Honors text: Ch 10 & 12
Unit 06
Notes: Balancing Chemical Equations
Effects of chemical reactions:
 Chemical reactions rearrange atoms in the reactants to form new products.
 The identities and properties of the products are completely different from that of the reactants.
 Production of gases and color changes are signs of chemical reactions.
Energy and Reactions
Energy must be ________________ to _________________ bonds.
Energy is ________________ when bonds are ________________.
Chemical energy is ________________ in chemical reactions.
EXOTHERMIC REACTIONS: release energy (More energy is released as the products form bonds than is
absorbed to break the bonds in the reactants.)
ENDOTHERMIC REACTIONS:
Chemical equations are used to represent or describe chemical reactions. An equation shows:





The “+” means “reacts with”
The “” means “yields” or “reacts to produce”
To show physical states of each substance:
• Consider the reaction of iron with oxygen to form iron (III) oxide, or rust.
COEFFICIENTS: numbers in front of compound that represents the number of molecules of that compound
SUBSCRIPTS: small numbers that help define the compound.
Ex:
H2O:
2H2O:
H2O2:
• During a chem. rxn.; atoms are rearranged (NOT created or destroyed!)
• Chemical equations must be balanced to show the relative amounts of all substances.
• Balanced means: each side of the equations has the same # of atoms of each element.
Honors text: Ch 10 & 12
Unit 06
RULES to follow in balancing:
1.
2.
3.
4.
5.
6.
___Fe +
Correct formulas for all reactants & products.
Reactants  Products
Count the # of atoms of each element in reactants & products.
Balance one at a time using coefficients.
Check for balance.
Are the coefficients in the lowest possible ratio?
____O2 
____Fe2O3
Examples:

a)
CuCl2(aq) + Al(s)
b)
Propane, C3H8, burns in oxygen, O2, to form carbon dioxide and water.
c)
Pentane, C5H12, burns in oxygen, O2, to form carbon dioxide and water.
d)
Silver nitrate reacts with copper to produce silver and copper nitrate.
e)
reacts with oxygen, O2, to produce disphosphorus pentoxide.
f)
C7H14
+
O2
Cu(s)

