Transcript 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
+
Cl
11p+
11e-
17p+
10e-
18e-
17e-
-
Making More Ionic Bonds
+2
Mg
12 p+
O
12e-
8p+
8e-
10e-
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
11 p+
+1
11e10e-
Na
p+
Na
11
11e-
10e-
+1
S
-2
-1
16p+
16e
Formula: Na2S
18e
Name: sodium sulfide
Total + charge = 2, Total - charge = 2
Total charge overall = 1 + 1 + (-2) = 0
-
Unequal charges
Determine the formula of
calcium bromide.
+2
+1
Ca
Formula:
CaBr2
Br
-1
-1
Br
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 decahydrate
CuSO4.5H2O
Copper (II) sulfate pentahydrate
Sodium 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
Single bond, 2 electrons
2 e-!
H
Shared!
Cl
8 e-!
Pi (p) bond
electron
density above and
below
8e-! nuclei
Multiple Bonds
Double bond
Needs 1e-,
makes 1 bond
2e-!
H
H
8e-!
C
O
Needs 4 e-,
makes 4 bonds
Needs 2 e-,
Makes 2 bonds
Sigma (s) bond
Electron density
between nuclei
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 Structure
Symmetry!
O
Count electrons!
6 + 6 + 6 + 4 + 2 = 24
C
O
O
-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