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REVIEW. We can tell how many electrons and atom will gain or lose by looking at its valence. Metals like to lose electrons. (Cations) Ex. Na + Nonmetals like to gain electrons. (Anions) Ex: O 2- All elements try to have a full valence of 8 electrons(OCTET RULE). REVIEW.

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  • We can tell how many electrons and atom will gain or lose by looking at its valence.

  • Metals like to lose electrons. (Cations)

    • Ex. Na +

  • Nonmetals like to gain electrons. (Anions)

    • Ex: O 2-

  • All elements try to have a full valence of 8 electrons(OCTET RULE).


  • Cation- is a positively charged ion.

  • How do cations form?

    • When atoms LOSE electrons they become positive.

  • Anion- is a negatively charged ion.

  • How do anions form?

    • When atoms GAIN electrons they become negative.



Chemical Bonding Notes

A chemical bond is the force of attraction that holds two atoms together.

Attractive Force


Why do elements form chemical bonds?

Atoms form chemical bonds in order to fill their outermost energy level with electrons.

A full valence shell causes an atom to be more stable.

A full valence shell consists of 8 valence electrons.

Ionic Bonding

Ionic bonds: Metal atoms transfer electrons to nonmetal atoms. Producing oppositely charged ions (cation & anion) which attract each other.

Na + Cl Na+Cl-

Remember: Non-metal atoms take electrons from metal atoms to form an octet.

How to write a formula.

  • Write cation first, followed by anion

  • Example:

    Anion : P3-

    Cation : Al3+

Formula : AlP

How to write a formula.

Compound must be neutral, so all charges must cancel

Write an ionic formula for Na+ bonding with F−

Balance the charges.

Na+ F−

(1+) + (1-) = 0

1 Na+ and 1 F− = NaF

Write an ionic formula for

Mg2+ bonding with Cl−

Balance the charges.

Mg2+ Cl−


(2+) + 2(1-) = 0

1 Mg2+ and 2 Cl− = Mg Cl2

Write an ionic formula for K+ bonding with S2−

Balance the charges.

K+ S2−


2(1+) + (2-) = 0

2 K+ and 1 S2− = K2S

Write the formula for…

an ionic compound composed of:

Al3+ and S2-


Write an ionic formula for Fe3+ bonding with OH−

Balance the charges.




(3+) + 3(1-) = 0

1 Fe3+ and 3 OH− = Fe(OH)3

Let’s play the Ionic Bonding Dating Game!

Step 4: AlCl


Example: Aluminum Chloride

Criss-Cross Rule

Aluminum Chloride

Step 1:

write out name with space

Al Cl



Step 2:

write symbols & charge of elements

Al Cl

Step 3:



criss-cross charges as subsrcipts

combine as formula unit

(“1” is never shown)

Example: Aluminum Oxide

Criss-Cross Rule

Step 1: Aluminum Oxide

Step 2: Al3+ O2-

Step 3: Al O



Step 4: Al2O3

Example: Magnesium Oxide

Criss-Cross Rule

Step 1: Magnesium Oxide

Step 2: Mg2+ O2-

Step 3: Mg O



Step 4: Mg2O2

Step 5: MgO

Criss-Cross Rule

criss-cross rule:

charge on cation / anion

“becomes” subscript of anion / cation

** Warning: Reduce to lowest terms.

In3+ and Br1–

Al3+ and O2–

Ba2+ and S2–

In1 Br3

Al2 O3

Ba2 S2




aluminum oxide

barium sulfide

indium bromide

Lesson Three--Transition Metal Compounds

Transition metals have electrons in d orbitals and can donate different numbers of electrons, thus giving them several different positive charges.

These can be determined from the Roman numeral which is written next to the metal's name.

Example: Cu1+is Copper I

Pb2+is Lead II

Fe3+is Iron III

Sn4+s Tin IV

Metals with more than 1 charge


  • Cu + Copper (I)

  • Cu+2Copper (II)

  • Fe+2Iron (II)

  • Fe+3Iron (III)


  • K and Cl

  • K and S

  • Ca and S

  • Cu (II) and S

Polyatomic Ions!!!!!!!!

