Pre-AP Chemistry
Covalent Bonding Review Sheet
 PART 1: COVALENT BONDING THEORY
- Chemical Bond: force holding two atoms together.
- Atoms share valence electrons to follow the octet rule (=8 valence electrons). This is a covalent
bond which is an intramolecular, electrostatic force.
- Covalent bonding occurs between two nonmetals to form molecules.
- Covalent compounds have relatively low melting and boiling points.
- A covalent bond can be a single bond (two e-), double bond (4 e-), or triple bonds (6 e-).
o A single bond is the longest and weakest covalent bond.
o A triple bond is the shortest and strongest covalent bond.
- In nature, there are seven diatomic molecules: H2, N2, O2, F2, Cl2, Br2, and I2.
o O2 contains a double bond, N2 contains a triple bond, and all other diatomics are single
bonded.
o The elements cannot be found in nature as single atoms.
- The type of bond can be determined using electronegativity (EN) – an element’s ability to
compete for electron aka the strength to pull electrons.
o Nonpolar covalent – equal sharing of electron – EN difference = 0
o Polar covalent – unequal sharing of electron – EN difference = between 0 and 2.
o Ionic – transfer of electron forming cation and anion – EN difference = 2 or greater.
- Metallic bond – between metals; nuclei embedded in a sea of electrons.
 PART 2: LEWIS STRUCTURES AND VSEPR
- Lewis Structures show the bonding of atoms in a molecule.
- VSEPR (Valence Shell Electron Pair Repulsion) Theory helps predict molecular shape around the
central atom based upon the theory that electron/negative groups position themselves as far
apart as possible.
- Steps to drawing Lewis Structures (L.S.):
o Count total valence electrons from each atom in formula (be sure to account for charge)
o Write the central atom. Hydrogen is NEVER the center.
o Write all other atoms surround central atom.
o Connects outer atoms with central atom using only single bonds. Subtract the number
of electrons used from total.
o Fill octets of surrounding atoms. (Note: H will only have 2 e- in octet.) Subtract the
number of electrons added from total.
o Place any remaining electrons on the central atom.
o Push electrons to form multiple bonds to ensure that all outside atoms have a full octet.
(Note: Carbon and Silicon will need to have an octet when they are the central atom.
Also, Period 3 and heavier elements can have expanded octets when in the center.)
o The L.S. of a charged molecule must be enclosed in brackets with the charge noted.
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VSEPR Theory
o Must be able to draw a correct Lewis Structure to determine VSEPR Notation.
o VSEPR Notation is AXnEm.
o See the handout/chart from class for more details and explanations.
o You are responsible for knowing the VSPER notation and electronic geometry.
 PART 3: HYBRIDIZATION AND FORMAL CHARGE
- Know the basic concepts of hybridization as presented in class. Don’t worry about the details.
- The hybridization of an atom can be determined using the finger method.
o The orbital or “fingers” : s p p p d d d d d
o Using your fingers, count the number of bonding events around the atom. (Note:
Multiple bonds count as one bonding event and lone pairs count as one bonding event.)
o This number will correspond to the orbital chart and the hybridization.
o Example: An oxygen atom has two single bonds and 2 lone pair (as in the water
molecule). This is 4 bonding events so we count four orbital: s p p p = sp3 hybridized.
- Each bond is an overlap of two hybridized orbitals. These orbitals can be determined using the
hybridization of the bonded atoms.
- Each bond contains a sigma (σ) bond and any other bonds are pi (π) bonds.
o A single bond has one sigma bond.
o A double bond has one sigma and one pi bond.
o A triple bond has one sigma and two pi bonds.
- While a molecule can have an overall charge, charges can vary within different parts of the
molecule.
- Formal charge is the charge of specific atoms within a molecule. This assumes that all electrons
are shared equally in each bond.
- The sum of the formal charges within a molecule must add up to the overall charge of the entire
molecule.
