Ch. 9 Chemical
Reactions & Equations
Zn + I2
Reactants
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Zn I2
Product
Introduction
– Chemical reactions occur when bonds
between atoms are formed or broken
– Chemical reactions involve changes in
matter, the making of new materials with
new properties, and energy changes.
– Symbols represent elements,
– Formulas represent compounds,
– Chemical equations represent chemical
reactions
Combustion of Methane
CH4 + 2O2  CO2 + 2H2O
Atoms are rearranged!
Chemical Equations
Yields or produces!
Their Job: Depict the kind of
reactants and products and their
relative amounts in a reaction.
4 Al (s) + 3 O2 (g)
> 2 Al2O3 (s)
The numbers in the front of formulas
are called COEFFICIENTS
The letters (s), (g), and (l) are the
physical states of compounds.
Starting Materials
What’s created
• The charcoal used in a grill is basically carbon. The
carbon reacts with oxygen to yield carbon dioxide.
Word Equations: show the names of the reactants and
the products
This is the “Word Equation” for that reaction
carbon + oxygen
carbon dioxide
A skeleton equation does NOT indicate the
relative amounts of reactants and products
(no coefficients!)
The skeleton equation for that reaction is:
C + O2  CO2 Lavoisier, 1788
Chemical Equations must be balanced in order to
conform to the Law of Conservation of Mass same # & type of atoms on each side of the yield
arrow.
Word Equation:
Methane + Oxygen gas
Skeleton equation:
carbon dioxide + water
CH4 + O2
CO2 + H2O
Symbols Used in Equations
• Solid (s)
• Liquid (l)
• Gas (g)
• Aqueous solution (aq) (dissolved in water)
H2SO4
• Catalyst
• Escaping gas ()
• Change of temperature ()
• Precipitate ( )
Balancing Equations
– When balancing a chemical reaction you
may add coefficients in front of the
compounds to balance the reaction, but
you may
not
change the subscripts.
• Changing the subscripts changes the
compound. Subscripts are determined
by the valence electrons (charges for
ionic or sharing for covalent)
Never put a coefficient in the middle of a formula
2 NaCl is ok Na2Cl is not.
Subscripts vs. Coefficients
• The subscripts
tell you how
many atoms of
a particular
element are in a
compound. The
coefficient tells
you about the
quantity, or
number, of
molecules of
the compound.
There are a number of ways to
interpret balanced equations
2 H2(s) + O2(g) ---> 2 H2O(s)
This equation means:
2 molecules H2 + 1 molecules O2 ---->2 molecules H2O(g)
Steps to Balancing Equations
There are four basic steps to balancing a chemical equation.
1. Write the correct formula for the reactants and the
products. DO NOT TRY TO BALANCE IT YET! You
must write the correct formulas first. And most
importantly, once you write them correctly DO NOT
CHANGE THE FORMULAS!
2. Find the number of atoms for each element on the left
side. Compare those against the number of the atoms
of the same element on the right side.
3. Determine where to place coefficients in front of
formulas so that the left side has the same number of
atoms as the right side for EACH element in order to
balance the equation.
4. Check your answer to see if:
– The numbers of atoms on both sides of the
equation are equal.
– The coefficients are in the lowest possible whole
number ratios. (reduced)
Some Suggestions to Help You
Helpful Hints for balancing equations:
• Take one element at a time.
• Save pure elements for last
• IF everything balances except for that last
pure element, and there is no way to
balance it with a whole number, double all
the coefficients and try again.
• (Shortcut) Polyatomic ions that appear on
both sides of the equation should be
balanced as independent units
That is, don’t separate them into individual
atoms!
Example
Make a table to keep track of atoms
H2 + O2  H2O
R
P
2
H
2
2
O
1
Need another O on the product side
Example
Place a coefficient of 2 in front of H2O
H2 + O2  2H2O
R
Changes the O
P
2
H
2
2
O
1
Example
H2 + O2  2H2O
R
P
2
H
2
2
O
1
This also changes the H
2
Example
H2 + O2  2H2O
R
P
2
H
2
4
2
O
1
2
Now we need twice as much H in the reactant
Example
2H2 + O2  2H2O
R
P
2
H
2
4
2
O
1
2
Add a coefficient of 2 in front of H2
Example
2H2 + O2  2H2O
Your answer
R
4
Recount to check
P
2
H
2
4
2
O
1
2
Balancing Equations
Sodium phosphate + iron (III) oxide 
sodium oxide + iron (III) phosphate
2 Na3PO4 +
3 Na2O + 2 FePO4
Fe2O3 --->
R
L
Na
3
6
2
6
PO4
1
2
1
2
Fe
2
1
2
O
3
1
3
Balancing Equations
5
3
4
__C3H8 + __O
---->
__CO
2
2 + __ H2O
EXAMPLE:
Leave oxygen for last!
Balance B first.
Now balance H
How many O’s are now on the right?
Can you get that many on the left?
__B4H10 + ___O
2
11 2
2
---->___
4 B 2O 3
11
NO
5
+ ___H
10 2O
When this happens, try doubling everything!
How many O’s on the right, now? 22
Can we get that on the left?
YES!!!
•
Types of Reactions
There are millions of chemical reactions.
The only way to be sure what your products
will be is to carry them out in the lab!
• Not very practical – or cost effective,
however…
• There are five types of chemical reactions
we can make some predictions for:
1. Combination reactions (Synthesis)
2. Decomposition reactions
3. Single Replacement reactions
4. Double Replacement reactions
5. Combustion reactions
• You need to be able to identify the type of
reaction and predict the product(s) given
the reactants
1. Combination Reactions
• Combination reactions occur when two or
more substances combine to form a single
compound.
• (Sometimes these are called synthesis
reactions.)
reactant + reactant  1 product
• Basically: A + B  AB
• Example: 2H2 + O2  2H2O
• Example: C + O2  CO2
Combination Reactions
(Synthesis)
2 Mg
Mg
+
O2
O

