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 ADDED to BREAK
bonds.
Energy is RELEASED when bonds are
FORMED.
Chemical energy is CONSERVED in
chemical reactions.
Exo- vs. Endo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: absorb
energy
Chemical Equations
Chemical equations are used to represent or
describe chemical reactions.
For example when hydrogen H2 burns, it
reacts with oxygen, O2, in the air to form
water. We write the chemical equation for
this reaction as follows:
2H2 + O2 —> 2H2O
Chemical Equations
An equation shows…
Formulas of reactants
Formulas of products
Molar ratios of all compounds in the reaction.
Chemical Equations
We read the (+) sign as “reacts with” and the
arrow (—>) as “produces” or “yields”.
2H2 + O2 —> 2H2O
Reactants
Products
To show physical states of each
substance:
• (s) or 
• (l)
• (g) or 
• (aq)
-
solid
liquid
gas
aqueous
• aqueous means dissolved in water
To show physical states of each
substance:
• Consider the reaction of iron with oxygen
to form iron (III) oxide, or rust.
• Fe(s) + O2(g)  Fe2O3(s)
(unbalanced)
Coefficients & Subscripts
COEFFICIENTS: numbers in front of
compound that represents the number of
molecules/moles of that compound
SUBSCRIPTS: small numbers that help
define the compound.
Coefficient
2H2SO4
Subscript
H2O
One molecule of water
2H2O Two molecules of water
H2O2 One molecule of Hydrogen Peroxide
•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.
CH4 + O2 —> H2O + CO2
Unbalanced
CH4 + 2O2 —> 2H2O + CO2
Balanced
In order to balance…
• Write correct formulas for all reactants
and products
• Reactants  Products
• Count the number 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?
Balancing Equations
NOTE: When balancing equations, you
may change coefficients as much as
you need to, but you may never
change subscripts because you can’t
change what substances are involved.
Balancing equations involves a great deal of
“trial and error” at first,
but there are some tricks…
For example…..
Sodium metal reacts with water to produce
sodium hydroxide and hydrogen gas.
Na + H2O —> NaOH + H2
Note that on the product side (right side) there are
an odd number of hydrogens. On the reactant side
(left side) there is an even number. This implies
there must be an even coefficient in front of the
NaOH. Lets start with 2
_Na + _H2O —> 2NaOH + _H2
Now lets balance sodium; we need a 2 in front of the Na…
2Na + _H2O —> 2NaOH + _H2
Now consider hydrogen…
2Na + 2H2O —> 2NaOH + H2
2Na + 2H2O —> 2NaOH + H2
Check to see if it balances…
2 Na on the left
2 Na on the right
4 hydrogen
2 + 2 = 4 hydrogen
2 oxygen
2 oxygen
the equation is balanced.
Examples
CuCl2(aq) + Al(s)  Cu(s) +AlCl3(aq)
3CuCl2(aq) + 2Al(s)  3Cu(s) +2AlCl3(aq)
(3:2:3:2)
Examples
Propane, C3H8, burns in oxygen, O2, to form
carbon dioxide and water.
C3H8 + O2 CO2 + H2O
Balance C – then H – then O
C3H8 + 5O2 3CO2 + 4H2O
(1:5:3:4)
Examples
Pentane, C5H12, burns in oxygen, O2, to form
carbon dioxide and water.
C5H12 + O2 CO2 + H2O
Balance C – then H – then O
C5H12 + 8O2 5CO2 + 6H2O
(1:8:5:6)
Examples
Silver nitrate reacts with copper to produce
silver and copper (II) nitrate.
AgNO3 + Cu Ag + Cu(NO3)2
2AgNO3 + Cu 2Ag + Cu(NO3)2
(2:1:2:1)
Examples
Phosphorus reacts with oxygen gas to produce
diphosphorus pentoxide.
P + O2  P2O5
4P + 5O2  2P2O5
(4:5:2)
Examples
C7H14 + O2 CO2 + H2O
Balance C – then H – then O
C7H14 + 10½O2 7CO2 + 7H2O
2C7H14 + 21O2 14CO2 + 14H2O
(2:21:14:14)
Types of Chemical Reactions
•
•
•
•
•
Synthesis / Combination
Decomposition
Single Replacement
Double Replacement
Combustion
Synthesis / Combination
Reactions
Definition: Reaction where two or more
substances react to form a single substance.
A + B  AB
Examples:
2K(s) + Cl2(g)  2KCl(s)
SO2(g) + H2O(l)  H2SO3(aq)
Decomposition Reactions
Definition: Reaction where a single compound
is broken down into two or more products.
AB  A + B
Examples:
2H2O(l)  2H2(g) + O2(g)
CaCO3  CaO + CO2
Single-Replacement Reactions
Definition: Reaction where atoms of one element replace
atoms of a second element in a compound.
XA + B  BA + X
Note: A reactive metal will replace any metal listed below it in the activity series.
Generally, nonmetal replacement is limited to the halogens. The activity of the
halogens decreases as you go down Group 7A of the periodic table. See handout.
Examples:
2AgNO3 + Mg  Mg(NO3)2+2Ag
Mg+LiNO3  no reaction
Li
K
Activity Series
Ca
Na
Mg
Al
Increasing Activity
Any element will replace
any element below it.
Zn
Fe
Pb
(H)*
Cu
Hg
Ag
*Metals from Li to Na will replace H
from acids and water; from Mg to Pb
they will replace H from acids only
For Example…
Ca + MgO  CaO + Mg
The Ca will replace the Mg
because Ca is more active than Mg.
That is to say…Ca is above Mg on the
activity list.
Double-Replacement Reactions
Definition: Reaction that involves an exchange of
positive ions between two compounds.
XA + BY  BA + XY
Note: These reactions generally take place between two ionic compounds in
aqueous solution, and are often characterized by one of the products coming
out of solution in some way.
Examples:
2NaCN(aq)+H2SO4(aq)  2HCN(g)+Na2SO4(aq)
Na2S(aq)+Cd(NO3)2(aq)  CdS(s)+2NaNO3(aq)
Combustion Reactions
Definition: Reaction where an element or a
compound reacts with oxygen, often
producing energy in the form of heat and
light.
Examples:
CH4+2O2  CO2+2H2O + heat + light
2Mg(s)+O2(g)  2MgO(s)
Combustion of Hydrocarbons
If the reactant is a hydrocarbon, the
products are always carbon dioxide and
water.
CH4 + 2O2  CO2 + 2H2O
Ionic Equations
• When a soluble substance is dissolved
in water, the substance often breaks
into ions. This solution is said to be an
aqueous solution.
• Pb(NO3)2(aq)  Pb2+ + 2NO3• NaI(aq)  Na+ + I-
Ionic Equations
• Consider the reaction…
• Pb(NO3)2(aq) + NaI(aq)  PbI2(s) + NaNO3(aq)
• What is really going on is…
• Pb2+ + NO3- + Na+ + I-  PbI2(s) + Na+ + NO3-
• Note that the Na+ ion and the NO3- ion are
not reacting. They are said to be spectator
ions.
Net Ionic Equations
• 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.
Pb2+ + NO3- + Na+ + 2I-  PbI2(s) + Na+ + NO3becomes….
Pb2+ + 2I-  PbI2(s)
From the types of rxns wkst…