Electrons and Bonding

Valence Electrons

• The electrons that are located in the outer energy shell of each atom• These electrons are available to be shared,

lost, or gained in the formation of a chemical compound.

Electron Shells

• Remember as you move down a period on the periodic table you add an energy level (shell) to that atom.

• Each shell has a maximum number of electrons that it can hold before becoming full (and stable)

• All Noble Gases are stable, this means their outer energy levels are full.• They do not want to share, gain or lose any of their

electrons.

Heisenberg Uncertainty Principle

• This is the theory that states that it is impossible to determine simultaneously both the position and velocity of an electron or any other particle.• In other words, this principle is stating that it is

impossible to know where the electron is exactly, how fast it is going nor the direction it is moving.

• Velocity is the speed and direction of the electron

Quantum Theory

• Schrodinger derived an equation that described energy & position of electrons in atom• Schrodinger along with other scientists

laid the foundation for the modern quantum theory, which describes mathematically the wave properties of electrons and other very small particles.

Orbitals and Quantum Numbers

• Electrons exist around the nucleus in certain regions called atomic orbitals - 3D regions around the nucleus that indicates the probable location of an electron.• Quantum numbers are used to describe

atomic orbitals. The quantum numbers are like an address for the electrons.

Orbitals

• There are 4 main types of orbitals• s• p• d• f

• Each one has:• a certain number of electrons that they can

hold • A certain shape

s-orbital

• 1 s orbital for every energy level• Spherical shaped• can hold 2 electrons• Called the 1s, 2s, 3s, etc.. orbitals.

P orbitals

• Start at second energy level • 3 different directions• 3 different shapes• Each can hold 2 electrons

P orbitals

D orbitals

• Start at third energy level • 5 different shapes• Each can hold 2 electrons

F orbitals

• Start at fourth energy level• 7 different shapes• 2 electrons per shape

Summary of Orbitals

Orbital

S

P

D

F

# of ShapesMax electrons

Starts at energy level

1

3

5

7

2

6

10

14

1

2

3

4

By energy level

• Orbitals do not fill up in a neat order.• Energy levels overlap• Lowest energy fill first.

Electron Configuration

•Way electrons are arranged in atoms.• Aufbau principle- electrons occupy

the lowest-energy orbital that can receive it. This means electrons enter the lowest energy first.

• This causes difficulties because of overlap of orbitals of different energies.

• Pauli Exclusion Principle- states that no 2 electrons in the same atom can have the same set of four quantum numbers. Basically, at most there are 2 electrons per orbital with different spins.

• Hund’s Rule- Orbitals of equal energy are each occupied by one electron before any orbital is occupied by a second electron. • This means when electrons occupy orbitals

of equal energy they don’t pair up until they have to. • (BUS RULE)

Start at the beginning of the periodic table and write down the energy level, subshell and number of electrons for each layer that you pass until you reach the element that you want.

Example: Manganese would be 1s22s22p63s23p64s23d5

Remember each s orbital can hold 2 electrons, p can hold 6, etc…

Add up all of the subscripts to find the number of electrons. 2+2+6+2+6+2+5=25

Click icon to add picture

2s2

Energy Level or shell Shape or

subshell

# of electrons

Manganese would be 1s22s22p63s23p64s23d5

What is the electron configuration for bromine?

1s2 2s2 2p6 3s2 3p6 3d10 4p54s2

Ionic Bonding

• Transfer of electrons• Involves the attraction of oppositely charged particles or

ions (cations &anions)• Referred to as “salts” or ionic compounds• Occur between a metal and a nonmetal• Electronegativity difference of 1.6 or higher• Individual piece called a formula unit

Characteristics/Properties of an Ionic Bond

• Called Salts• Forms a crystalline solids at room temperature• High melting and boiling points• Generally soluble in water• Brittle• Conduct electricity when molten or dissolved in water• Breaks into ions when dissolved• Electrolytes

Covalent Bonds

• Sharing of electrons• Occurs between two or more nonmetals• Referred to as molecular compounds• Individual piece called a molecule• Two types• Polar• Unequal sharing• Electronegativity difference of 0.2 to 1.6

• Nonpolar• Electronegativity difference of 0 to 0.2• Equal sharing

Characteristics/Properties of Covalent Bonds

• Found as amorphous solids, liquids & gases• Low melting and boiling points• Generally only slightly to insoluble in water• Generally do NOT conduct electricity• Do not break into pieces when dissolved• Remain as molecules

Metallic Bonding

• Made of closely packed cations• Electron Sea Model – electrons are free to move between

the atoms – this explains the properties.• Electrons can become delocalized• Characteristics/Properties• Good conductors of electricity & heat• Malleable• Ductile