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Chapter 4 Electron Configurations
Atom Structure The Atom
-‐ the smallest part of an element that still retains the properties of that element
Subatomic particles protons (p+) 1.674 x 10-‐24 = 1 amu neutron (n0) 1.674 x 10-‐24 = 1 amu electron (e-‐) 9.11 x 10-‐28 = 0 amu
The History of the Development of the Human Understanding of the Atom. aka History of Atomic Theory
Dalton
+ + + +
+
+ -
-
- - -
-
Thomson
+ + + + + +
-
-
- - -
-
Rutherford
+ + + + + + - -
- - -
-
Bohr
4
1913 -‐ Niels Bohr • Hydrogen atoms were known to emit specific wavelengths of light a8er being excited.
• Focusing on the par>cle proper>es of electrons, Bohr constructed a quantum model to explain this emission phenomenon.
• He proposed that electrons orbited the nucleus at specific radii, also called energy levels.
+ + + + + + - -
- - -
-
Bohr
Electromagnetic Waves • c = λ x ν • Velocity = c = speed of light
• 2.997925 x 108 m/s • All types of light energy travel at the same speed.
• Amplitude = A = measure of the intensity of the wave, i.e.“brightness”
• Magnitude of the change
amplitude
Electromagnetic Waves (cont.) • Velocity = c = speed of light
• 2.997925 x 108 m/s
• Wavelength = λ = distance between two consecutive peaks or troughs in a wave • Generally measured in nanometers
(1 nm = 10-9 m) • Same distance for troughs
Lower frequency
Higher frequency
• Frequency = ν = the number of waves that pass a point in space in one second – Generally measured in Hertz (Hz), – 1 Hz = 1 wave/sec = 1 sec-1
• c = λ x ν
wavelength
Electromagnetic Radiation
• Electromagnetic radiation is given off by atoms when they have been excited by any form of energy
Types of Electromagnetic Radiation
high energy
http://www.youtube.com/watch?v=cfXzwh3KadE
Max Planck’s Revelation • Showed that for certain applications light energy could be thought of as particles or photons
Planck’s Revelation (cont.)
• The energy of the photon is directly proportional to the frequency of light.
• Wavelength and frequency are inversely related
λ
ν
Emission of Energy by Atoms/Atomic Spectra
• Atoms that have gained extra energy release that energy in the form of light.
Atomic Spectra
• Atomic Emission Spectra: the set of frequencies of the electromagnetic spectrum emitted by excited elements of an atom
• Line spectrum: very specific wavelengths
of light that atoms give off or gain
• Each element has its own line spectrum, which can be used to identify that element.
Atomic Spectra (cont.)
Atomic Spectra (cont.)
Atomic Spectra (cont.)
• The line spectrum must be related to energy transitions in the atom
Atomic Spectra (cont.)
• The atom is quantized, i.e. only certain energies are allowed.
(a) Varies continuously (b) certain elevations i.e steps
Bohr’s Model
• Explained spectrum of hydrogen • Energy of atom is related to the distance
of electron from the nucleus
+ + + + + + - -
- - -
-
Bohr’s model
Bohr’s Model (cont.)
• Energy of the atom is quantized • Atom can only have certain specific energy
states called quantum levels or energy levels.
• When atom gains energy, electron “moves” to a higher quantum level
• When atom loses energy, electron “moves” to a lower energy level
• Lines in spectrum correspond to the difference in energy between levels
Bohr’s Model (cont.)
• Ground state: minimum energy of an atom • Therefore electrons do
not crash into the nucleus
• The ground state of hydrogen corresponds to having its one electron in the n=1 level
• Excited states: energy levels higher than the ground state
Flame test
Glowing pickel: http://www.youtube.com/watch?v=5FO0IKUTkJ0
Understanding the Arrangement of Electrons Inside Atoms • Electron configura>on
• Orbital Nota>on
Heisenberg Uncertainty Principle
+ + + + + + - -
- - -
-
Bohr’s model
• Can not determine the exact position and momentum of moving objects
• Probability of finding electrons in certain regions of the atom described by orbitals
• Orbitals have particular shapes, sizes, and energies
• Orbital: s, p, d, and f
Suppose you needed to communicate the seating in the auditorium by email without the use of the picture with just letters and/or numbers. You might symbolize the seats in the following manner:
• Sec>ons • Rows • Seat numbers • etc
24
First Row -‐ Electron ConZiguration • H: 1s1
• s orbitals are sphere-‐shaped and there is one on each and every energy level, each orbital can hold 2 electrons
• He: 1s2 25
first energy level
1 electron
“s” orbital
Second Row -‐ Electron ConZiguration • Li: 1s2 2s1 • Be: 1s2 2s2 • B: 1s2 2s2 2p1
• There are 3 “p” orbitals on any given energy level, each can hold 2 e-‐
• They are lobe-‐shaped, oriented in the x, y, z planes.
