*
* What is:
1. A metal…
2. A non metal…
3. Molecular compound…
4. Ionic compound…
5. Group…
6. Period…
7. Charge…
*
H2(g)
hydrogen
N2(g)
nitrogen
O2(g)
oxygen
F2(g)
fluorine
Cl2(g)
chlorine
Br2(l)
bromine
I2(g)
iodine
S8(s)
sulfur
P4(s)
phosphorus
*
Are now found in the data booklet!!!!
*
1.
2.
3.
4.
write the first name as given on the periodic table of
elements.
write the last name using an “ide” ending.
place the appropriate prefix in front the first and last
name to describe the number of atoms there are of
each element.
where the first element has only one atom, “mono” is
not necessary.
Example:
P4O3(g) =
4 atoms
tetraphosphorus trioxide
3 atoms
*
1.
2.
3.
4.
5.
Combining elements of the periodic table
that come from the metals and nonmetals
(left and right side of the “staircase” only)
forms an ionic compound.
When naming ionic compounds:
write the first name as given on the periodic
table of elements.
write the last name using an “ide” ending.
use no prefixes.
Example: CaCl2 - calcium chloride
*
1.
Using the crossover method:
1.
2.
3.
put down the metallic element first.
put down the nonmetallic element last.
cross the elements’ ionic charge to become the subscript for each
other element.
numerically simplify the subscripts.
4.
Example:
magnesium phosphide
Mg2+ and P3–
Mg3
Mg3P2 (s)
P2
***all ionic compounds are solids at room temperature.
* use the same format as above but whenever a complex ion is
named, use brackets to keep that complex ion as a group.
Example: sodium sulfate Na+ and SO42–
Put the two together grouping the complex ion: Na+ (SO42–)
Now cross the charges:
Na2(SO4)1
Since 1’s are not necessary:
Na2SO4 (s)
Example: calcium nitrate Example:
*
Ca2+ and NO3–
Ca2+(NO3–)
Ca(NO3)2 (s)
sodium hydroxide Na+ and OH–
Na+(OH–)
NaOH (s)
complex ions
*
* use the first ion listed as the most common. For example, Cu2+ is
more common than Cu+, so Cu2+ would be used if no choice is given.
When naming these compounds containing
elements with more than one charge:
• use Roman numerals to indicate the charge of the ion used.
Example:
CuCl is copper (I) chloride
Example:
CuCl2 is copper (II) chloride
* Follow ionic rules:
* Then use a dot along with the number of water
molecules required.
* When naming hydrated compounds, follow all
ionic rules described above.
* Then use a prefix in front of the word
“hydrate”.
Example:
Example:
*
CuSO4  6H2O is:
copper (II) sulfate hexahydrate
aluminum chloride trihydrate is:
AlCl3  3H2O
*
*
*
*
*
Formation of a precipitate
Formation of a gas
Colour change
Energy change
Simple Composition
element + element  compound
2 Na(s) + Br2(g)  2 NaBr(s)
*
Simple Decomposition
compound  element + element + element
2CaCO3(s)  2Ca(s) + 2C(s) + 3O2(g)
Single Replacement
element + compound  element + compound
Mg(s) + 2 NaOH(aq)  2 Na(s) + Mg(OH)2(aq)
Double Replacement
compound + compound  compound + compound
3 HCl(aq) + Al(OH)3(aq)  AlCl3(s) + 3 HOH(l)
Hydrocarbon Combustion
hydrocarbon + oxygen  carbon dioxide + water vapour
C3H8(g) + 5O2(g)  3CO2(s) + 4 HOH(g)
Oxidation
Metal + oxygen  Compound
2Fe(s) + O2(g) 
FeO (s)
Ionic compounds
Molecular compounds
may be soluble (turn into
aq)
may be soluble (turn into
aq)
DO form ions (individual
do not form ions (individual
charged elements in water) charged elements in water)
Are electrolytes (because
ions are electrolytes
Are not electrolytes
Examples:
NaOH(aq)  Na+(aq) + OH–(aq)
Al2(SO4)3(aq)  2 Al3+(aq) + 3 SO42–(aq)
*
*
Example: A silver nitrate solution reacts with a solution of barium chloride.
AgNO3(aq)
+
BaCl2(aq)

Ba(NO3)2(aq)
+ AgCl(s) (unbalanced)
Nonionic Equation: (regular balanced equation)
2 AgNO3(aq)
+
BaCl2(aq)

Ba(NO3)2(aq)
+
2 AgCl(s)
Total Ionic Equation: (list dissociations for electrolytes only)
2 Ag+(aq) + 2 NO3–(aq) + Ba2+(aq) + 2 Cl–(aq)  Ba2+(aq) + 2 NO3–(aq)
+
2 AgCl(s)
(do not write dissociations for solids, liquids or gases)
Net Ionic Equation: (list only what reacts or changes)
or, simplified:
2 Ag+(aq) + 2 Cl–(aq)  2 AgCl(s)
Ag+(aq) + Cl–(aq)  AgCl(s)
Do page 2 of your workbook
*
*
*
*
* All numbers listed are significant except zeros
before or after a decimal that must be used as
placeholders.
Example:
100.0010 - 7 significant digits
0.001010 - 4 significant digits
*
* Multiplication or Division Rules:
* Count the number of digits in each number being
multiplied or divided.
* Perform the multiplication or division.
* Round off to the least number of digits found in
each of the individual numbers being multiplied.
*
*
1.Determine the balanced chemical
equation.
2.Determine information given.
3.Determine your wanted.
4.Determine the number of moles of what
is given.
5.Use a wanted over given to determine
the number of moles of the unknown.
6.Solve for the answer.
Example: If 200 mL of 0.100 mol/L silver nitrate solution reacts
with a piece of copper, determine the mass of metal reacted.
2 AgNO3 (aq)
+
Cu(NO3)2(aq)
+
v = 0.200 L
Cu(s)

2 Ag(s)
m=?
C = 0.100 mol/L
nAgNO3 = Cv
nAgNO3= (0.200 mol/L)(0.100 L)
nAgNO3 =
0.0200 mol
nCu = 0.0200 mol x ½ = 0.0100 mol
mCu = nM
mCu = (0.0100 mol)(63.55 g/mol)
mCu =
0.636 g
Do page 3 of workbook!!