2 MgO
Mg2+ O2-
O
Mg
O2- Mg2+
General form: A
+
B

AB
element or element or
compound
compound compound
Synthesis Reactions
• example
Balancing
Equations
___
2 Al(s) + ___
3 Br2(l) ---> 2AlBr3(s)
Practice
• Predict the products. Write and balance
the following synthesis reaction equations.
• Sodium metal reacts with chlorine gas
2 Na(s) + Cl2(g)  2 NaCl
• Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g)  MgF2
• Aluminum metal reacts with fluorine gas
2 Al(s) + 3 F2(g)  2 AlF3
2. Decomposition Reactions
• Decomposition reactions occur when a
single compound breaks up into two or
more simpler substances
• 1 Reactant  Product + Product
• In general: AB  A + B
• Example: 2 H2O  2H2 + O2
• Example: 2 HgO  2Hg + O2
Decomposition Reactions
• Another view of a decomposition reaction:
Example:
Decomposition of hydrogen peroxide.
• Word equation:
hydrogen peroxide  water + oxygen
Example:
Decomposition of hydrogen peroxide.
• Skeleton equation:
H2O2

H2O
+
O2
Example:
Decomposition of hydrogen peroxide.
• Balanced Equation:
2H2O2

2H2O
+
O2
Example: Decomposition of
hydrogen peroxide in the presence
of a catalyst.
Balanced equation showing
• the catalyst (MnO2)
• the state of the reactants and products:
MnO2
2H2O2(l)

