1. UNIT 2 : Bonding, Reactions and Stoichiometry

2. Nuclear reactions involve the nucleus
Chemical reactions involve valence electrons
Chemical bonds form when electrons
are attracted to 2 different atoms
When bonds form, energy is released
ex.) Ca + S  CaS + energy
When bonds break, energy is absorbed
ex.) CaS + energy  Ca + S

3. Atoms are most stable when their last PEL (electron shell) is full—
usually 8 valence electrons (2 for 1st energy level)
Like the noble gases
He Ne Ar Kr Xe Rn
Atoms can reach a stable octet by forming chemical bonds:
1. Atoms gain or lose electrons
ionic bond
2. Atoms share electrons
covalent bond

4. Ionic Bonds Na+ Cl-- Na 2 - 8 - 1 Cl 2 - 8 -7
-- form when electrons are transferred from one atom to another
Example: Na
+11 p
-10 e +1
Na+
sodium ion
more stable if 1 e- is lost
( 2 – 8 matches Ne) Cl
+17 p
-18 e -1
Cl--
chloride ion
more stable if 1 e- is gained
( 2 – 8 – 8 matches Ar)
Since + and – attract, the Na+ and Cl- bond
to form NaCl, sodium chloride
An electron has transferred from Na+ to Cl-

5. Dot Structures for Ionic Compounds:
Ionic formulas depend on electron gain/loss (look at oxidation state)
Examples: calcium ___ oxide ___ iodide ___ aluminum ___
Ca+2
O-2 I-
Al+3
The formula for an ionic compound is neutral
(+ = - )
Examples: Ca+2I-
Ca+2I-
CaI2 calcium iodide 2
Al+3O-2
Al+3O-2
Al2O3 aluminum oxide
Dot Structures for Ionic Compounds: x
Example: Ca x I x x I
[Ca]+2
[ ]- I
[ ]- I

6. Ionic Substances --high melting and boiling points -- hard and brittle
--poor electrical conductors
when solid
--good electrical conductors
(electrolytes)
when melted or dissolved –
they ionize: separate into + and - ions

7. Covalent Bonds -- when 2 atoms share electrons
Fluorine molecule -- F2
When 2 identical atoms share electrons, they share equally
--the result is a non-polar covalent bond
When 2 different atoms share electrons, they share unequally
--one of the atoms attracts the electrons more strongly
-- the “stronger” atom becomes slightly negative, the
“weaker” atom becomes slightly positive

8. The bond has 2 opposite (+ and -) ends, or poles
The result is a polar covalent bond
The tendency of an atom to attract electrons is measured by its
Electronegativity (see ref table S)
High electronegativity = strong attraction for e F = 4.0
Low electronegativity = weak attraction for e Fr = 0.7

9. Bonds can be classified by the difference in electronegativities
of the atoms
difference of 1.7 or more = ionic
difference of less than 1.7 = polar covalent
difference of 0 = nonpolar covalent
Atoms share pairs of e- They can share 2 or 3 pairs to form
double or triple bonds Ex., CO2, N2

10. Covalent bonds produce molecular substances
-- groups of atoms, or molecules, are separate from each other
-- soft
-- low melting and boiling points
-- poor conductors of heat and
electricity

11. Metals and Metallic Bonding
-- positive metal ions form a framework, surrounded by
mobile electrons
-- good conductors of heat and electricity
-- malleable (can be shaped by hammering; not brittle)
-- most have high melting points

12. } Summary of Bonds and Substances BOND SUBSTANCE Ionic Ionic
Polar Covalent
Molecular
Non-Polar Covalent
Metallic Metal

13. Intermolecular Forces – Between molecules
are important in molecular substances not ionic
are weaker than ionic, covalent, or metallic bonds
depend on the attraction between + and - charge
A molecule with a separation of + and – is called a
polar molecule, or a dipole
To be a dipole, a molecule must have:
1. polar covalent bonds
2. an asymmetric distribution of charge
dipole
not a dipole

