Transcript Ch. 5 PPT Part 3

Section 5.3
Quantum numbers and Atomic Orbitals
• Quantum numbers are numbers that
specify the properties of atomic orbitals
and of the electrons in that orbital
• It’s the electrons “address”
Four Quantum Numbers
•
•
•
•
Principal quantum number
Orbital quantum number
Magnetic quantum number
Spin quantum number
Principal quantum number
• Symbol, n
• Indicates the main energy levels
• To this point, only 1-7
• Where do we see 7 main energy levels in
this room?
Orbital quantum number
• Shape of an
orbital
• Four shapes
• s, p, d, and f
• Within each
main energy
level there are
different
shapes of
orbitals
Shapes of orbitals
• s orbital
p orbitals
Shapes of d orbitals
Examples of f-shaped orbitals
Magnetic quantum number
• Indicates the orientation (or position) of an
orbital around the nucleus
– s orbital has 1 orientation
– p orbitals have 3 orientations
– d orbitals have 5 orientations
– f orbitals have 7 orientations
• Each orbital can contain only 0, 1, or 2
electrons.
Spin quantum number
• Indicates the spin of the electron
– +1/2
– -1/2
– So if there are two electrons in one orbital,
they spin in opposite directions
• *** no two electrons can have the same 4
quantum numbers***
Electron configurations
(electron arrangements)
• Pauli Exclusion Principle
– No two electrons in the same atom will have
the same set of 4 quantum numbers
How to “read” orbitals
• How we determine which orbital gets filled with electrons first?
• Must follow the ________________:
– Orbital of Lowest energy gets filled before going to the next
lowest energy orbital
– In other words we fill from lowest energy to highest energy
– “building up” principle: electrons occupy the lowest-energy orbital
that is available.
– For example, Hydrogen’s electron goes into the __ orbital,
because it is the lowest energy orbital
Electron configurations
(electron arrangements)
• How do we know which orbitals are higher or
lower in energy?
– Read Periodic Table from Left to Right, Top to Bottom
Periodic Table Sections
3 types of notation
• Orbital Notation
• Electron-Configuration Notation
• Electron Dot Notation
Orbital Notation
• Unoccupied orbital
• Orbital with1 e• Orbital with 2 e-
__
↑ or ↓
↓↑
• Example: Hydrogen
Example: Lithium
• Example: Helium
Example: Oxygen
Electron configurations
(electron arrangements)
• Hund’s rule
– Orbitals of equal energy are each occupied by 1
electron before a 2nd electron is added.
– All electrons in singly occupied orbitals must have the
same spin
– For example, there are 3 p orbitals. If you have 3
electrons, there will be one in each orbital and all will
have spin quantum number of +1/2 or -1/2
– Example N:
Electron-Configuration notation
• Similar to orbital notation, but uses superscripts
instead of lines
• Example: Hydrogen
• Example: Helium
• Example: Lithium
Electron-Dot Notation
• Uses only the Valence electrons
• Valence electrons = the electrons in the highest
(outermost) main energy level
•
H
•
He
•
K
Practice Problems (orbital and
dot notation)
Carbon
Sodium
Sulfur
Shorthand Notation
• Use the last noble gas before your
element as a “building block”
• Example: Phosphorous
Practice Problems (d and f
orbitals)
Fe
Au
Trick to Electron Dot Notation
• Use the group number that the element is in
• Hydrogen is in group 1, 1 valence electron
• Oxygen is in group 6, 6 valence electrons
• These 8 groups are sometimes called the 8
“main groups”