- The formal charge for an atom can be determined using a finger method.
o Determine the number of valence electrons in a neutral atom of the element of interest.
o Subtract the number bonds.
o Subtract the number of lone electrons on that atom (lone pair = subtract 2 electrons)
o The remaining value (positive or negative) will be the formal charge of the atom.
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 PART 4: NOMENCLATURE
Type I: Ionic Bonding with FIXED oxidation states (charge)
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How to identify: Metal is in the main groups (Group 1, 2, 3) and Ag+, Zn2+, Cd2+
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SYMBOL  NAME
1) Write full name of the metal/cation.
2) Write name of the nonmental/anion with –ide ending.
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NAME  SYMBOL
1) Write symbol for the metal/cation.
2) Write symbol for the nonmetal/anion.
3) Use the charge from the periodic table (based on element’s group #) to determine the
subscripts of the metal and nonmetal to give an overall charge of zero (pos = neg).
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Examples:
1) Na2S
2) Calcium fluoride
sodium sulfide
CaF2
Type II: Ionic Bonding with VARIABLE oxidation states (charge)
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How to identify: Metal is a transition metal except Ag, Zn, Cd
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SYMBOL  NAME
1) Write full name of the metal/cation.
2) Write name of the nonmental/anion with –ide ending.
3) Use the fixed charge of the anion (from the periodic table) to determine the total
positive charge to give an overall compound charge of zero.
4) Divide the overall positive charge by the number of metal atoms to determine the
charge of EACH metal atom. Write this number as a Roman Numeral between the metal
and nonmetal in the chemical name.
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NAME  SYMBOL
1) Write symbol for the metal/cation.
2) Write symbol for the nonmetal/anion.
3) The charge of the EACH metal cation comes from the Roman Numeral in the given
name.
4) The charge of the anion comes from the periodic table.
5) Determine the subscripts of the metal and nonmetal to give an overall charge of zero
(pos = neg).
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Examples:
1) Sn3P4
2) Cobalt(II) nitride
tin(IV) phosphide
Co3N2
Type III: Ionic Bonding with Polyatomic Ions
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How to identify: A polyatomic ion will be present in the name or formula.
The list of polyatomic ions on Page 257 (textbook) MUST be memorized (name, formula, charge)
Consider the polyatomic ion as one large unit that does not change
NEVER touch or change the polyatomic ion --- use parentheses to denote multiple quantities
The charge of a polyatomic ion is the charge for the WHOLE unit/group of atoms. The sulfate ion
(SO42-) has a charge of 2- on the entire sulfate molecule…one SO4 has a charge of 2-.
Do NOT change the ending of the polyatomic anion to an –ide ending. Write the full name of the
ion.
1) If the metal/cation has a fixed charge, follow rules for Type I (except do not use –ide
ending).
2) If metal/cation has a variable charge, follow rules for Type II (except do not use –ide
ending.)
Examples:
1) Mn3(PO4)7
2) Barium acetate
permanganate(VII) phosphate
Ba(C2H3O2)2
Type IV: Covalent Bonding
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How to identify: 1st element is a nonmetal.
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SYMBOL  NAME
1) Write full name of the 1st nonmetal.
2) Write name of the 2nd nonmetal with –ide ending.
3) Put Greek prefix before both nonmetal based on the number of each nonmetal atom.
4) Remove “mono-” if placed in front of the first nonmetal. NEVER put “mono-” at
beginning of the chemical name.
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NAME  SYMBOL
1) Write symbol for the 1st nonmetal.
2) Write symbol for the 2nd nonmetal.
3) Use prefix before each metal to give the subscript of the corresponding nonmetal.
Greek Prefixes
1 = mono- 6 = hexaExamples:
2 = di7 = hepta1) BrO3
bromine trioxide
3 = tri8 = octa2) Diphosphorus pentoxide
P2O5
4 = tetra9 = nona5 = penta- 10 = decaPage 4 of 4
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