Light
Low Energy
High Energy
Long =Low Frequency
=Low ENERGY
Short =High Frequency
=High ENERGY
ATOMIC SPECTRA
White light all colors of visible spectrum
• prismprism separates it
according to λ
IF THE LIGHT IS NOT WHITE
heating gas with electricity will emit colors
• this light thru prism is different
ATOMIC SPECTRUM
elements emit own characteristic colors
• composition of stars determined thru spectral analysis
• atomic emission spectrum
• Unique to each element, like fingerprints!
• ID’s elements
LIGHT IS A PARTICLE?Energy is quantizedLight is energy…..light must be quantized
photons photons smallest pieces of lightPhotoelectric effect –
• Matter emits e- when it absorbs energy• Albert Einstein Nobel Prize in chem
Energy & frequency: directly related
Planck-Einstein Equation: E = hEE = Energy, in units of Joules (kg·m = Energy, in units of Joules (kg·m22/s/s22)) (Joule…metric unit of energy)(Joule…metric unit of energy)
hh = Planck’s constant (6.626 x 10 = Planck’s constant (6.626 x 10-34-34 J·s) J·s)(reflecting sizes of energy quanta)(reflecting sizes of energy quanta)
= frequency, units of hertz (hz, sec= frequency, units of hertz (hz, sec-1-1))
ENERGY (E ) OF ELECTROMAGNETIC RADIATION DIRECTLY PROPORTIONAL TO FREQUENCY () OF RADIATION.
THE MATH IN CHAPTER 5There are 2 equations:
1) c = 2) E = h Put these on your 3 x 5
notecard!
EXAMPLES1) What is the wavelength of
blue light with a frequency of 8.3 x 1015 hz?
2) What is the frequency of red light with a wavelength of 4.2 x 10-5 m?
3) What is the energy of a photon of each of the above?
EXPLANATION OF ATOMIC SPECTRA
electron configurations written in lowest energy.
energy level, and where electron
starts from, called it’s ground ground statestate - lowest energy level.
CHANGING THE ENERGYLet’s look at a hydrogen atom, with only one electron, and in the first energy level.
Changing the energyHeat, electricity, or light can move e-’ up to different energy levels. The electron is now said to be ““excitedexcited””
Changing the energyAs electron falls back to ground state,
it gives energy back as lightlight
Experiment #6, page 49-
may fall down in specific steps
Each step has different energy
Changing the energy
{{{
Lyman series (UV) Balmer series
(visible)
Paschen series
(infrared)
further they fall, more energy released = higher frequency
orbitals also have different energies inside energy levels
All electrons can move around.
Ultraviolet Visible Infrared
WHAT IS LIGHT?Light is a particle - it comes in chunks.
Light is a wave - we can measure its wavelength and it behaves as a wave
combine E=mc2 , c=, E = 1/2 mv2 and E = hthen we can get:
= h/mv (from Louis de Broglie)Calculates wavelength of a particle.
called de Broglie’s equation • He said particles exhibit properties of waves
WAVE-PARTICLE DUALITYJ.J. Thomson won the Nobel prize for describing the electron as a particle.
His son, George Thomson won the Nobel prize for describing the wave-like nature of the electron.
The electron is a particle!
The electron is an energy
wave!
CONFUSED? YOU’VE GOT COMPANY!
“No familiar conceptions can be woven around the electron;
something unknown is doing we don’t know what.”
Physicist Sir Arthur Eddington
The Nature of the Physical World
1934
THE PHYSICS OF THE VERY SMALL
Quantum mechanics explains how very small particles behave
•Quantum mechanics is an explanation for subatomic particles and atoms as waves
Classical mechanics describes the motions of bodies much larger than atoms
HEISENBERG UNCERTAINTY PRINCIPLEimpossible to know exact location and velocity of particle
better we know one, less we know other
Measuring changes properties.
True in quantum mechanics, but not classical mechanics
HEISENBERG UNCERTAINTY PRINCIPLE
You can find out where the electron is, but not where it is going.
OR…
You can find out where the electron is going, but not where it is!
“One cannot simultaneously determine both the position and momentum of an electron.”
Werner Heisenberg
IT IS MORE OBVIOUS WITH THE VERY SMALL OBJECTS
To measure where e-, we use light
But light energy (photon) moves e- due to small mass
And hitting e- changes frequency of light
Moving Electron
Photon
Before
Electron velocity changes
Photon wavelengthchanges
After
Fig. 5.16, p. 145