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Bohr’s Atomic Model of Hydrogen
Bohr - electrons exist in energy levels AND defined orbits around the nucleus.
Each orbit corresponds to a different energy level.
The further out the orbit, the higher the energy level
De Broglie
Heisenburg
Modeled electrons as waves
Heisenberg Uncertainty Principle: states one cannot know the position and energy of an electron
Electrons exist in orbital’s of probability
Orbital - the area in space around the nucleus where there is a 90% probability of finding an electron
Schrödinger Schrödinger Wave Equation - mathematical solution of an electron’s energy in an atom
quantum mechanical model of the atom – current model of the atom treating electrons as waves.
Quantum Numbers
Wave Equation generates 4 variable solutions n - size l – shape: azimuthal
quantum
m – orientation s – spin
Address of an electron
n – Primary Quantum Number
Describes the size and energy of the orbital
n is any positive #
n = 1,2,3,4,….
Found on the periodic table
Like the “state” you live in
l – Orbital Quantum Number
Sub-level of energy
Describes the shape of the orbital
l = 0,1,2,3,4,….(n-1)
“City” you live in
n = 3
l = 0,1,2n = 2
l = 0,1n =
1
l = 0
l – Orbital Quantum Number
# level = # sublevels 1st level – 1 sublevel 2nd level – 2 sublevels 4th level = 4 sublevels
Energy Sublevels
• Labeled s, p, d, or f– Based on shape of the atom’s orbitals
– Each sublevel can only contain at most 2 e-
m – Magnetic Quantum Number
Describes the orientation of the orbital in space
Also denotes how many orbital's are in each sublevel
For each sublevel there are 2l +1 orbital's
m = 0, ±1, ±2, ±3, ±l
“Street” you live on
Look at Orbital's as Quantum Numbers
l = 0 m = 0
Can only be one s orbital
l = 1 m = -1, 0, +1
For each p sublevel there are 3 possible orientations, so
three 3 orbital's
Assigning the NumbersAssigning the Numbers The three quantum numbers (n, l, and m) are integers. The principal quantum number (n) cannot be zero. n must be 1, 2, 3, etc. The angular quantum number (l ) can be any integer between 0 and n - 1. For n = 3, l can be either 0, 1, or 2. The magnetic quantum number (ml) can be any integer between -l and +l. For l = 2, m can be either -2, -1, 0, +1, +2.
Energy Energy LevelLevel
Possible Possible sub-sub-
levelslevels
Number of Number of Sub-levelsSub-levels
nn
No. of No. of OrbitalsOrbitals
nn22
No. of No. of ElectronElectron
ss
2n2n22
44 s, p, d, fs, p, d, f 44 1616 3232
33 s, p, ds, p, d 33 99 1818
22 s, ps, p 22 44 88
11 ss 11 11 22
Hog Hilton Time
Read the scenarioComplete the questionsCompleted packet due tomorrowHW: Finish Packet
Energy Level Diagrams
Aufbau Principle
• Electrons occupy the lowest energy level orbital available.
Aufbau Principal Lowest energy orbital
available fills first
“Lazy Tenant Rule”
Pauli Exclusion PrinciplePauli Exclusion Principle
No two electrons in an atom can have the same four quantum numbers.
Wolfgang Pauli
Every house has a different address
Pauli Exclusion Principle
No two electrons have the same quantum #’s
Maximum electrons in any orbital is two
()
Hund’s Rule
When filling degenerate orbital's, electrons will fill an empty orbital before pairing up with another electron. Empty room rule
RIGHT WRONG
Outermost sub-shell being filled with electrons
The order of sublevel filling is arranged according to increasing energy level. Electrons first fill the 1s sublevel followed by the 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p and 6s
Increasing Energy
1s
2p
6s
4s
5s
3s
2s
4d
3d
5p
4p
3p
Think about piggy palace….
p ______ ______ ______
3 s ______
p ______ ______ ______
2 s ______
1 s ______
An energy diagram for Neon
Incre
asin
g E
nerg
y
Electron Spin
1s2
2s2 2px22py
22pz2
2p61s2
2s2
Electron Configuration Notation
Orbital Notation shows each orbital O (atomic number 8)
____ ____ ____ ____ ____ ____
1s 2s 2px 2py
2pz 3s
1s22s22p4 electron configuration!
Write the orbital notation for S S (atomic number 16)
____ ____ ____ ____ ____ ____ ____ ____ ____
1s 2s 2p 3s 3p
1s22s22p63s23p4
How many unpaired electrons does sulfur have? 2 unpaired electrons!2 unpaired electrons!
