atomic concepts.....
honors
.....aims
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atomic concepts..... |
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honors |
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.....aims |
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NOTES
Atomic theory through time
The Dalton model of the atom = hard sphere model
The Thompson model of the atom = plum-pudding model
The Bohr model of the atom = planetary model
The modern model of the atom = wave-mechanical model
Millikan’s oil drop experiment
Millikan determined the mass and charge of electrons.
Louis de Broglie Postulate
Particles such as electrons must have both wave-like and particle-like properties.
The Heisenberg uncertainty principle
The position and the momentum (mass times velocity) of a particle cannot be accurately measured.
The Schrodinger Wave Mechanics Theory
An electron can only have certain states or orbital (a mathematical expression which represents the behavior of electrons in curved paths)
Electron Configuration
An electron has four quantum numbers corresponding to its location within an atom.
The Quantum Numbers are
The principal quantum number (n) or shell, or level
The azimuthal quantum number (l) or subshell, or sublevel
The magnetic quantum number (ml) or orbital
The magnetic direction (ms) or spin
The principal quantum number (n) represents the distance from the nucleus, shell 1 is the closest. The potential energy of e- increases as the distance from the nucleus increases therefore electrons are most stable when close to the nucleus.
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n |
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---7--- |
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---6--- |
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---5--- |
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---4--- |
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---3--- |
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---2--- |
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---1--- |
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NUCLEUS |
The azimuthal quantum number (l) or angular momentum quantum number represents the shape of an orbital.
There are 4 sublevels (s, p, d, and f)
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n |
l |
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---7--- |
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---6--- |
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---5--- |
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---4---
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---3---(f) |
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---2---(d) |
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---1---(p) |
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---0---(s) |
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---3---
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---2---(d) |
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---1---(p) |
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---0---(s) |
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---2--- |
---1---(p) |
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---0---(s) |
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---1--- |
---0---(s) |
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NUCLEUS |
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The magnetic quantum numbers (ml) represents the orientation of a sublevel in space along the x, y, and z axis of a coordinate system.
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---n--- |
---l--- |
---ml--- |
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---4--- |
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---3---
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---2--- (d) |
--- (+2) --- |
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--- (+1) --- |
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--- (0) --- |
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--- (-1) --- |
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--- (-2) --- |
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---1--- (p) |
--- (+1) --- |
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--- (0) --- |
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--- (-1) --- |
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---0--- (s) |
--- (0) --- |
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---2---
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---1--- (p) |
--- (+1) --- |
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--- (0) --- |
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--- (-1) --- |
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---0--- (s) |
--- (0) --- |
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---1--- |
---0--- (s) |
--- (0) --- |
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NUCLEUS |
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The magnetic direction (ms) represents the spin of electrons. Values can be -½ or +½.
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---n--- |
---l--- |
---ml--- |
---ms-- - |
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---4--- |
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---3---
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---2--- (d) |
--- +2 --- |
+ or - ½ |
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--- +1 --- |
+ or - ½ |
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--- 0 --- |
+ or - ½ |
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--- -1 --- |
+ or - ½ |
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--- -2 --- |
+ or - ½ |
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---1--- (p) |
--- +1 --- |
+ or - ½ |
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--- 0 --- |
+ or - ½ |
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--- -1 --- |
+ or - ½ |
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---0--- (s) |
--- 0 --- |
+ or - ½ |
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---2---
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---1--- (p) |
--- +1 --- |
+ or - ½ |
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--- 0 --- |
+ or - ½ |
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--- -1 --- |
+ or - ½ |
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---0--- (s) |
--- 0 --- |
+ or - ½ |
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---1--- |
---0--- (s) |
--- 0 --- |
+ or - ½ |
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NUCLEUS |
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Summary of Electron Configuration
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Principle Quantum Numbers |
Sublevels |
Orbitals |
Spin numbers |
Total Number of sublevels |
Total Number of Orbitals |
Total Number of Electrons |
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n |
l |
ml |
ms |
n |
n2 |
2n2 |
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1 |
0 (s) |
0 |
+ or - ½ per electron |
1 |
1 |
2 |
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2 |
0 (s) |
0 |
2 |
4 |
8 |
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1 (p) |
-1,0,+1 |
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3 |
0 (s) |
0 |
3 |
9 |
18 |
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1 (p) |
-1,0,+1 |
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2 (d) |
-2,-1,0,+1,+2 |
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4 |
0 (s) |
0 |
4 |
16 |
32 |
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1 (p) |
-1,0,+1 |
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2 (d) |
-2,-1,0,+1,+2 |
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3 (f) |
-3,-2,-1,0,+1,+2,+3 |
The Pauli Exclusion Principle
Only two electrons can occupy any one orbital, and they must have opposite spins. Therefore no two electrons can have the same values for all four quantum numbers.
The Hund’s rule
Electrons occupy the orbitals of a subshell singly and with parallel spins until each orbital has one electron, this way electrons are farther away from each other.
The Aufbau or energy order
Electrons fill the orbital that have the lowest energy first.
Example
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1s |
2s |
2p |
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H |
1 |
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1s1 |
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He |
2 |
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1s2 |
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Li |
2 |
1 |
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1s2 2s1 |
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Be |
2 |
2 |
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1s2 2s2 |
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B |
2 |
2 |
1 |
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1s2 2s2 2p1 |
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C |
2 |
2 |
1 |
1 |
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1s2 2s2 2p2 |
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N |
2 |
2 |
1 |
1 |
1 |
1s2 2s2 2p3 |
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O |
2 |
2 |
2 |
1 |
1 |
1s2 2s2 2p4 |
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F |
2 |
2 |
2 |
2 |
1 |
1s2 2s2 2p5 |
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Ne |
2 |
2 |
2 |
2 |
2 |
1s2 2s2 2p6 |
Since some energy levels have lower potential energy than others follow the table below to write electron configurations.
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1s |
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2s |
2p |
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3s |
3p |
3d |
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4s |
4p |
4d |
4f |
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5s |
5p |
5d |
5f |
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6s |
6p |
6d |
6f |
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7s |
7p |
7d |
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Example
In the electron configuration of K, 4s is filled before the 3d.
1s2 2s2 2p6 3s2 3p6 4s1
Electron configurations can be abbreviated by replacing the inner electrons by the electron configuration of the corresponding noble gas.
Example:
F is [He] 2s2 2p5
K is [Ar] 4s1
Transition and Inner Transition Elements
In addition when a sublevel is half or fully filled it gives the atom more stability therefore next level or sublevel is filled once a half or full sublevel occurs.
Example:
Cr is [Ar] 3d5 4s1 and not [Ar] 3d4 4s2
Cu is [Ar] 3d10 4s1and not [Ar] 3d9 4s2