+
AlCl3(aq)
CO2
+
H2O
Net Ionic Equations – Reactions in Aqueous Solutions
(solutions where water is the solvent)
When a soluble substance is dissolved in water, the substance often breaks into ions. This solution is said to
be an aqueous solution.
An aqueous solution is
.
OVERALL CHEMICAL REACTION:
Consider the reaction…
TOTAL IONIC EQUATION:
What is really going on is…
Note that the Na+ ion and the NO3- ion are not reacting. They are said to be spectator ions.
NET IONIC EQUATION: It is often useful to write an equation showing only the species that are actually reacting. This is called a
net ionic equation. It does not show the spectator ions.
Honors text: Ch 10 & 12
Unit 06
Precipitates:
To predict precipitates, use the Precipitation Diagram (see Table 1 below)
Table 1: Precipitation Diagram - Generalized Solubility of Salts in Water
Type of cation
NO3-
ClBrI-
SO42-
OH-
BaSO4
CaSO4
Mg(OH)2
S2CO32CrO42PO43-
(labeled “X”
below)
Group 1 cations
(Li+, Na+, K+, Cs+, Rb+)
and NH4+
Group 2 cations
(Mg2+, Ca2+, Ba2+)
Transition metal cations
AgX
PbX2
Hg2Cl2
Ag2SO4
PbSO4
Hg2SO4
Choose the cation row and read across to the anion column. If the block is colored, a precipitate will form from dilute solution.
Where a formula is given, that is the only cation-anion combination in that block that will precipitate.
For a more complete list of solubility rules, visit http://www.chemteam.info/Equations/Solubility-Table.html
Examples: For each, write the overall reaction, the total ionic equation and the net ionic equation. Circle
spectator ions.
NaOH(aq) + Cu(NO3)2(aq) 
Na3PO4(aq) + Fe(NO3)3(aq) 
MgSO4(aq) + AgNO3(aq) 
KOH(aq) + NaNO3(aq) 
Honors text: Ch 10 & 12
Unit 06
Types of Chemical Reactions
In chemistry, there are 5 general types of reactions:
1) Synthesis or Combination: 2 or more reactants combine to form 1 product.
2) Decomposition: 1 reactant decomposes to form 2 or more products.
3) Single Replacement: One metal replaces another metal in an ionic compound, producing a new ionic
compound and a metal.
4) Double Replacement: Two positive ions “switch places” forming 2 new ionic compounds:
5) Combustion: a hydrocarbon (containing C and H) or other substance burns in the presence of oxygen gas
(O2) to produce CO2 and H2O.
Honors text: Ch 10 & 12
Unit 06
Chemical Reaction Reference Tables
Table 2: Activity Series (more reactive elements are capable of replacing less reactive elements in a compound)
DECREASING ACTIVITY
ACTIVITY SERIES OF
METALS
Name
Symbol
Lithium
Li
Potassium
K
Calcium
Ca
Sodium
Na
Magnesium
Mg
Aluminum
Al
Zinc
Zn
Iron
Fe
Lead
Pb
(Hydrogen)
(H)*
Copper
Cu
Mercury
Hg
Silver
Ag
*Metals from Li to Na will replace H from acids and water; from Mg to Pb they will replace H from acids only.
Table 3: Keys to identifying general types of chemical reactions
General Equation
R + S  RS
Reactants
Two elements
Probable Rxn Type
Combination
Probable Products
A single compound
Two compounds (or at least one is a
molecular compound)
Combination
A single compound
A single binary compound
Decomposition
Two elements
A single tertiary compound
Decomposition
2 or more elements and/or compounds
T + RS  TS + R
An element and a compound
Single replacement
A different element and a new compound
R+S- + T+U-  R+U- + T+S-
Two ionic compounds
Two new compounds
CxHy + #O2  CO2 + H2O
Oxygen and a compound of C, H,(O)
Double
Replacement
Combustion
RS  R +
S
CO2 and H2O (with incomplete
combustion, C and/or CO may be
additional products)
STOICHIOMETRY
Calculations Based on Balanced Chemical Equations
Iron (III) oxide reacts with carbon monoxide to form iron and carbon dioxide.
How many CO molecules are required to react with 25 particles of Fe2O3?
How many iron atoms can be produced by the reaction of 2.5 x 105 particles of Fe2O3 with excess CO?
Honors text: Ch 10 & 12
Unit 06
*Formulas can also represent MOLES of substances involved in chem. rxns.
*Equations define reaction ratios, i.e. the mole ratios of reactants and products.
What mass of CO is required to react with 146 grams of iron (III) oxide?
What mass of iron (III) oxide is required to produce 8.65 g of carbon dioxide?
Limiting Reactants (Reagents)
Calculations need to be based on the limiting reactant, or the reactant that is present in lesser amount.
Example: Suppose a box contains 87 bolts, 110 washers, and 99 nails. How many sets consisting of one bolt,
two washers and one nail, can you construct from the contents of the box?
Example: What is the maximum mass of sulfur dioxide that can be produced by the reaction of 95.6 g carbon
disulfide with 100 g oxygen?
CS2
+
O2

CO2
+
SO2
Example: What mass of CO2 could be formed by the reaction of 8.0 g CH4 with 48 g O2?
CH4
+
O2  CO2
+
H2O
Honors text: Ch 10 & 12
Unit 06
Percent Yield
Many chem. rxns do not go to completion (reactants are not completely converted to products)
Percent yield indicates how much of a desired product is obtained.
Percent yield may be represented as:
% Yield =
actual yield
theoretical yield
x 100
So far the masses we have calculated from chemical equations were based on the assumption that each
reaction occurred 100%.
• The theoretical yield is the yield calculated assuming 100% reaction and isolation of 100% of the desired
product.
• The actual yield is the yield actually obtained form a particular “run” of a rxn.
Example: A 10.0 g sample of ethanol, C2H5OH, was boiled with excess acetic acid, CH3COOH, to produce 14.8 g
of ethyl acetate, CH3COOC2H5. What percent yield of ethyl acetate is this?
CH3COOH + C2H5OH

CH3COOC2H5 + H2O