2H2O (l)
+ O2 (g)
Decomposition Types
• These are hardest to predict products for!
• There are many special cases, but we will not
explore those in Chemistry. For Ex.:
• Carbonates and chlorates are special case
decomposition reactions that do not go to the
elements.
• Carbonates (CO32-) decompose to carbon dioxide
and a metal oxide
• Example: CaCO3  CO2 + CaO
• Chlorates (ClO3-) decompose to oxygen gas and a
metal chloride
• Example: 2 Al(ClO3)3  2 AlCl3 + 9 O2
• I
Practice
• Predict the products. Then, write and
balance the following decomposition
reaction equations:
• Solid Lead (IV) oxide decomposes
PbO2 
Pb + O2
• Aluminum nitride decomposes
2 AlN  2 Al
+
N2
3. Single Replacement Reactions
• Single Replacement Reactions occur when
one element replaces another in a compound.
• A metal can replace a metal (+) OR
a nonmetal can replace a nonmetal (-).
• element + compound element + compound
A + BC  AC + B (if A is a metal) OR
A + BC  BA + C (if A is a nonmetal)
*(remember the cation always goes first!)
*Use Activity Series of Metals
to see if reaction works.
Single Replacement Reactions
• Another view:
Single Replacement Reactions
• Single Replacement Reaction
Mg
+
CuSO4
General form:
A
+ BC


MgSO4
AC
+
+
Cu
B
Here is a mixture of aluminum 2 Al + Fe2O3
powder and iron (III) oxide
powder.
Once ignited the aluminum rips
the oxygen off of the iron
oxide resulting in molten iron,
which falls down to the
pot of sand below. One time
someone used wet sand,
which was a mistake because
the molten iron turned
the water to steam instantly.
This sprayed molten iron
all over the front two rows of
students.
Write the balanced equation for this reaction.
2 Fe + Al2O3
We have looked at several reactions:
Such experiments reveal trends. The Activity
Series of Metals ranks the relative
No,reactivity
Ni is
of metals.
Yes,
Li is
below
Na
Used predict if single displacement
abovereactions
Zn
will occur: metals near the top are most
Yes, Al
is the
reactive and will displacing metals
near
above Cu
bottom.
Q: Which of these will react?
(SO4Fe
)3 is
Fe + CuSO4  Cu + Fe2Yes,
above Cu
Ni + NaCl
 NR (no reaction)
Li + ZnCO3  Zn + Li2CO3
Al + CuCl2  Cu + AlCl3
p. 295
Li
K
Ca
Na
Mg
Al
Zn
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Au
Activity Series
H is the only nonmetal listed.
H2 may be displaced from acids
can be given off when a metal
reacts with H2O.
Ca
Foiled again –
Aluminum loses to Calcium
Metal Reactivity
Li
Rb
K
or Ba
Ca
Na
Halogen Reactivity
Mg
Al
Mn
F2
Zn
Cl2
Cr
Br2
Fe
Ni
I2
Sn
Pb
H2
Cu
Hg
Ag
Pt
Au
Single Replacement Reactions
• Write and balance the following single
replacement reaction equation:
• Zinc metal reacts with aqueous
hydrochloric acid
Zn(s) + 2HCl(aq) 
ZnCl2 + H2(g)
Note: Zinc replaces the hydrogen ion in the
reaction
Li
K
Ca
Na
Mg
Al
Zn
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
Au
Single Replacement Reactions
• Sodium chloride solid reacts with fluorine gas
2 NaCl(s) + F2(g)  2 NaF(s) + Cl2(g)
Note that fluorine replaces chlorine in the compound
• Aluminum metal reacts with aqueous copper
(II) nitrate
Al(s)+ Cu(NO3)2(aq)
Li
Complete these reactions:
Cu + Sn2S  NR
Ca + Cu2SO4  2 Cu + CaSO4
2 Na + PbO  Pb + Na2O
K
Ca
Na
Mg
Al
Zn
Fe
Ni
Sn
Pb
H
Cu
Hg
Ag
4. Double Replacement Reactions
• Double Replacement Reactions occur when
a metal in one compound replaces a metal in
another compound.
compound + compound  compound + compound
• AB + CD  AD + CB
Double replacement reactions
usually occur between 2 ionic
compounds in aqueous solution.
THEY ARE USUALLY DRIVEN BY THE
FORMATION OF EITHER A GAS, A
PRECIPITATE (INSOLUBLE SOLID),
OR A MOLECULAR COMPOUND (LIKE
WATER).
Double Replacement Reactions
Double-replacement reaction
CaCO3
+
2 HCl
General form:
AB
+ CD