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1A" 2A" 8A"
3A" 4A" 5A" 6A" 7A"
Li"Be" B"1B" 2B" 3B" 4B" 5B" 6B" 7B" 8B" 9B" 10B"
Second Row, continued....
• Li: 1s2 2s1
• Be: 1s2 2s2
• B: 1s2 2s2 2p1
• C: 1s2 2s2 2p2
• N: 1s2 2s2 2p3
• O: 1s2 2s2 2p4
• F: 1s2 2s2 2p5
• Ne: 1s2 2s2 2p6 27
1A" 2A" 8A"
3A" 4A" 5A" 6A" 7A"
Li" Be" B" C" N" O" F" Ne"1B" 2B" 3B" 4B" 5B" 6B" 7B" 8B" 9B" 10B"
The Periodic Table is Shaped to Help You
• s -‐ two columns, 2 electrons maximum, 1 orbital • p -‐ six columns, 6 electrons maximum, 3 orbitals
28
1A" 2A" 8A"
3A" 4A" 5A" 6A" 7A"
1B" 2B" 3B" 4B" 5B" 6B" 7B" 8B" 9B" 10B"
s orbitals
p orbitals
Third Row -‐ Electron ConZiguration • Na: 1s2 2s2 2p6 3s1
• Mg: 1s2 2s2 2p6 3s2
• Al: 1s2 2s2 2p6 3s2 3p1
• Si: 1s2 2s2 2p6 3s2 3p2
• P: 1s2 2s2 2p6 3s2 3p3
• S: 1s2 2s2 2p6 3s2 3p4
• Cl: 1s2 2s2 2p6 3s2 3p5
• Ar: 1s2 2s2 2p6 3s2 3p6
29
1A" 2A" 8A"
3A" 4A" 5A" 6A" 7A"
Na"Mg" 1B" 2B" 3B" 4B" 5B" 6B" 7B" 8B" 9B" 10B" Al" Si" P" S" Cl" Ar"
⨂ ⊘⊘〇
⨂ ⊘⊘⊘ ⨂ ⨂⊘⊘ ⨂ ⨂⨂⊘ ⨂ ⨂⨂⨂
Orbital Notation Useful to ID paired and unpaired electrons
⨂ 〇〇〇
⨂ ⊘〇〇
Fourth Row
• K: 1s2 2s2 2p6 3s2 3p6 4s1 • Ca: 1s2 2s2 2p6 3s2 3p6 4s2 • Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1 • There are five “d” orbitals on any (allowed) energy level.
• 21 protons is enough + aZrac>on to pull the electrons closer to the nucleus to the 3rd energy level.
• Thus you need to remember that when you are in the 4th row of the table, you are filling the 3d orbitals.
• Lets con>nue.... 30
Why is it 3d not 4d?
Fourth Row
• K: 1s2 2s2 2p6 3s2 3p6 4s1 • Ca: 1s2 2s2 2p6 3s2 3p6 4s2 • Sc: 1s2 2s2 2p6 3s2 3p6 4s2 3d1 • Ti: 1s2 2s2 2p6 3s2 3p6 4s2 3d2 • V: 1s2 2s2 2p6 3s2 3p6 4s2 3d3 • Cr: 1s2 2s2 2p6 3s2 3p6 4s2 3d4 • Mn: 1s2 2s2 2p6 3s2 3p6 4s2 3d5 • Fe: 1s2 2s2 2p6 3s2 3p6 4s2 3d6 • Co: 1s2 2s2 2p6 3s2 3p6 4s2 3d7 • Ni: 1s2 2s2 2p6 3s2 3p6 4s2 3d8 • Cu: 1s2 2s2 2p6 3s2 3p6 4s2 3d9 • Zn: 1s2 2s2 2p6 3s2 3p6 4s2 3d10
31
Zinish the Fourth Row s (d) & p
• K: 1s2 2s2 2p6 3s2 3p6 4s1 • Ca: 1s2 2s2 2p6 3s2 3p6 4s2 • (Transi(on Metals -‐ “d” group) • Ga: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p1 • Ge: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p2 • As: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p3 • Se: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p4 • Br: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p5 • Kr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6
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What are the shape of “d” orbitals?
