REVIEW

REVIEWREVIEW• 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).

REVIEWREVIEW• 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 NotesChemical Bonding Notes

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

Attractive Force

Na Cl

• http://www.wisc-online.com/objects/ViewObject.aspx?ID=GCH2204

Why do elements form Why do elements form chemical bonds?chemical bonds?

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

• A full valence shell causes an atom to be A full valence shell causes an atom to be more stable.more stable.•A full valence shell consists of 8 valence electrons.

Ionic BondingIonic Bonding

Ionic bondsIonic bonds: Metal atoms transfer electrons : Metal atoms transfer electrons to nonmetal atoms. Producing to nonmetal atoms. Producing oppositely charged ions oppositely charged ions

(cation & (cation & anion) which attract each anion) which attract each other.other.

Na + ClNa + Cl Na Na++ Cl Cl--

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

How to write a formula.How to write a formula.

Write cation first, followed by anionExample:

Anion : P3-

Cation : Al3+

Formula : AlP

How to write a formula.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−

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−

K+

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

2 K+ and 1 S2− = K2S

Write the formula for…Write the formula for…

an ionic compound composed of:

Al3+ and S2-

Al2S3

Write an ionic formula for Fe3+ bonding with OH−

Balance the charges.Fe3+ OH−

OH− OH−

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

1 Fe3+ and 3 OH− = Fe(OH)3

Let’s play the Ionic Bonding Let’s play the Ionic Bonding Dating Game!Dating Game!

Example: Aluminum ChlorideExample: Aluminum Chloride

Step 1:

Step 2:

Step 3: 1 3

Step 4: AlCl3

Criss-Cross Rule

Al Cl

Al Cl

3+ 1-

write out name with space

write symbols & charge of elements

criss-cross charges as subsrcipts

combine as formula unit(“1” is never shown)

Aluminum Chloride

Example: Aluminum OxideExample: Aluminum Oxide

Step 1: Aluminum Oxide

Step 2: Al3+ O2-

Step 3: Al O2 3

Step 4: Al2O3

Criss-Cross Rule

Example: Magnesium OxideExample: Magnesium Oxide

Step 1: Magnesium Oxide

Step 2: Mg2+ O2-

Step 3: Mg O2 2

Step 4: Mg2O2

Step 5: MgO

Criss-Cross Rule

InBr3BaS

Criss-Cross RuleCriss-Cross Rule

criss-cross rule: charge on cation / anion

“becomes” subscript of anion / cation

** Warning: Reduce to lowest terms.

Al2O3

Al3+ and O2–

Al2 O3

Ba2+ and S2–

Ba2 S2

In3+ and Br1–

In1 Br3

aluminum oxide barium sulfide indium bromide

Lesson Three--Transition Metal Lesson Three--Transition Metal CompoundsCompounds

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 IIFe3+is Iron IIISn4+s Tin IV

Metals with more than 1 chargeMetals with more than 1 charge

Examples:

• Cu + Copper (I)

• Cu+2 Copper (II)

• Fe+2 Iron (II)

• Fe+3 Iron (III)

PracticePractice

• K and Cl

• K and S

• Ca and S

• Cu (II) and S

Polyatomic Ions!!!!!!!!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

H +

NH4+ H N H

H

• Al + PO4-3 K + SO4

-2

• Al +3 + PO4-3 K +1 + SO4

-2

• Al(PO4

) K2

(SO4

)

• Ca + PO4-3

• Ca +2 + PO4-3

• Ca3

(PO4

)2

Covalent bonding NotesCovalent 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 MoleculesSingle Covalent Bonds Diatomic Molecules

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

Cl + Cl Cl Cl

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

Single Covalent Bonds Diatomic MoleculesSingle 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 Cl Cl

Single Covalent Bonds in compoundsSingle 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 TurnYour Turn

• Example OF2

Double Covalent BondsDouble Covalent Bonds

Two pair of electrons are being shared.

S + S S S

Triple Covalent BondsTriple Covalent Bonds

Three pair of electrons are being shared.

P + P P P

Charged CompoundsCharged Compounds

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

• Example- NH4 +1

Exceptions to the Octet RuleThe octet rule cannot be satisfied in molecules whose total

number of valence electrons is an odd number. However, these molecules do exist in nature.

Examples:

Nitrogen dioxide (NO2)

Boron trifluoride (BF3)

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

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

Nonpolar Covalent BondNonpolar Covalent Bond

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

• Cl2• O2

• N2

• H2

Polar Covalent BondPolar 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

Example: H-Br

polar molecules are also called dipoles.

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

ExamplesExamples

• H-Br

• H2S

• SCl2

• CO2

C. JohannessonC. Johannesson

• Nonpolar Covalent Bond– e- are shared equally– symmetrical e- density– usually identical atoms

C. Bond PolarityC. Bond Polarity


+ -

C. Bond PolarityC. Bond Polarity

• Polar Covalent Bond– e- are shared unequally– asymmetrical e- density– results in partial charges (dipole)


+ -

+

B. Lewis StructuresB. Lewis Structures• Nonpolar Covalent - no charges

• Polar Covalent - partial charges


A. Dipole MomentA. Dipole Moment

• Direction of the polar bond in a molecule.

• Arrow points toward the more e-neg atom.

H Cl+ -


B. Determining Molecular PolarityB. Determining Molecular Polarity

• Nonpolar Molecules– Dipole moments are symmetrical and

cancel out.

BF3

F

F F

B



• Polar Molecules– Dipole moments are asymmetrical and don

’t cancel .

netdipolemoment

H2OH H

O


CHCl3

H

Cl ClCl


• Therefore, polar molecules have...– asymmetrical shape (lone pairs) or – asymmetrical atoms

netdipolemoment

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