Honors Chemistry
Chapter 11 Note Packet
(student edition)
Chapter 11 problems:
36, 40, 52, 54, 57, 67, 71, 72
good figures in the book to look at....
Hey, we’re over halfway through the
curriculum. The word on the street is that looking at illustrations, tables, and figures in the book
is a good idea. I’ve done my best to convince you of that fact. I hope that you carry that habit
with you through this course, your other courses, and college.
11.1
Describing Chemical Reactions
chemical equations describe chemical reactions
reactants
chemical equation:

word equation: say
H2
+
Cl2


products
2 HCl
“
”
other example - creation of sodium chloride
Showing Phases in Chemical Reactions:
gas = (
) or
liquid = ( ) solid = (
ex: NaCl ( ) = solid sodium chloride
NaCl (
H2O ( ) - liquid water (not aqueous - redundant)
other symbols:

= heat
one arrow
aqueous (
)
) = solution of sodium chloride (in water)
= one way reaction
= catalyst*
)
2 arrows
= reversible reaction
= precipitate (solid - only found on products side)
* a substance that speeds up a reaction without being used up in the reaction
Balancing Chemical Equations: truly a trial and error process if there ever was one
Helpful hints:
1. 1 atom at a time
2. Balance atoms that appear only 1X per side first
3. Balance polyatomic ions as whole units
4. Balance diatomic elements last
5. Save H + O for last
if this doesn’t succeed, try
Showing Energy Changes in Equations:
everything (particularly with combustion)
1
A + B  C + heat ( - ∆H )
A + B + heat  C ( + ∆H ) OR…
exothermic - releases heat – 1 way to represent this:
endothermic - put in heat – 2 ways to represent this:
Interpreting Chemical Equation - what do chemical equations really mean anyway?
H2O

H2
+
O2
(not balanced)
H2O

H2
+
O2
(
)
atoms
=
atoms
+
atoms (
)
g
=
(
)
molecules
≠
molecules
+
molecules
(
)
moles
≠
moles
+
mole
(
)
11.2
g
+
g
Types of Chemical Reactions
5 types - synthesis, decomposition, combustion, single replacement, double replacement
remember – “First you’ve got to find the right products, then you gotta balance!”
1. Synthesis (direct combination) - needs energy to happen
general formula (
+
):
Ba
+
S

Mg
+
Cl2

Al
+
Cl2

Na
+
O2

2. Decomposition (analysis) - needs energy to happen (usually
)
general formula (
):
FeCl3

HgO

MgSO4
(
.
or

7 H2O
)
2
3. Combustion - the reaction of hydrocarbons and oxygen to yield... “when you just say CHO,...”
general formula:
ex.
combustion of CH4, C3H8, C4H10, C2H5OH
4. double replacement – take place in aqueous solution - little energy involved - usually forms
one soluble ionic product (aka - aqueous) and either a ppt, water, or a gas that bubbles out of
water.
general formula
.
FeCl3 +
NaOH 
H2SO4 +
NaOH 
NH4Cl +
NaOH 

5. single replacement - take place in
solutions and need very little energy
Type 1: positive ions switch
general formula:
AB
+ M
HI
+
Mg

AlCl3
+
Ca

Ca
+
HOH 

MB
+
A
(compound + metal)
Type 2: negative ions switch
general formula:
AB
+ X
NaCl
+
F2

BaS
+
O2

AX +
B (compound + nonmetal)
Hey, some reactions happen and some don’t…
3
For synthesis, combustion, and decomposition, we will assume they all
given sufficient activation energy (
).
For single replacement, use the
!
Ca
+
H2O

Al
+
H2O

Al
+
HI

Cu
+
HI

Rules for the single replacement activity series:
1. Any single element above an element in a compound will
it.
2. The top 5 elements react with
.
3. Metals above H react with
(molecules that start with H – not
water).
4. The nonmetal reactivity series is
For double replacement reactions, use a solubility table and the following rules:
1. If one of the products formed is
, the reaction happens.
2. If a
is formed, the reaction happens.
3. If an
forms (
), the reaction
happens (actually a reaction may happen when two soluble products form, but
it doesn’t go to completion and is not directly observable).
**** refer to chapter 7 lab, “ions in solution”
examples from the lab - page 1, page 3
Na2CrO4
+
KCl
FeCl3
+
KOH


Not balanced yet… “first you gotta find the right products,…”
Note: precipitate symbols include
4
11.3
Reactions in Aqueous Solution
Ions in
solutions for aqueous solutions.
Ionic compounds are
.
When you put them in water, they
Ex1:
NaBr(s) +
Ex2:
MgCl2(s)
H2O(l)
+
H2O(l)
:


Also happens with some molecular substances - particularly acids: we call it
Ex3:
HCl(g)
+
Ex4:
H2CrO4(s)
H2O(l)

H2O(l) 
+
Writing Ionic Equations
so far we’ve looked at molecular equations
now it’s time for ionic equations - show what happens to the ions in a reaction.
first find products, balance, look up solubilities - remember water is a liquid, write net ionic,
then total ionic, identify spectators.....
(“molecular” equation)
Fe(s)
+

2HF(aq)
FeF2(aq)
+
H2(g)
(total ionic)
(net ionic)
spectator ion =
Spectator ions are ions that undergo
(“molecular” equation)
H2SO4(aq)
+
2NaOH(aq)

Na2SO4(aq)
in a chemical reaction.
+
2H(OH)(l)
(total ionic)
(net ionic)
spectator ions =
and
Conservation of Charge - charges must
in a chemical reaction.
5
NIB
Solubility Trends
Cations -
very soluble very insoluble-
Anions -
very soluble for monotomics
very insoluble
sulfides
general trend:
As size decreases, solubility
Some helpful notes on writing phases in chemical reactions:
1. Metals are solids (except
)
2. In single and double replacement reactions, reactants that are compounds are always aqueous.
3. In single and double replacement reactions, products that are compounds should have their
phases identified using a solubility chart (aqueous vs. precipitate)
4. In synthesis and decomposition reactions, ionic compounds are solids.
5. In combustion reactions, the water, CO2, and O2 are gases. The hydrocarbon is hard to tell,
but is usually a liquid after C=5 or higher.
6. Most other covalent compounds are gases.
7. Acids (chemicals starting with hydrogen) are always aqueous.
6