• yikes ! You do not need to know these shapes.
33
Fifth Row s & d • Rb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s1 • Sr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 • Y: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d1 • Zr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d2 • Nb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d3 • Mo: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d4
• Tc: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d5 • Ru: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d6 • Rh: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d7 • Rd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d8 • Ag: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d9 • Cd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
34
Fifth Row Representative Elements s (d) & p • Rb: 1s2 2s1 3s2 3s2 3p6 4s2 3d10 4p6 5s1 • Sr: 1s2 2s2 3s2 3s2 3p6 4s2 3d10 4p6 5s2 • Finish the transition Metals - “d” group Cd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
• In: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1 • This is sooooo tedius, we o8en write “condensed” electron configura>ons
• Sn: [Kr] 5s2 4d10 5p2
• Sb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p3 • Te: [Kr] 5s2 4d10 5p4
• I: [Kr] 5s2 4d10 5p5
• Xe: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 35
Fifth Row Representative Elements s (d) & p • Rb: 1s2 2s1 3s2 3s2 3p6 4s2 3d10 4p6 5s1 • Sr: 1s2 2s2 3s2 3s2 3p6 4s2 3d10 4p6 5s2 • Transi(on Metals -‐ “d” group • In: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p1 • Sn: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p2 • Sb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p3 • Te: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p4 • I: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p5 • Xe: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6
36
Sixth Row • Cs: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1 • Ba: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 • so pause to note where we are in the periodic table • clearly we need a new orbital type as we are headed into a new “block” on the table.
• This type is called “f” • La: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f1 • Ce: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f2 • Pr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f3 • Nd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f4 • Pm: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f5
• Sm: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f6
• Etc, etc, etc through • Yb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 37
So what shape are “f” orbitals? • 7 different orbitals, each of which is 4-‐lobed • you do NOT need to know these shapes
38
Sixth Row continued..... • Yb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 • So where do we go from here? • on to the “d” orbitals • Lu: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d1 • Hf: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d2 • Ta: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d3 • Etc, etc, etc through • Hg: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 • Tl: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p1 • Pb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p2 • Bi: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p3 • Etc, etc, etc
39
Sixth Row • Cs: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s1 • Ba: [Xe] 6s2
• so pause to note where we are in the periodic table • clearly we need a new orbital type as we are headed into a new “block” on the table.
• This type of orbital is called “f” • La: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f1 • Ce: [Xe] 6s2 4f2
• Pr: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f3 • Nd: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f4 • We could even represent the next element by just wri>ng the “last orbital” filled, assuming all lower energy orbitals filled up.
• Pm: 4f5
• Sm: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f6
• Etc, etc, etc through • Yb: 4f14
40
Sixth Row continued..... • Yb: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 • So where do we go from here? • on to the “d” orbitals • Lu: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d1 • Hf: (condensed version) [Xe] 5d2
• Ta: (single highest orbital) 5d3
• Etc, etc, etc through • Hg: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 • Tl: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p1 • Pb: (condensed version) [Xe] 6p2
• Bi: 1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10 5p6 6s2 4f14 5d10 6p3 • Po: (single highest orbital) 6p4 • At: (condensed version) [Xe] 6p5
• Rn: (single highest orbital) 6p6
41
Write the entire electron conZiguration • 16S
• 1s22s22p63s23p4
• 28Ni • 1s22s22p63s23p64s23d8
• 60Nd • 1s22s22p63s23p64s23d104p65s24d105p66s24f4
42
Turn these entire e.c. into the condensed version of e.c.
• 16S • 1s22s22p63s23p4 • [Ne] 3s23p4
• 28Ni • 1s22s22p63s23p64s23d8 • [Ar] 4s23d8
• 60Nd • 1s22s22p63s23p64s23d104p65s24d105p66s24f4 • [Xe] 6s24f4
43
Orbital Notation • Electrons spin just like a clock (clockwise or counterclockwise)
• Electrons will fill up the orbitals one at a >me THEN double up in an orbital
• Draw the electron out as an arrow up and down
Orbital notation…Lets Try!!
