Chapter 7 (Chemical Formulas and Bonding)
INTRODUCTION
Question: Why do atoms of different elements react to form compounds?
(Increased stability due to completion of octet of valence electrons. The Octet Rule
states that atoms tend to gain, lose or share electrons in order to acquire a full set of
valence electrons, generally 8.)
Question: What is happening in this process? (Electrons are transferred from
one atom to another or they are shared between atoms; ionic, covalent and polar
covalent bonds form.)
Lewis Dot Diagrams
To help us understand and draw atoms quickly, we will use the Lewis Dot Diagrams, first
introduced by Gilbert N. Lewis in 1916. This method shows only the valence electrons,
‘where the action is’ in chemistry.
Text page 230 and worksheet on Lewis Dot Diagrams. (Noble gas core + valence
electrons are shown in Fig. 7-9, p230.)
Transparency of Lewis Dot Diagrams (Have students work them out).
Four basic types of Chemical Bonds:
-Ionic Bonds: electrons are transferred from one atom to another. (NaCl)
-Covalent Bonds: electrons are shared between two atoms. (H-H or Cl-Cl)
-Polar Covalent Bonds: electrons are shared unequally between two atoms.
-Metallic bonds: metal atoms in a ‘sea of electrons’ that are shared.
Properties of a material depends on the type of bonding:
Ionic: crystalline, water soluble, conduct electricity in water solution.
Covalent: often liquids/gases, insoluble in water, do not conduct electricity.
However, there are exceptions since some materials have properties of both ionic
and covalent substances. (Alcohol, glycerine, vitamin C).
Properties are related to how the electrons are shared.
-Bonds are a ‘tug of war’ between atoms for the electrons. A measure of this ‘tug
of war’ is ELECTRONEGATIVITY (from Ch 5) since it is a measure of the ability of an
atom in a bond to attract electrons.
Electronegativity Examples (Fig. 5-23 on page 184 for EN values; also pg 241):
F: highest electronegativity, EN = 4.0 (wants one electron)
Fr: low EN = 0.7 (wants to give up one electron)
C: mid range EN = 2.5 (just wants to share electrons)
Remember, Electronegativity increases in a Period (left to right) on Periodic Table (Na to
Cl); EN decreases going down a Group (Li vs. Cs); EN for Noble Gases = zero (octet!).
(Use transparency.)
General Rule: The farther apart two atoms are on the Periodic Table, the greater their
difference in electronegativity. Use Delta EN to predict bonding. (p242 for Fig 7-21)
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Delta EN is just the difference between the largest and smallest value, so it is always a
positive number:
Delta EN > or = 2.0
Delta EN = 0.4 to 2.0
Delta EN < or = 0.4
Ionic Bond (electron transfer)
Polar Covalent Bond (electron sharing unequally)
Nonpolar Covalent Bond (‘pure’ covalent; equal sharing)
EXAMPLES:
1). Cs (0.7) & F (4.0), so Delta EN = 3.3 (IONIC BOND)
2.) H (2.1) & H (2.1), so Delta EN = 0.0 (‘Pure’ COVALENT BOND)
3.) H (2.1) & C (2.5), so Delta EN = 0.4 (electrons NOT shared equally, so polar
covalent sharing)
4.) H (2.1) & Cl (3.0), so Delta EN = 0.9 (Polar covalent sharing, but Cl “gets the
biggest half”!)
Show drawings of NaCl (ionic), H-H (pure covalent) & HCl (polar covalent) with dipoles
and δ+/- of partial charges.
Ch.7-1 (Ionic Bonding)
Ionic Bond: a bond in which a positively charged ion (cation) is electrostatically
attracted to a negatively charged ion (anion) to form an ionic compound. (In the solid
state these ions are arranged in a crystal lattice or network. See transparency.)
