Chapter 8
“Chemical
Reactions”
Chemistry
Judson High School
Mr. Trotts
1
Describing Chemical Reactions
 OBJECTIVES:
–Describe how to write a
word equation.
2
Describing Chemical Reactions
 OBJECTIVES:
–Describe how to write a
skeleton equation.
3
Describing Chemical Reactions
 OBJECTIVES:
–Describe the steps for
writing a balanced
chemical equation.
4
All chemical reactions…
have two parts:
–Reactants - the substances you
start with
–Products- the substances you
end up with
 The reactants turn into the products.
 Reactants  Products

5
Products
Reactants
6
CHEMICAL REACTION:
a reaction in which one or
more substances are changed
to new substances.
7
REACTANTS:
substances involved in a
chemical reaction.
8
PRODUCTS:
new substances produced in
a chemical reaction.
9
The relationship between
reactants and products can be
written as:
REACTANTS  PRODUCTS
10
In a chemical reaction
Atoms aren’t created or destroyed.
 A reaction can be described several ways:

1. In a sentence (every item is a word)
Copper reacts with chlorine to form copper (II)
chloride.
2. In a word
equation (some symbols used)
Copper + chlorine  copper (II) chloride
11
Symbols in equations
 the
arrow separates the reactants
from the products
–Read as “reacts to form” or yields
 The plus sign = “and”
 (s) after the formula = solid: AgCl(s)
 (g) after the formula = gas: CO2(g)
 (l) after the formula = liquid: H2O(l)
12
Symbols used in equations
 (aq) after the formula = dissolved
in water, an aqueous solution:
NaCl(aq) is a salt water solution
 used after a product indicates a
gas has been produced: H2↑
 used after a product indicates a
solid has been produced: PbI2↓
13
Symbols used in equations
■
indicates a reversible
reaction (more later)

heat
■   ,    shows that
heat is supplied to the reaction
Pt
■   is used to indicate a
catalyst is supplied, in this case,
platinum.
14
What is a catalyst?
A substance that speeds up a
reaction, without being
changed or used up by the
reaction.
 Enzymes are biological or
protein catalysts.

15
3. The Skeleton Equation
 Uses formulas and symbols to
describe a reaction
–but doesn’t indicate how many;
this means they are NOT
balanced
 All chemical equations are a
description that describe reactions.
16
Write a skeleton equation for:
1.
2.
17
Solid iron (III) sulfide reacts with
gaseous hydrogen chloride to form
iron (III) chloride and hydrogen
sulfide gas.
Nitric acid dissolved in water reacts
with solid sodium carbonate to form
liquid water and carbon dioxide gas
and sodium nitrate dissolved in
water.
Now, read these:
Fe(s) + O2(g)  Fe2O3(s)
Cu(s) + AgNO3(aq)  Ag(s) + Cu(NO3)2(aq)
Pt
NO2 (g)   N2(g) + O2(g)
18
4. Balanced Chemical Equations
 Atoms
can’t be created or
destroyed in an ordinary reaction:
–All the atoms we start with we
must end up with
 A balanced equation has the same
number of each element on both
sides of the equation.
19
I. CONSERVATION OF
MASS:
In a chemical reaction, matter
is not created or destroyed but
is conserved.
20
II. IN OTHER WORDS:
The starting mass of the
reactants equals the final
mass of the products.
21
Rules for balancing:
1) Assemble the correct formulas for all the
reactants and products, use + and →
2) Count the number of atoms of each type
appearing on both sides
3) Balance the elements one at a time by
adding coefficients where needed (the
numbers in front) - save balancing the H
and O until LAST!
(I prefer to save O until the very last)
4) Check to make sure it is balanced.
22
Never change a subscript to balance an
equation.
– If you change the formula you are
describing a different reaction.
– H2O is a different compound than H2O2
 Never put a coefficient in the middle of a
formula
2NaCl is okay, but Na2Cl is not.