• H • He • Li • Be • B • F
Lets Continue!!
• 24Cr • 30Zn • 59Pr • 64Gd
Rules for orbital notation
• Au#au Principle: Electrons are added one at a >me to Lowest orbitals available first un>l filled Rule violated below!!!
!
Rules for orbital notation
• Pauli Exclusion Principle: Orbital can hold 2 electrons and spins must be opposite. Rule violated below!!!
Rules for orbital notation
• Hund’s Rule: Electrons like to be unpaired (or alone!!) so each orbital will fill up with one electron un>l you have to double them up Rule violated below!!!
!
Write the orbital notation for these condensed electron conZigurations
• 16S • [Ne] 3s23p4 • ⊗ ⊗WW
• 28Ni • [Ar] 4s23d8 • ⊗ ⊗⊗⊗WW
• 60Nd • [Xe] 6s24f4 • ⊗ WWWW���
50
Name the element described by the condensed version of e.c.
• [Ne] 3s23p3
• 15P • [Ar] 4s23d104p5
• 35Br • [Xe] 6s24f145d3
• 73Ta • [Rn] 7s25f8
• 96Cm 51
Name the element described by the single highest energy orbital. (Assume all lower orbitals are Zilled.) • 2p1
• 5B
• 4d2 • 40Zr
• 6p5 • 85At
• 5f2 • 90Th
• 4p8 • No such element
52
Write the single highest energy orbital to describe the element. (Assume all lower orbitals are Zilled.)
• 12Mg • 3s2
• 43Tc • 4d5
• 65Tb • 4f9
• 82Pb • 6p2
53
Electron conZigurations and Valance electrons • Now you now where all the electrons reside (electron config.)
• Valence electrons are the electrons found in the outermost energy level (s and p orbitals)
• These valence electrons are where bonding happens!
• ALL elements want a FULL outer energy level (just like Halogens) and they will gain or lose electrons to get it!!!
• MAGIC NUMBER in outer shell is 8 : Octet Rule
Valence electrons continued • All elements want to be like the Halogens! • Reason for what type of ion they become • For examples below, write the condensed E.C., how many valence electrons? Will it want to gain or lose electrons to achieve octet rule?
• 16S • 8O • 17Cl • 13Al • 20Ca • 85At • 86Rn
Other Periodic trends
• Atomic Size • Ioniza>on Energy • Electronega>vity
57
The Size of Atoms • The size of atoms
increases down the chart. • due to more energy levels
• The size of atoms decreases across to the right on the chart. • due to increased + pull on
electrons • electrons are not further
away from the nucleus. • and the shielding does not
increase
58
Ionization Energy • The amount of energy required to forcibly remove
an electron from an atom (from outermost energy level)
• How strongly an atom holds onto its outermost electron
• High ioniza>on energy = hold onto electron >ghtly! • Low ioniza>on energy = will lose an electron easily (become a ca>on)
+
Energy in electron out
-
atom becomes a positively charged ion
59
First Ionization Energy (The energy required to remove only one electron from an atom.)
• IE decreases down the chart. • Larger size of atom
(∴ e- further from protons) makes it easier to remove a valence electron.
• IE increases across to the right on the chart. • The smaller size and the
increased effective nuclear charge of the atom makes it harder to remove an electron.
Fi rs t Ioni zati on Energ ies (k J/mole)
1 H 1311
He 2370 1
2 Li 521
Be 899
B 799
C 1087
N 1404
O 1314
F 1682
Ne 2080 2
3 Na 496
Mg 737
Al 576
Si 786
P 1052
S 1000
Cl 1245
Ar 1521 3
4 K 419
Ca 590
Ga 579
Ge 762
As 944
Se 941
Br 1140
Kr 1351 4
5 Rb 403
Sr 550
In 558
Sn 709
Sb 832
Te 869
I 1009
Xe 1170 5
6 Cs 376
Ba 503
Tl 589
Pb 716
Bi 703
Po 812
At Rn 1037 6
7 Fr Ra 7
IE decreases down
IE increases across
Electronegativity • The ability to attract
electrons in a chemical bond
• EN increases as you move to the right of the periodic table (i.e. easier to gain e- to have a full outer energy level)
• EN decreases go down (i.e. increase size so further away from nucleus)