Types of Ions: Monoatomic ( Na+1, Ca+2, Al+3, Cl-1, O-2, N-3 for example) and
Polyatomic (Fig 7-10 and 7-11) (Sulfate, phosphate, etc.).
Class Activity: Use magnetized cards to make balanced chemical equations of ionic
compounds.
Criss-Cross Method of balancing Ionic Formulas. Also Worksheet 7-1PP, IF sheet 38,
44, 48 & 45. for practice. (Note that it is not necessary to indicate a ‘1’.)
Naming Binary (two elements) Ionic Compounds:
-Name cation first, then anion (chlorine become chloride, -ide ending). Note
names of polyatomic anions and cations. Note use of Roman numerals for various ions
(copper (I) & (II); iron (II) & (III); lead (II) & (IV) for example.)
Empirical Formula: the chemical formula that gives the simplest whole-number ratio
of atoms of the elements in a compound (numbers written as subscripts).
NaCl, not Na2Cl2; Al2O3, not Al4O6; CaBr2, not Ca2Br4
Review Worksheet: 7-1R&R.
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Ch7-2 (Covalent Bonding)
Covalent Bond: a bond that is formed by the sharing of electrons between two atoms
(Worksheet IF 39)
Molecule: a group of atoms that are united by covalent bonds. H2O, CCl4
Molecular Substance: a substance that is made of molecules.
Molecular Formula: the chemical formula that indicates the numbers of each atom in a
molecular compound. (See above.)
Structural Formula: a formula that indicates how the atoms in a molecule are bonded to
each other. (Water, H2O is H-O-H, not H-H-O, for example.)
Use Lewis Dot Diagrams to draw covalent bonds and covalent molecules:
H. + H. -> H:H
Cl + Cl -> Cl-Cl
Show CH4, NH3
Single bonds (single covalent bonds): two atoms share ONE PAIR of electrons (H:H)
Double bonds: two atoms share TWO PAIRS of electrons (H2C::CH2 or H2C=CH2)
Triple bonds: two atoms share THREE PAIRS of electrons (HC:::CH)
(Use dots or dashes, the latter being simpler to convey the idea of two electrons in a
covalent bond.)
EXCEPTIONS TO THE OCTET RULE
-Boron compounds often have less than an octet (BF3 on page 240).
-Sulfur and phosphorus compounds sometimes have more than an octet (SF4 on page
241).
-Sometimes molecules have an odd number of electrons (nitrogen monoxide, NO).
Polarity of Molecules
Recall the discussion above about polar covalent bonds as a result of electronegativity
differences. Nonpolar bonds form between identical atoms (H-H, F-F) or atoms with
very similar EN. Polar bonds result when EN differences are 0.4 to 2.0. (Show
examples of polar molecular structures, H2O, CH3Cl etc. with dipole and delta charges.)
Worksheet 7-2 Apply(??) and 7-2RR
7-3 (Naming Chemical Compounds)
Ionic Compounds (See the section above.)
Hydrates: Special ionic compounds that have water associated with their crystal
structure that may be removed by simple heating to make them anhydrous. Formulas
are generally indicated by the water as (nH2O or .nH2O). Prefixes are used (mono-, ditri-, tetra-, penta- etc.) (See p246 for full list.)
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Covalent (Molecular) Compounds: Mono is generally not used, but the other
prefixes denote the numbers of atoms in the formula. There are lots of exceptions for
common substances (carbon monoxide, not carbon monoxide; NH3 is ammonia; H2O is
water.)
Acids: For now we will use the definition of a molecular substance that dissolves
in water to produce hydrogen ions (H+). Many acids, many names and many
exceptions, but you should know the common acids:
HF, HCl, CBr, HI
H2S
HNO3
H2CO3
H2SO4
H3PO4
HC2H3O2
Worksheets: IF page 46, 47, 40.
7-3PP and 7-3RR
PRACTICE, PRACTICE, PRACTICE WRITING FORMULAS AND NAMING
COMPOUNDS!
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