Ionic Compounds
Ch.6 & 7
Ionic Bonding – strongest of
all bonds
• definiton – results from the electrical
attraction between cations (+) and
anions (-)
• “ions” means charges in our case
• Calculated by the difference between
electronegativity values (over 1.7)
• Figure 2 -- p.176
Ionic Bonding
• Positive ions are formed from the
metal elements when they lose
electrons
• Negative ions are formed from the
nonmetal elements when they gain
electrons
• The two, (+) and (-) come together to
form a new compound
Ionic Compounds
• definition – composed of positive and
negative ions that are combined so
that the numbers of positive and
negative charges are equal
• In a neutral compound, the net charge
will be zero
• Balance the charges from cations and
anions until the compound is neutral
Forming Ionic Compounds
1.
2.
3.
4.
5.
Na+ + Cl-  NaCl
Ca2+ + Cl-  CaCl2
Mg2+ + O2-  MgO
Fe3+ + S2-  ______
beryllium + iodine  ______
Naming Ionic Compounds
• Name the metal first (complete name)
• Name the root of the nonmetal the
change the ending to –ide
• If there is a polyatomic (more than one
atom carrying one charge) use the name
as it is
• Transition metals have a Roman numeral
after it to tell its charge
Naming Ionic Compounds
1.
2.
3.
4.
5.
NaCl
CaCl2
MgO
Fe2S3
BeI2
____________________
____________________
____________________
____________________
____________________
Ionic Compound
Characteristics
• A formula unit is the smallest whole
number ratio of cation to anion
ex. NaCl
• Ionic compounds mostly have a
crystalline structure
• Lattice energy is the energy released
when one mole of an ionic crystalline
compound is formed from gaseous ions
Ionic Compound
Characteristics
• Generally have a high melting point
and boiling point
• Hard but brittle
• Not electrical conductors as solids
• Most can dissolve in water
• Strongest bond
Polyatomic Ions
• Definition: a charged group of
covalently bonded atoms
• Lewis structures can be used to show
the electron placement and where
the charge of the ion comes from
(p.194)
Polyatomic Ions
Names have a system:
– ending -ate means the highest number
of oxygen bonded for that nonmetal
(sulfate, chlorate, iodate, etc)
– ending –ite means one less oxygen than
-ate (sulfite, chlorite, iodite, etc)
– Chlorine is a special case:
ClO4-, ClO3-, ClO2-, ClO-, Cl- (p.226)
Polyatomic Ions to Study
•
•
•
•
•
Phosphate
Chlorate
Nitrate
Sulfate
Carbonate
PO4-3
ClO3NO3SO3-2
CO3-2
Oxidation Numbers
Definition: the number of electrons
that must be added to or removed
from an atom in a combined state to
convert the atom into the elemental
form
Example: H+ + Cl-  HCl
KMnO4 K = ___ Mn = ___ O = ___
Metallic Bonding
• Definition: the chemical bonding that
results from the attraction between
two metal atoms and the surrounding
sea of electrons
• Properties include:
– High electrical and thermal conductivity
– Strong absorber and reflector of light
– Conforms to shape easily
• Malleability – hammered into sheets
• Ductility – drawn into wires
Covalent Molecules
Ch. 6 & 7
Covalent Bonding
• Definition: bond formed by the
sharing of electron pairs between two
nonmetals
• There are two types of covalent bonds:
– Nonpolar: electrons are shared equally
– Polar-covalent: unequally shared electrons
Electronegativities
• Every element on the PT has an
electronegativity value
• Those values are subtracted to see
whether is:
– Nonpolar
– Polar-covalent
– Ionic
(0-0.3)
(0.3-1.7)
(over 1.7)
Lewis Structures
• Definition – formulas in which atomic
symbols represent nuclei and innershell electrons, dot-pairs in covalent
bonds
• Sample Problem C - p.185
• Practice - p.186
Molecular Geometry
• VSEPR theory – repulsion between the sets
of valence electrons surrounding an atom
causes these sets to be oriented as far as
possible
• Use VSEPR with Lewis structures to come up
with shapes
• Shapes are:
– Linear, trigonal planar, tetrahedral, bent, and
trigonal pyramidal, etc. (p.200)
• Sample Problem F – p.201
Intermolecular Forces
• Definition: forces in between molecules
• Weaker than bond strength
• Types:
– Dipole-dipole: strongest in polar covalent
– Hydrogen bonding: hydrogen bonded to
highly electronegative atom
– London dispersion forces: weakest force due
to motion of atoms in compound
Naming Covalent Molecules
• Use prefixes listed on p.228
• Prefixes tell how many atoms of each nonmetal
make up the molecule
• The nonmetal farthest left on PT is written
first, then the most electronegative atom
(farthest right on the PT)
• Mono does not appear on the first atom to
notate one, it is understood
• Omit vowels on the prefix if there is a vowel on
the element
• Second atom ends in -ide
Covalent Prefixes
mono
di
tri
tetra
penta
1
2
3
4
5
hexa
hepta
octa
nona
deca
6
7
8
9
10
Two Rules of Thumb:
1. If two vowels end up next to each other, the vowel on the
prefix will be deleted.
2. Mono is never used for the first element.
Practice
Name the following covalent compounds:
1. CO2
_________________
2. P2O5
_________________
3. OF3
_________________
4. SO2
_________________
Acids
• Formed from H+ + a nonmetal or a
polyatomic
• If H+ is bonded to a nonmetal, use the
prefix hydro- + the root of the nonmetal +
ending with –ic acid
• If H+ is bonded to a polyatomic, do NOT
use the prefix hydro! Use the root of the
polyatomic and use ending from –ate to –ic
or –ite to –ous acid
• Section Review p.231 #4 f-h
Practice
Name the following acids:
1. HCl
________________
2. HClO4
________________
3. H2SO3
________________
4. HI
________________