23
WRITING EQUATIONS
24
III. Chemical reactions can be
described with words such as:
25
solid lead (II) nitrate,
dissolved in water, plus solid
potassium iodide, dissolved
in water, produces solid lead
(II) iodide plus potassium
nitrate, dissolved in water
26
A. All of this information is
important to letting a scientist
know what the reactants and
products are as well as their
physical states.
27
B. A shorthand method has
been developed to describe
chemical reactions.
28
C. This method uses:
1. Chemical formulas
(NaCl for sodium chloride)
29
2. Coefficients
(numbers to indicate how
many molecules)
3. Symbols
(for physical state, catalysts,
direction of reaction, etc.)
30
The reaction described above
would look like this:
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
31
IV. COEFFICIENTS: are the
numbers placed to the left of
the formulas for the reactants
and products.
32
A. The coefficients represent
the number of units of each
substance taking part in a
reaction.
33
Pb(NO3)2(aq) + 2KI(aq)  PbI2(s) + 2KNO3(aq)
B. In the reaction above, there
are:
1. 2 units each of KNO3(aq) and KI(aq)
2. 1 unit each of Pb(NO3)2(aq) and
PbI2(s)
34
V. Symbols are used to
indicate what is happening in
the reaction.
35
A. The physical state of the
reactants
B. Things added to help the
reaction take place
36
C. An indicator to show
which chemicals are the
reactants and which are the
products
37
D. Some commonly used symbols
are in the following table.
38
SYMBOLS USED IN CHEMICAL
EQUATIONS
SYMBOL