Electron Configuration
Shorthand way of writing electron configuration of atoms
10Ne: 1s2 2s2 2p6
Elemental Symbol and atomic number
Principal energy level
Energy sublevel
Number of electrons
Shorthand Configuration
S 16e-
Valence Electrons
Core Electrons
S 16e- [Ne] 3s2 3p4
1s22s22p63s23p4
Longhand Configuration
[Ar]4s23d104p2
Example - Germanium
X X X X X X X X X X X X X
Let’s Practice P (atomic number 15)
1s22s22p63s23p3
Ca (atomic number 20) 1s22s22p63s23p64s2
As (atomic number 33) 1s22s22p63s23p64s23d104p3
W (atomic number 74) 1s22s22p63s23p64s23d104p65s24d105p66s24f145d4
Noble Gas Configuration
[Ne] 3s23p3
[Ar] 4s2
[Ar] 4s23d104p3
[Xe] 6s24f145d4
Your Turn N (atomic number 7)
1s22s22p3
Na (atomic number 11)
1s22s22p63s1
Sb (atomic number 51)
1s22s22p63s23p64s23d104p65s24d105p3
Cr (atomic number 24)
1s22s22p63s23p64s23d4
Noble GasConfiguration
[He] 2s22p3
[Ne] 3s1
[Kr]5s24d105p3
[Ar] 4s23d4
End of information for the test on Thursday 1/14
Valence Electrons
As (atomic number 33) 1s22s22p63s23p64s23d104p3
The electrons in the outermost energy level.
s and p electrons in last shell
5 valence electrons
Full energy levelFull sublevelHalf full sublevel
1
2
3
4 5
6
7
Copper Expect: [Ar] 4s2 3d9
Actual: [Ar] 4s1 3d10
Silver Expect: [Kr] 5s2 4d9
Actual: [Kr] 5s1 4d10
Chromium Expect: [Ar] 4s2 3d4
Actual: [Ar] 4s1 3d5
Molybdenum Expect: [Kr] 5s2 4d4
Actual: [Kr] 5s1 4d5
Exceptions are explained, but not
predicted!
Atoms are more stable with half full
sublevel
Atoms create stability by losing, gaining or sharing electrons to obtain a full octet
Isoelectronic with noble gases
1
2
3
4 5
6
7
+1 +
2-3 -2 -1
0
+3
+4
Atoms take electron configuration of the closest noble gas
Na (atomic number 11) 1s22s22p63s1
1s22s22p6 = [Ne]1
2
3
4 5
6
7
Na
1 Valence electronMetal = Loses
Ne
P-3 (atomic number 15)
1s22s22p63s23p6
Ca+2 (atomic number 20)
1s22s22p63s23p6
Zn+2 (atomic number 30)
1s22s22p63s23p63d10
Last valence electrons (s and p)
Full Octet
Element Configuration notation
Orbital notation Noble gas notation
Lithium 1s22s1 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s1
Beryllium 1s22s2 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2
Boron 1s22s2p1 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p1
Carbon 1s22s2p2 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p2
Nitrogen 1s22s2p3 ____ ____ ____ ____ ____
1s 2s 2p
[He]2s2p3
Oxygen 1s22s2p4 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p4
Fluorine 1s22s2p5 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p5
Neon 1s22s2p6 ____ ____ ____ ____ ____ 1s 2s 2p
[He]2s2p6
Half of the distance between nucli in covalently bonded diatomic molecule
"covalent atomic radii"
Periodic Trends in Atomic Radius
Radius decreases across a period Increased effective nuclear charge dueto decreased shielding
Radius increases down a group Addition of principal quantum levels
Determination of Atomic Radius
Table of Atomic Radii
Increases for successive electrons taken from the same atom
Tends to increase across a period
Electrons in the same quantum level do not shield as effectively as electrons in inner levels
Irregularities at half filled and filled sublevels due to extra repulsion of electrons paired in orbitals, making them easier to remove Tends to decrease down a group
Outer electrons are farther from thenucleus
Ionization Energy: the energy required to remove an electron from an atom
Affinity tends to increase across a period
Affinity tends to decrease as you go down in a period
Electrons farther from the nucleusexperience less nuclear attraction
Some irregularities due to repulsive forces in the relatively small p orbitals
Electron Affinity - the energy change associated with the addition of an electron
A measure of the ability of an atom in a chemicalcompound to attract electrons
Electronegativities tend to increase across a period
Electronegativities tend to decrease down a group or remain the same
Electronegativity
Cations Positively charged ions
Smaller than the corresponding atomAnions
Negatively charged ions Larger than the corresponding atom
Ionic Radii
Table of Ion Sizes
Summary of Periodic Trends
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