Ionic Compounds Ionic Compounds Ionic Compounds Ionic Compounds Ch.6 & 7.
Ionic Compounds
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Transcript of Ionic Compounds
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