CaCl2
AD
+
H2CO3
+
CB
Double Replacement Reactions
• Think about it like “foil”ing in algebra, first and
last ions go together + inside ions go together
• Example:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
• Another example:
K2SO4(aq) + Ba(NO3)2(aq) 2 KNO3(aq) + BaSO4(s)
MnO2 + CO --> Mn2O3 + CO2
Carbon monoxide is
commonly used to strip off
oxygen atoms from metals.
This reaction shows the first
step. If more CO is present,
eventually all oxygen atoms
will be grabbed by CO and
manganese (Mn) metal will
be left.
This is how metal ores get
converted to metals.
Balance this equation! 2 MnO2 + CO --> Mn2O3 + CO2
Practice
•
Predict the products. Balance the equation
1.
2.
3.
4.
5.
HCl(aq) + AgNO3(aq) 
CaCl2(aq) + Na3PO4(aq) 
Pb(NO3)2(aq) + BaCl2(aq) 
FeCl3(aq) + NaOH(aq) 
H2SO4(aq) + NaOH(aq) 
6. KOH(aq) + CuSO4(aq) 
5. Combustion Reactions
• Combustion reactions occur whenever
something reacts with oxygen gas –
often producing noticeable heat & light.
• This is also called burning!!!
Combustion of Hydrocarbons
Hydrocarbons are compounds made of
hydrogen & carbon. Surprise!
There are *gobs of energy stored in the
bonds between carbon atoms!
When we burn them (react them with O2)
we release that stored energy to do work
for us!
*Important scientific term!
Combustion is used to heat homes, cook food,
and power automobiles. -propane is C3H8
-octane, as in gasoline, is C8H18
-ethane, as in ‘natural gas’, is C2H6
Complete Combustion of Hydrocarbons
• Meaning – enough oxygen to
react with all of the carbon!
• Products are ALWAYS carbon
dioxide and water.
• In general:
CxHy + O2  CO2 + H2O
Incomplete Combustion of Hydrocarbons
•Incomplete Combustion does cause
some by-products like carbon monoxide &
elemental carbon. (CO & C)
It occurs when oxygen can’t get to all the
carbon atoms fast enough.
*Hint: If you see CO and/or C as products
it is incomplete combustion!
Combustion
Reactions
Edgar Allen Poe’s
drooping eyes and
mouth are potential
signs of CO
poisoning.
Combustion of Hydrocarbons
• *Balance C’s, then H’s. Save O’s for last!
• Example:
•
C5H12 + 8 O2  5 CO2 + 6 H2O
• Write the products and balance the
following combustion reaction:
•
C10H22 + O2 
•
Some steps for Writing Reactions
1. Identify the type of reaction
2. Predict the product(s) using the type of
reaction as a model
Remember! If you write a formula – You MUST
balance it!
3. Balance the equation (using coefficients)
Don’t forget about the diatomic elements!
(BrINClHOF or HON 7) For example,
Oxygen is O2 as an element.
In a compound, it can’t be a diatomic element
because it’s not an element anymore, it’s part
of a compound!
Practice
Identify the type of reaction for each of the
following synthesis or decomposition
reactions, and write the balanced equation:
N2(g) + O2(g)  Nitrogen monoxide
BaCO3(s) 
Co(s)+ S(s)  (make Co be +3)
NH3(g) + H2CO3(aq) 
NI3(s) 
Mixed Practice
•
1.
2.
3.
4.
5.
State the type, predict the products, and
balance the following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 