Transcript n-2 - Elgin Local Schools

Unit 5- Electrons in Atoms
Quantum Model
of the Atom
Reminder
Thomson Model - Plum Pudding
 Rutherford Model - Nuclear model
1912-1913 Rutherford gathered
physicists, including Neils Bohr to work on
WHY atoms don’t collapse (negative eattract positive p+)
Theory
Planks Quantum Theory suggested the eexist in energy levels that have discrete
amounts of energy.
e- can jump(quantum leap) to higher
levels after gaining a Quantum of
energy.
Same amount of energy is given off
upon return to GROUND STATE.
Energy Levels.
A ladder represents the idea.
Lowest rung represents energy level 1,
n= 1 (closest to the nucleus)
Each successive rung is a higher
energy level. ( n=2, n=3, n=4…)
The energy level is the most likely
location an electron can be found within
the cloud.
Energy Levels
n = 1 can hold 2 e- max
n = 2 can hold 8 e- max
n = 3 can hold 18 e- max
n = 4 can hold 32 e- max
Try these:
Mg atomic number 12 (12 p+ and 12 e-)
n = 1 2 e- n = 2 8 e- n = 3 2 e-
Valence electrons
The outer most electrons are called
VALENCE ELECTRONS
They are the bonding electrons – VERY
IMPORTANT
B. Quantum Mechanics
Schrödinger Wave Equation (1926)
finite # of solutions  quantized energy
levels
defines probability of finding an e-
Ø 1s =
() e
1 Z 3/2 -ó
∂ a0
B. Quantum Mechanics
Orbital (“electron cloud”)
Region in space where there is 90%
probability of finding an e-
Orbital
Radial Distribution Curve
C. Quantum Numbers- Address of an e-
Four Quantum Numbers:
Specify the “address” of each electron
in an atom
UPPER LEVEL
C. Quantum Numbers
1. Principal Quantum Number ( n )
Energy level
Size of the orbital
n2 = # of orbitals in
the energy level
C. Quantum Numbers
2. Angular Momentum Quantum # ( l )
Energy sublevel
Shape of the orbital
s
p
d
f
C. Quantum Numbers
n = # of sublevels per level
n2 = # of orbitals per level
Sublevel sets: 1 s, 3 p, 5 d, 7 f
C. Quantum Numbers
3. Magnetic Quantum Number ( ml )
Orientation of orbital
Specifies the exact orbital
within each sublevel
C. Quantum Numbers
px
py
pz
C. Quantum Numbers
4. Spin Quantum Number ( ms )
Electron spin  +½ or -½
An orbital can hold 2 electrons that spin
in opposite directions.
B. Notation
Orbital Diagram (drawing of e- location)
O
8e-
1s
2s
2p
Electron Configuration (distribution of e- in
the orbitals)
2
2
4
1s 2s 2p
C. Orbital Shapes
S shape orbital
C. Orbital Shapes
p shaped orbitals (dumbbells)
px
py
pz
C. Orbital shapes
d shaped orbitals (cloverleaf)
E- configuration
Electrons fill the atom from Low energy to
High energy.
1s2 2s2 2p6 3s2 3p6 4s2 3d10 4p6 5s2 4d10
5p6 6s2 4f14 5d10 6p6 7s2 5f14 6d10 7p6
How do we keep it straight?
Feeling overwhelmed?
Read
B. Notation
Longhand Configuration
P 15e- 1s2 2s2 2p6 3s2 3p3
Core Electrons
Valence Electrons
Shorthand Configuration
P
15e
2
3
[Ne] 3s 3p
B. Notation
Longhand Configuration
Ca 20e- 1s2 2s2 2p6 3s2 3p6 4s2
Core Electrons
Valence Electrons
Shorthand Configuration
Ca
20e
2
[Ar] 4s
C. Periodic Patterns
s
p
1
2
3
4
5
6
7
f (n-2)
d (n-1)
6
7
© 1998 by Harcourt Brace & Company
A. General Rules
Aufbau Principle
Electrons fill the
lowest energy
orbitals first.
A. General Rules
Pauli Exclusion Principle
Each orbital can hold TWO electrons
with opposite spins.
A. General Rules
Hund’s Rule
Within a sublevel, place one e- per
orbital before pairing them.
WRONG
RIGHT