14. Examples: H |
H—C—H
H—N—H | H
H—Cl
Cl—Cl
The + end of one dipole attracts the – end of another

15. Hydrogen Bonds -- a dipole force formed when polar molecules contain H
bonded to a small, electronegative atom: N, O, or F
-- are stronger than other dipole forces
Example: Boiling Point, ˚C
H2Te
H2Se
H2S
H2O
There are also forces that attract non-polar molecules to each other –
but these are weaker than dipole or hydrogen bonds

16. Strength of intermolecular forces:
1. H bonds
2. regular dipole
The intermolecular forces determine:
3. non-polar forces
1. melting and boiling points
(strong forces = high)
2. hardness
(strong forces = hard)
3. solubility -- “like dissolves like”
polar substances mix with polar substances
non-polar substances mix with non-polar substances
Example: water, a polar liquid, dissolves ionic solids like salt

17. NAMING COMPOUNDS

18. REMEMBER THESE LAWS…
In the 1700’s, scientists making accurate measurements
discovered several new laws --
1. Law of Conservation of Matter
--matter is not created or destroyed during chemical or physical changes
2. Law of Definite Composition
-- compounds have an unvarying
composition
3. Law of Multiple Proportions
-- elements combine in simple ratios

19. Then Amedeo Avogadro, Italy, 1800’s
at the same temp and pressure, equal volumes
of different gases have the same number
of molecules
Example: 1.0 L of H2 has the same number of
molecules as 1.0 L of CO2.
The CO2 weighs more because each CO2 molecule is
heavier (formula mass H2 = 2, CO2 = 44)
Since it’s difficult to work with single molecules, chemists usually
work with groups of them:
Avogadro’s number = 6.02 x 1023
6.02 x 1023 molecules = 1 mole
Why 6.02 x 1023?
--because 6.02 x 1023 molecules of any substance weigh the
formula mass of that substance, in grams

20. Examples: 1
6.02 x 1023 Na atoms = ______moles
= _______ grams 23
6.02 x 1023 H2O molecules = ______ moles 1
= _______ grams 18
3.01 x 1023 B atoms = _____ moles
0.5
= _______ grams
5.5
80 g calcium = _____ moles 2
12.04 x 1023
= _______________atoms
24.08 x 1023 4
48 g carbon = _____ moles
= ______________ atoms
3.55 x 1023
16 g aluminum = ______ moles
0.59
= _____________ atoms
grams
formula mass
moles =

21. By Avogadro’s law, 1 mole of any gas must have the same volume
The molar volume of a gas at STP = 22.4 L
Examples:
89.6 L
1. What volume would 4 moles NO2 occupy, at STP? __________
12.04 x 1023 2
L of hydrogen at STP = ____ moles
= ___________molecules
3. What is the mass of 22.4 L of O2 at STP? _________
32 g
4. a. What volume does 1 mole NO occupy at STP? _________
22.4 L
30 g
b. What is the mass of 1 mole of NO? _________
1.3 g/L
c. What is the density of NO? __________
mass
volume
density =
5. A gas has a density of 1.78 g/L at STP
a. what is the mass of 1 L of this gas? _________
1.78 g
b. how much does 22.4 L of this gas weigh? ________
39.9 g
c. what is the formula mass of this gas? ________
39.9 g

22. Types of Chemical Reactions
1. Synthesis -- A + B  AB
Ex.) Fe + O2  Fe2O3 4 3 2
Ex.) SO3 + H2O  H2SO4
2. Decomposition -- AB 
A + B
Ex.) H2O  H2 + O2 2 2
Ex.) CaCO3  CaO + CO2
Single Replacement -- A + BC 
AC + B
or BA + C
Ex.) Zn + NaCl  ZnCl Na 2 2
Ex.) Cl KBr  KCl + Br2 2 2

23. 4. Double Replacement -- AB + CD 
AD + CB
Ex.) MgI2 + K2S  MgS + KI 2
Ex.) Ba(NO3)2 + (NH4)2SO4  BaSO NH4NO3 2