+
(s)
(l)
(g)
(aq)
heat

light

Zn

39
MEANING
produces, yields or forms; placed between the reactants and
the products
plus; placed between individual reactants or products in the
equation
solid; placed after the formulas in parentheses on the same
line
liquid
gas
aqueous, a solid is dissolved in water
the reactants are heated; things added to make the reaction
happen are written over the arrow.
the reactants are exposed to light
Zinc was added as a catalyst. A catalyst increases the rate of
a reaction, but is not consumed in the reaction.
Reactants
Products
C3H8(g) + 5O2(g)  3CO2(g) + 4H2O(g)
1
5
3
4
40
unit propane gas;
units oxygen gas
units carbon dioxide gas;
units water vapor
VII. REACTION ENERGY -Chemical reactions either
absorb heat (endothermic) or
release heat (exothermic).
41
A. Endothermic reactions must
have energy added
(usually thermal, sometimes
electrical, or light)
for the reaction to take place.
42
1. This is due to more energy
being required to break bonds
than to make new bonds.
43
2. Some endothermic
reactions cause the reaction
container to feel cool to the
touch.
44
b. Exothermic reactions have
some form of energy released
by the reaction (usually it is
thermal or light).
45
1. This is due to less energy
being required to break bonds
than to form new bonds.
46
2. Exothermic reactions cause
the reaction container to feel
warm or even hot to the touch.
47
BALANCING EQUATIONS
THERE ARE 5 STEPS TO
BALANCING
EQUATIONS.
48
Step 1: Remember these two
rules:
49
Each capital letter
begins a new element
50
Coefficients
(numbers that multiply the
whole formula)
can only be placed in front
of the chemical formula
51
Step 2: Determine each type
of element present on both
sides of the equation
52
Write the equation.
53
Make a table under the
equation that lists each type
of element.
54
List the elements in the
center of the table
55
Example:
MgO + Br2
56
 MgBr2 +
Mg
Br
O
O2
Step 3: Determine the number of
each type of element on both
sides of the equation.
57
Example:
MgO + Br2  MgBr2 +
1
Mg 1
2
Br 2
1
O
2
58
O2
Step 4: Determine which
elements are not balanced.
In the example, the unbalanced
element is oxygen.
59
Step 5: to balance the number of
atoms for the element,
Put a coefficient in front of the
formula with the lower number of
atoms for that element to
produce an equal number of
atoms.
60
Remember:
The coefficient multiplies all of
the elements in the chemical
formula.
61
Start balancing by trying to
make the numbers of elements
in all the other formulas equal
with the most complicated
chemical formula (the formula
with the most elements and
subscripts).
62
Always balance Hydrogen
and Oxygen last.
63
Water can always be added
to the equation
64
Example: Place a coefficient of 2
in front of the MgO.
Then multiply the number of
atoms in the table by the
coefficient.
65
2 MgO + Br2  MgBr2 + O2
2 1 Mg 1
2 Br
2
2 1
O
2
The coefficient changed the number of
Mg also. Now the Mg has to be balanced.
66
Go to the other side of the
equation and place a
coefficient of 2 in front of the
MgBr2.
67
2 MgO + Br2  2 Mg Br2 + O2
2 1
Mg
1 2
2
Br
2 4
2 1
O
2
The coefficient changed the number of Br
also. Now the Br has to be balanced.
68
Place a coefficient of 2 in front
of the Br2.
69
2 MgO + 2 Br2
2 1
4 2
2 1
70
 2
Mg
Br
O
Mg Br2 + O2
1 2
2 4
2
The equation is now balanced
with 2-Mg, 4-Br, and 2-O.
71
OTHER HINTS
72
The diatomic molecules
(Br I N Cl H O F)
are always written as two-atom
molecules
(Br2, I2, N2, Cl2, H2, O2, F2).
73
A coefficient can be changed
when needed to get every
element to balance.
74
Many times there can be an
even amount of atoms on one
side and an odd amount of
atoms on the other side.
]Find a common multiple.
75
Example:
N2 + H2  NH3
2
N 1
2
H 3
There are 2 hydrogen on the left side.
There are 3 hydrogen on the right side.
76
1. Determine the common
multiple between 2 and 3 (it is
6).
77
2. Multiply the H2 by 3 and the
NH3 by 2.
N2 + 3 H2  2 NH3
2 N 1 2
6 2 H 3 6
78
The equation is now
balanced with 2-N and 6-H.
79
Sometimes hydrogen and
oxygen can cause problems.
Two or more sources of
hydrogen or oxygen on one
side of the equation can occur.
80
Example:
ZnS + O2  ZnO + SO2
The product side has oxygen
from 2 sources.
81
1. Make the odd source of
oxygen even by multiplying
ZnO by 2.
ZnS + O2  2 ZnO + SO2
This
1
Zn 1 2
unbalances
1
S 1
the Zinc.
2 O 3 4 (2-O’s
from 2ZnO plus
2-O’s from SO2)
82
2. Place a 2 in front of the
ZnS.
2 ZnS
2
2
This
unbalances
the sulfur.
83
+ O2  2 ZnO + SO2
1
Zn 1 2
1
S
1
2
O
3 4
3. Place a 2 in front of the
SO2.
2 ZnS + O2  2 ZnO + 2 SO2
2 1
Zn
1 2
2 1
S
1 2
This
2
O
3 4 6
unbalances
(2 O’s from 2ZnO plus
the oxygen more.
4 O’s from 2SO2)
84
4. Multiply the O2 by 3 to
balance the equation.
2 ZnS + 3 O2  2 ZnO + 2 SO2
2 1 Zn 1 2
2 1 S
1 2
6 2 O 3 4 6
85
The equation is now balanced
with 2-Zn, 2-S, and 6-O on
each side.
86
A polyatomic ion that is not
broken into parts by the
reaction (it stays the same on
both sides) is treated like an
element.
87
Example: The equation,
H3PO4 + Ca(OH)2  Ca3(PO4)2 + H2O,
has two polyatomic ions,
phosphate (PO4) and hydroxide
(OH).
88
The phosphate has remained
intact.
The hydroxide was broken up
to form water.
89
1. Make a table and list all of
the elements.
H3PO4 + Ca(OH)2  Ca3(PO4)2 + H2O
List the PO4 1 PO4
2
as if it is an 1
Ca
3
element.
5
H
2 List the H
2
O
1
and O
separately and last.
90
Note that the oxygen in the
phosphate ion is not included
when the number of oxygen
atoms is listed.
91
2. The most complicated
formula is the Ca3(PO4)2.
Make the other formulas
balance with the Ca3(PO4)2.
92
Start by balancing the
calcium, multiply Ca(OH)2
by 3.
H3PO4 + 3 Ca(OH)2  Ca3(PO4)2 + H2O
The PO4
1 PO4
2
is now
Ca
3
3 1
unbalanced. 9 5
H
2
O
1
6 2
Do not worry about H and O yet.
93
3. Multiply the H3PO4 by 2.
2 H3PO4 + 3 Ca(OH)2  Ca3(PO4)2 + H2O
2 1 PO4 2
3 1 Ca
3
12 9 5
H
2
6 2
O
1
This leaves 12 H’s and 6 O’s. The ratio of
H to O is 2:1.
This means that the reaction made
6 molecules of water.
94
4. Finish balancing the
equation by multiplying the
water by 6.
2 H3PO4+3 Ca(OH)2  Ca3(PO4)2 + 6 H2O
2 1 PO4
2
3 1 Ca
3
12 9 5
H
2
6 2
O
16
The equation is now balanced with 2-PO4,
3-Ca, 12-H and 6-O.
95
Practice Balancing Examples
1. _AgNO3 + _Cu  _Cu(NO3)2 + _Ag
2. _Mg + _N2  _Mg3N2
3. _P + _O2  _P4O10
4. _Na + _H2O  _H2 + _NaOH
5. _CH4 + _O2  _CO2 + _H2O
96
Types of Chemical Reactions
 OBJECTIVES:
–Describe the five general
types of reactions.
97
Types of Chemical Reactions
 OBJECTIVES:
–Predict the products of the
five general types of
reactions.
98
Types of Reactions
There are millions of reactions.
 We can’t remember them all, but luckily
they will fall into several categories.

