Chapters 7-9: Chemical Composition
Intra-chemical Forces
Intra = within
Atoms (elements) held together by an
attractive force
Metallic Bonding
atoms of a metal “share” the valence
electrons because they move from one
element to another
Ionic Bonding
valence electrons
are transferred
between two
elements
strongest bonds
Covalent Bonding
valence electrons
are shared
between two
elements
Weaker bonds
than ionic bonds
Ionic Bonding
Electrons are transferred from one
element to another.
Ionic Bonding
Opposite charges is
attractive force
Commonly referred to as
“salts”
Atoms that donates
electron = cation
Atom that accepts
electron = anion
Oxidation state refers to
the charge of an atom
Lewis Dot Formulas
Octet Rule: every element wants 8 electrons in
its outer shell.
a) potassium + chlorine → potassium chloride
a) magnesium + fluorine → magnesium fluoride
Types of Ions
monatomic cation: cation with one element
K+ Mg2+ Fe3+ Fe2+ Mn7+ Au3+ Au+
monatomic anion: anion with one element
name ends in – ide
Cl– O2–
N3–
S2–
F–
P3–
Br –
polyatomic ion: many atoms covalently bonded
that have a net charge.
NO3– SO42–
C2H3O2–
PO43–
NH4+
Writing Ionic Chemical Formulas
1.
2.
3.
4.
5.
6.
Composition
number of elements
Writing chemical formulas (from the names)
Recognize the (+) and (–) ions
Write the symbols of the elements with their charge
A Roman numeral will tell you what the charge is on the
cation if there is more than one possibility
7. Adjust the number of each ion (with subscripts) as needed
so the positive charge is equal and opposite the negative
charge.
8. If the ions are polyatomic and there is more than one, the
ion is enclosed with parentheses with a subscript on the
outside.
Writing Ionic Chemical Formulas
1. sodium chloride
2. calcium sulfide
3. calcium sulfate
4. barium phosphate
Naming Ionic Compounds
a)
b)
c)
d)
Consists of two words:
Name the cation
Name the anion
If the cation has more than one possible charge,
a Roman Numeral is used to show the charge.
e) All transition metals need roman numerals
except:
i. Zinc always has a charge of +2
ii. Silver always has a charge of +1
Naming Ionic Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
FeCl3
Fe3+  iron(III) chloride
FeCl2
Fe2+  iron(II) chloride
NH4Cl
Cu2SO4
NaC2H3O2
Ca(NO3)2
Zn(ClO)2
Cu2O
CuO
Naming Ionic Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
FeCl3
Fe3+  iron(III) chloride
FeCl2
Fe2+  iron(II) chloride
NH4Cl
ammonium chloride
Cu2SO4
copper (I) sulfate
NaC2H3O2
sodium acetate
Ca(NO3)2
calcium nitrate
Zn(ClO)2
zinc hypochlorite
Cu2O
copper (I) oxide
CuO
copper (II) oxide
Covalent Bonding
A.Valence electrons are shared
between two elements
B.Weaker than ionic bonding
polar & nonpolar covalent bonds
Polar Covalent (stronger): unequal sharing
of electrons (the more electronegative
element pulls more)
Nonpolar Covalent (weaker): equal sharing
of electrons
Writing Formulas
for Covalent Compounds
1. carbon dioxide
2. carbon monoxide
3. dinitrogen monoxide
4. carbon tetrafluoride
5. triphosphorus pentachloride
Naming Formulas
for Covalent Compounds
Binary covalent compounds (2 elements)
Formulas with two nonmetals
Rules:
i. First word:
1. prefix indicating the number of atoms for the first
element (if there is more than one)
2. name of first element
ii. Second word:
1. prefix for the number of atoms of the second
element (prefixes on supplement notes sheet)
2. name of second element
3. suffix –ide
Naming Formulas
for Covalent Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
NO
NO2
CBr4
P4O10
BF3
SiI5
H2O
S6Cl8
Se7O9
Naming Formulas
for Covalent Compounds
1.
2.
3.
4.
5.
6.
7.
8.
9.
NO
NO2
CBr4
P4O10
BF3
SiI5
H2O
S6Cl8
Se7O9
nitrogen monoxide
nitrogen dioxide
carbon tetrabromide
tetraphosphorus decoxide
boron trifluoride
silicon pentaiodide
dihydrogen monoxide
hexasulfur octochloride
heptaselenium nonoxide
Lewis Structures
The number of covalent bonds formed by an
atom equals the number of unpaired
electrons in the Lewis Dot Formula.
i. water (H2O)
Lewis Structures
ii. Hydrogen gas (H2)
iii. Hydrochloric acid (HCl)
Lewis Structures
iv. ammonia (NH3)
v. methane (CH4)
Multiple Bonds
i. double bonds: two pairs of
electrons shared O2
ii. triple bonds: three pairs of
electrons shared N2
Hybridization
Combining of two or more orbitals of nearly the
same energy into new orbitals of equal energy
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 2 (IIA):
Beryllium: [He]2s2
sp hybrid
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 13 (IIIA):
Boron: [He]2s22p1
sp2 hybrid
Hybridization
Most common hybridizations occur in groups 2, 13, 14 (IIA, IIIA, IVA)
Group 14 (IVA):
Carbon: [He]2s22p2
sp3 hybrid
Molecular Polarity
Molecules with more
than one element (polar
or nonpolar)
depends on:
i. electronegativity
difference (2 elements)
ii. Non-bonded electron
pairs (2+ elements)
iii. Structure (symmetry)
(2+ elements)
“Inter-chemical” forces
A. Inter = between
B. Whole salts or molecules attract
and bond with one another
“Inter-chemical” forces
1.
2.
Ion – dipole
3.
Hydrogen Bonding 4.
Dipole – Dipole
London Dispersion
Ion – Dipole forces
Strongest
inter-chemical
force
hydrogen bonding
hydrogen bonding is a unique case of
dipole – dipole bonding
occurs because
hydrogen’s
exposed
proton results
in a slight
positive
charge.
hydrogen bonding
medium strength inter-chemical bond.
occurs in molecules when hydrogen is
bonded with F, O, or N.
hydrogen bonding
hydrogen
bonding is
responsible for:
water’s high
boiling point,
and the low
density of ice
dipole – dipole bonding
weaker than hydrogen bonding.
occurs between polar molecules
London Dispersion Forces
named after Fritz London
the weakest
inter–molecular force
the random movement of electrons can create an
instantaneous dipoles