Ionic Compounds

• Ions are atoms that have gained or lost electron(s)

• Atoms tend to make ions with characteristic oxidation states (charges)

• Metals are losers (+ ions)• Nonmetals are gainers (- ions)• Some atoms do not readily make ions

(C, Si, many metalloids)

Ionic Bonds

• An ionic bond is formed between two or more oppositely charged ions

• Ionic compounds are made of a metal (+) and a nonmetal (-)

• Ionic compounds are called salts• The overall charge on an ionic

compound is zero • When a metal and nonmetal react,

electrons are transferred

Making Ionic Bonds

Na Cl11p+

11e-

17p+

17e-

10e-

+

18e-

-

Making More Ionic Bonds

Mg O

12 p+

12e-

8p+

8e-

10e-

+2

10e-

-2

Formulas and Names

• NaCl, sodium chloride• MgO, magnesium oxide• Binary salts

– metal, then nonmetal– nonmetal ending changed to

“ide”– no subscripts when ratio is 1:1

Unequal charges

NaS

11 p+

11e-

16p+

16e-

10e-

+1-1

Na11 p+

11e-

10e-

+1

18e-

-2

Formula: Na2S

Name: sodium sulfideTotal + charge = 2, Total - charge = 2

Total charge overall = 1 + 1 + (-2) = 0

Unequal charges

Determine the formula of calcium bromide.

Ca Br

BrFormula:CaBr2

+2-1

-1

+1

Polyatomic ions

Transition metal salts

Salts of polyatomic ions

Solubilities

• Salts are soluble in water if ion-water interactions can supply enough energy to break apart the crystal lattice

• Salts of lower-charged ions are more likely to be soluble (lower lattice energy)

• All alkali metal and ammonium salts are soluble

• All nitrates are soluble• All oxides are insoluble (alkali metal oxides

react to form hydroxides)

Ionic compound properties

• Made of metal and nonmetal (except ammonium and organic base salts)

• High MP (chemical bonds are broken in melting)

• Crystal lattice• Brittle• Form ions in water solution (ionization)

NaCl Na+ + Cl -

• Conduct electricity when melted

Hydrates

• Water can get trapped in crystal lattice of a crystallized salt

Na2CO3.10H2O

Sodium carbonate decahydrateCuSO4

.5H2O

Copper (II) sulfate pentahydrateSodium acetate trihydrate

NaC2H3O2.3H2O

Hydrates

• Some salts take water out of the air to become hydrates: hygroscopic

• Example: Na2CO3

• Others take enough water to become solutions: deliquescent

• Example: CaCl2

Crystal Lattices and Energy

• Regular repeating arrangement of ions is a crystal lattice

• Energy holding lattice together is the lattice energy

• Energy is released when lattice is formed (from gaseous ions) and absorbed when it is broken

Crystal Lattices and Energy

• Lattice energy is measured from the viewpoint of the system

• When gaseous ions come together to form a crystal energy leaves the system

• Since system energy is lower, lattice energy is always given as a negative value

Crystal Lattices and Energy

• Magnitude of lattice energy is directly proportional to charge density

• Charge density is related to charge magnitude and ion size

• Crystallization from gaseous ions is always negative; crystallization from solution can be negative or positive

Metallic Bonds

• Metals form molecular orbitals that cover the entire crystal

• Electrons can move anywhere in the orbital, so metals conduct heat and electricity well

• Metallic bonds are non-directional, so metals are malleable and ductile

• Strength of metallic bonds depends on the number of mobile electrons in the bond per atom

• Transition metals have mobile s and d electrons, so they are stronger and harder than alkali metals (only 1 s electron is mobile)

Metal Alloys

• Alloys are solid solutions of one or more metals

• Substitutional alloy: made by metals with atoms of similar size

• Interstitial alloy: made by metals with very different atomic sizes

• Adding nonmetals (such as carbon to iron) makes directional bonds

• Directional bonds make alloys harder, stronger and more brittle

Covalent Bonds

• Nonmetals of similar electronegativity cannot form ionic bonds

• These atoms share electrons to complete their octet

• Shared electrons “count” for both atoms• Each atom’s nucleus attracts the other

atom’s electrons

Forming Covalent Bonds

H ClShared!Shared!

Single bond, 2 electronsSingle bond, 2 electrons

8 e-!2 e-!

Multiple Bonds

C OH

HNeeds 1e-,makes 1 bond

Needs 4 e-,makes 4 bonds

Needs 2 e-,Makes 2 bonds

2e-!

8e-!

8e-!

Sigma bondElectron densitybetween nuclei

Pi () bond electron density above and below nuclei

Double bond

Molecular Dot Structures

• Count electrons – all valence electrons must appear in final structure

• Follow octet rule• Remember how many bonds each

type of atom makes (one for each extra electron needed)

Polyatomic Ion Dot Structures

• Same as molecular dot structures, except electrons must be added or subtracted to account for ion charge

• Subtract electrons for + charge, add for – charge

• Make all structures as symmetrical as possible

Carbonate (CO3-2) Dot

StructureSymmetry!

C

O

OO

Count electrons!

6 + 6 + 6 + 4 + 2 = 24

-2

Molecular Substances

• Made of molecules, which are loosely held together – van der Waals or London Dispersion forces

• Tend to be liquids, gases or low melting solids

• Melting molecular solids involves separating molecules from each other

• Most are insulators

Formulas and Names of Small Molecules

• Many have common names (i.e. water, ammonia)

• Systematic names use prefixes for each element

• P2O5 – diphosphorus pentoxide• N2O – dinitrogen monoxide • “mono” is not used for the first

element in a compound

Formulas and Names of Small Molecules

• CO2 – carbon dioxide

• CO – carbon monoxide

• SO3 – sulfur trioxide

• CCl4 – carbon tetrachloride