We will learn 5 major types.
 Will be able to predict the products.
 For some, we will be able to predict
whether or not they will happen at all.


We recognize them by their reactants
99
#1 - Combination Reactions or
Synthesis Reactions
Combine = put together
 2 substances combine to make one
compound.
 Ca +O2 CaO
 SO3 + H2O  H2SO4
 We can predict the products if the reactants
are two elements.
 Mg + N2 

100
Complete and balance:
+ Cl2 
 Fe + O2  (assume iron (II) oxide is product)
 Al + O2 
 Remember that the first step is to write
the correct formulas – you can still
change the subscripts at this point, but
not later!
 Then balance by using the coefficients
only
 Ca
101
#1 - Combination
 Note:
a) Some nonmetal oxides react
with water to produce an acid:
SO2 + H2O  H2SO3
b) Some metallic oxides react with
water to produce a base:
CaO + H2O  Ca(OH)2
102
#2 - Decomposition Reactions
 decompose
= fall apart
 one reactant breaks apart into two
or more elements or compounds.
electricity
 Na + Cl2
 NaCl   

 CaCO3   CaO + CO2
that energy (heat, sunlight,
electricity, etc.) is usually required
 Note
103
#2 - Decomposition Reactions
 Can
predict the products if it is a
binary compound-Made up of only
two elements
–breaks apart into its elements:
electricity

 H2O   

 HgO  
104
#2 - Decomposition Reactions
 If
the compound has more than
two elements you must be given
one of the products
–The other product will be from
the missing pieces

 NiCO3   CO2 + ___
heat
 H2CO3(aq) CO2 + ___
105
#3 - Single Replacement
 One
element replaces another
 Reactants must be an element and
a compound.
 Products will be a different element
and a different compound.
 Na + KCl  K + NaCl
 F2 + LiCl  LiF + Cl2
106
#3 Single Replacement
Metals replace other metals (and they
can also replace hydrogen)
 K + AlN 
 Zn + HCl 
 Think of water as: HOH
– Metals replace one of the H, and then
combine with the hydroxide.
 Na + HOH 

107
#3 Single Replacement
 We
can even tell whether or not a single
replacement reaction will happen:
–Some chemicals are more “active”
than others
–More active replaces less active
 There is a list on page 333 - called the
Activity Series of Metals
 Higher
108
on the list replaces lower.
The Activity Series of the Metals
Higher
activity
Lower
activity
109
Lithium
Potassium
Calcium
Sodium
Magnesium
Aluminum
Zinc
Chromium
Iron
Nickel
Lead
Hydrogen
Bismuth
Copper
Mercury
Silver
Platinum
Gold
1) Metals can replace other
metals provided that they are
above the metal that they are
trying to replace.
2) Metals above hydrogen can
replace hydrogen in acids.
3) Metals from sodium upward
can replace hydrogen in
water.
#3 Single Replacement
Practice:
6. Fe + CuSO4 
7. Pb + KCl 
8. Al + HCl 
110
#4 - Double Replacement
Two things replace each other.
– Reactants must be two ionic
compounds.
– Usually in aqueous solution
 NaOH + FeCl3 
– The positive ions change place.
 NaOH + FeCl3 Fe+3 OH- + Na+1 Cl-1
 NaOH + FeCl3 Fe(OH)3 + NaCl