  • A polyatomic ion is a charged species (ion) composed of two or more atoms covalently bonded.

  • PO4-3 NH4+1

Lewis Dot Structures for Polyatomic ions


NH4+ H N H


  • Al + PO4-3 K + SO4-2

  • Al +3 + PO4-3K +1 + SO4-2

  • Al(PO4)K2 (SO4)

  • Ca + PO4-3

  • Ca +2 + PO4-3

  • Ca3(PO4)2

Covalent bonding Notes

  • Covalent bond: The sharing of a pair of electrons between 2 nonmetal atoms in order to fill its valence shell.

    • Each atom gains 1 electron from each covalent bond it forms with another atom.

  • When electron sharing usually occurs so that atoms attain a stable electron configuration and have 8 valence electrons.

Single Covalent Bonds Diatomic Molecules

Each chlorine needs to gain one electron by sharing electrons each atom achieves stability .

Cl + ClClCl

The pair of shared electrons is often represented as a dash. Cl-Cl

Single Covalent Bonds Diatomic Molecules

The chlorine atoms only share one pair of valence electrons. The electrons pairs not shared are called unshared electron pairs or lone pairs.

Cl + ClClCl

Single Covalent Bonds in compounds

  • H20 is a molecule containing three atoms with two single covalent bonds.

  • Count up the electrons you have!!!

  • 2 H + O H O H

  • The hydrogen and oxygen attain stable configurations by sharing electrons.

Your Turn

  • Example OF2

Double Covalent Bonds

Two pair of electrons are being shared.

S + SS S

Triple Covalent Bonds

Three pair of electrons are being shared.

P + PP P

Charged Compounds

  • Some compounds do not satisfy their stable configuration and therefore have a charge on the compound.

  • Example- NH4+1

Exceptions to the Octet Rule

The octet rule cannot be satisfied in molecules whose total number of valence electrons is an odd number. However, these molecules do exist in nature.


Nitrogen dioxide (NO2)

Boron trifluoride (BF3)

Phosphorus pentachloride (PCl5) = 10 v.e- Expanded octet

Sulfur hexafluroride (SF6)= 12 v.e- Expanded octet

Nonpolar Covalent Bond

  • When atoms bond equally it is considered a nonpolar covalent bond.

  • Cl2

  • O2

  • N2

  • H2

Polar Covalent Bond

  • When electrons are shared unequally it is a polar covalent bond.

  • An atom that strongly attracts electrons is more electronegative and therefore gains a slightly negative charge.

  • The less electronegative atom has a slightly positive charge.

  • This results in a polar bond!

An arrow is used to show which element is donating the unshared pair of electrons.

The crossed end of the arrow indicates a pos. end and the arrow points in the direction of the neg. end



polar molecules are also called dipoles.

A dipole is a molecule with two partially charged ends or poles.


  • H-Br

  • H2S

  • SCl2

  • CO2

C. Bond Polarity

  • Nonpolar Covalent Bond

    • e- are shared equally

    • symmetrical e- density

    • usually identical atoms

C. Johannesson



C. Bond Polarity

  • Polar Covalent Bond

    • e- are shared unequally

    • asymmetrical e- density

    • results in partial charges (dipole)

C. Johannesson




B. Lewis Structures

  • Nonpolar Covalent - no charges

  • Polar Covalent - partial charges

C. Johannesson



H Cl

A. Dipole Moment

  • Direction of the polar bond in a molecule.

  • Arrow points toward the more e-neg atom.

C. Johannesson






B. Determining Molecular Polarity

  • Nonpolar Molecules

    • Dipole moments are symmetrical and cancel out.

C. Johannesson








B. Determining Molecular Polarity

  • Polar Molecules

    • Dipole moments are asymmetrical and don’t cancel .

C. Johannesson









B. Determining Molecular Polarity

  • Therefore, polar molecules have...

    • asymmetrical shape (lone pairs) or

    • asymmetrical atoms

C. Johannesson

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