Section 5.3Quantum 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- ↑ or ↓• Orbital with 2 e- ↓↑

• 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”