111
The Activity Series of the Halogens
Higher Activity
Fluorine
Chlorine
Bromine
Iodine
Lower Activity
Halogens can replace other
halogens in compounds,
provided that they are above the
halogen that they are trying to
replace.
2NaCl(s) + F2(g) 
MgCl2(s) + Br2(g) 
112
2NaF
??? (s) + Cl2(g)
No
???Reaction
#4 - Double Replacement
 Has
certain “driving forces”
–Will only happen if one of the
products:
a) doesn’t dissolve in water and forms
a solid (a “precipitate”), or
b) is a gas that bubbles out, or
c) is a molecular compound (usually
water).
113
Complete and balance:
 assume all of the following
reactions actually take place:
9. CaCl2 + NaOH 
10. CuCl2 + K2S 
11. KOH + Fe(NO3)3 
12. (NH4)2SO4 + BaF2 
114
How to recognize which type
 Look
at the reactants:
E + E = Combination (synthesis)
C
= Decomposition
E+C
= Single replacement
C + C = Double replacement
115
Practice Examples: reaction
type and product
13. H2 + O2 
14. H2O 
15. Zn + H2SO4 
16. HgO 
17. KBr +Cl2 
18. AgNO3 + NaCl 
19. Mg(OH)2 + H2SO3 
116
 Means
#5 - Combustion
“add oxygen”
 Normally, a compound composed of
only C, H, (and maybe O) is reacted
with oxygen – usually called “burning”
 If the combustion is complete, the
products will be CO2 and H2O.
 If the combustion is incomplete, the
products will be CO (or possibly just
C) and H2O.
117
Combustion Examples:
 C4H10
+ O2  (assume complete)
 C4H10
+ O2  (incomplete)
 C6H12O6
 C8H8
118
+ O2  (complete)
+O2  (incomplete)
SUMMARY: an equation...
 Describes
a reaction
 Must be balanced in order to follow the
Law of Conservation of Mass
 Can only be balanced by changing the
coefficients.
 Has special symbols to indicate
physical state, if a catalyst or energy is
required, etc.
119
Reactions
 Come
in 5 major types.
 We can tell what type they are by
looking at the reactants.
 Single Replacement happens
based on the Activity Series
 Double Replacement happens if the
product is a precipitate (insoluble
solid), water, or a gas.
120
Reactions in Aqueous Solution
 OBJECTIVES:
–Describe the information
found in a net ionic
equation.
121
Reactions in Aqueous Solution
 OBJECTIVES:
–Predict the formation of a
precipitate in a double
replacement reaction.
122
Net Ionic Equations
 Many
reactions occur in water- that is,
in aqueous solution
 Many ionic compounds “dissociate”, or
separate, into cations and anions when
dissolved in water
 Now we are ready to write an ionic
equation
123
Net Ionic Equations
 Example:
–AgNO3 + NaCl  AgCl + NaNO3
1. this is the full equation
2. now write it as an ionic equation
3. can be simplified by eliminating ions
not directly involved (spectator ions)
= net ionic equation
124
PO43–
125
S2–
3Cl–
2Ca2+
Na+
Al3+
Net ionic equations
Review: forming ions
Ionic (i.e. salt) refers to +ve ion plus -ve ion
 Usually this is a metal + non-metal or metal +
polyatomic ion (e.g. NaCl, NaClO3, Li2CO3)
 Polyatomic ions are listed on page 95
 (aq) means aqueous (dissolved in water)
 For salts (aq) means the salt exists as ions
 NaCl(aq) is the same as: Na+(aq) + Cl–(aq)
 Acids form ions: HCl(aq) is H+(aq) + Cl–(aq),
Bases form ions: NaOH(aq) is Na+ + OH–
Q - how is charge determined (+1, -1, +2, etc.)?
A - via valences (periodic table or see pg. 95)
 F, Cl gain one electron, thus forming F–, Cl–
 Ca loses two electrons, thus forming Ca2+

126
Background: valences and formulas
Charge can also be found via the compound
E.g. in NaNO3(aq) if you know Na forms Na+, then
NO3 must be NO3– (NaNO3 is neutral)
 By knowing the valence of one element you can
often determine the other valences
Q - Write the ions that form from Al2(SO4)3(aq)?
Step 1 - look at the formula:
Al2(SO4)3(aq)
Step 2 - determine valences: Al3 (SO4)2


(Al is 3+ according to the periodic table)
Step 3 - write ions:
2Al3+(aq) + 3SO42–
(aq)
 Note that there are 2 aluminums because
Al
127has a subscript of 2 in the original
formula
Practice with writing ions
Q - Write ions for Na2CO3(aq)
A - 2Na+(aq) + CO32–(aq) (from the PT Na is 1+.
There are 2, thus we have 2Na+. There is only
one CO3. It must have a 2- charge)
 Notice that when ions form from molecules,
charge can be separated, but the total charge
(and number of each atom) stays constant.
Q - Write ions for Ca3(PO4)2(aq) & Cd(NO3)2(aq)
A - 3Ca2+(aq) + 2PO43–(aq)
A - Cd2+(aq) + 2NO3–(aq)
Q - Write ions for Na2S(aq) and Mg3(BO3)2(aq)
A - 2Na+(aq) + S2–(aq), 3Mg2+(aq)+ 2BO33–(aq)
128
Types of chemical equations
Equations can be divided into 3 types
1) Molecular, 2) Ionic, 3) Net ionic
 Here is a typical molecular equation:
Cd(NO3)2(aq) + Na2S(aq)  CdS(s) + 2NaNO3(aq)
We can write this as an ionic equation (all
compounds that are (aq) are written as ions):
Cd2+(aq) + 2NO3–(aq) + 2Na+(aq) + S2–(aq)
 CdS(s) + 2Na+(aq) + 2NO3–(aq)
 To get the NET ionic equation we cancel out all
terms that appear on both sides:

Net:
129
Cd2+(aq) + S2–(aq)  CdS(s)
Equations must be balanced
There are two conditions for molecular,
ionic, and net ionic equations
Materials balance
Both sides of an equation should have the
same number of each type of atom
Electrical balance
Both sides of a reaction should have the
same net charge
Q- When NaOH(aq) and MgCl2(aq) are mixed,
_______(s) and NaCl(aq) are produced.
Mg(OH)
Write
balanced
molecular, ionic & net ionic
2
equations
130

First write the skeleton equation
2 NaOH(aq) + MgCl2(aq)
 Mg(OH)2(s) +2 NaCl(aq)
Next, balance the equation
Ionic equation:
2Na+(aq) + 2OH-(aq) + Mg2+(aq) + 2Cl-(aq)
 Mg(OH)2(s) + 2Na+(aq) + 2Cl-(aq)
131
LiNO3
Ca3(PO4)2
NaC2H3O2
Net ionic equation:
2OH-(aq) + Mg2+(aq)  Mg(OH)2(s)
Write balanced ionic and net ionic equations:
CuSO4(aq) + BaCl2(aq)  CuCl2(aq) + BaSO4(s)
Fe(NO3)3(aq) + LiOH(aq)  ______(aq) + Fe(OH)3(s)
Na3PO4(aq) + CaCl2(aq)  _________(s) + NaCl(aq)
Na2S(aq) + AgC2H3O2(aq)  ________(aq) + Ag2S(s)
132
Cu2+(aq) + SO42–(aq) + Ba2+(aq) + 2Cl–(aq) 
Cu2+(aq) + 2Cl–(aq) + BaSO4(s)
Net: SO42–(aq) + Ba2+(aq)  BaSO4(s)
Fe3+(aq) + 3NO3–(aq) + 3Li+(aq) + 3OH–(aq) 
3Li+(aq) + 3NO3–(aq) + Fe(OH)3(s)
Net: Fe3+(aq) + 3OH–(aq)  Fe(OH)3(s)
133
For more lessons, visit
www.chalkbored.com
2Na3PO4(aq) + 3CaCl2(aq) Ca3(PO4)2(s)+
6NaCl(aq)
6Na+(aq) + 2PO43–(aq) + 3Ca2+(aq) + 6Cl–(aq)
 Ca3(PO4)2(s)+ 6Na+(aq) + 6Cl–(aq)
Net: 2PO43–(aq) + 3Ca2+(aq)  Ca3(PO4)2(s)
2Na+(aq) + S2–(aq) + 2Ag+(aq) + 2C2H3O2–
(aq)  2Na+(aq) + 2C2H3O2–(aq) + Ag2S(s)
Net: S2–(aq) + 2Ag+(aq)  Ag2S(s)
134
Predicting the Precipitate


Insoluble salt = a precipitate
General solubility rules are found:
a) Table 8.3, p. 227
b) Reference section - page 887
Table A.7 (back of textbook)
gives combinations of cations
and anions that are soluble or
insoluble in water
135