1. Types of Chemical Reactions
Synthesis (Combination) reaction
A B  AB
Decomposition reaction
AB  A + B
ASingle-replacement reaction
A BC  AC + B
element
compound
compound
element
BDouble-replacement reaction
AB CD  AD + CB
compound
compound
compound
compound
Neutralization reaction
HX BOH  BX + HOH
Objective:
To classify a chemical reaction as one of the following types: combination (synthesis), decomposition, single replacement, double replacement, or neutralization.
acid
base
salt
water
CH O2  CO2 + H2O
Combustion reaction (of a hydrocarbon)
Polymerization
Polymer = monomer + monomer + …
Ause activity series to predict
Bdriving force…water, gas, or precipitate

2. Types of Chemical Reactions
Synthesis (Combination) reaction
A B  AB
Decomposition reaction
AB  A + B
ASingle-replacement reaction
A BC  AC + B
BDouble-replacement reaction
AB CD  AD + CB
Neutralization reaction
HX BOH  BX + HOH
CH O2  CO2 + H2O
Objective:
To classify a chemical reaction as one of the following types: combination (synthesis), decomposition, single replacement, double replacement, or neutralization.
Combustion reaction (of a hydrocarbon)
Polymerization
Polymer = monomer + monomer + …
Ause activity series to predict
Bdriving force…water, gas, or precipitate

3. Chemical Equations

4. Synthesis Reaction Direct combination reaction (Synthesis)
2 Na Cl  NaCl Na Na Cl Cl  Cl Cl Na Na
General form: A B  AB
element or element or compound
compound compound

5. Synthesis Reaction Direct combination reaction (Synthesis)
2 Na Cl  NaCl Na Cl
Na+
Cl - Cl
Cl -
Na+ Na
General form: A B  AB
element or element or compound
compound compound

6. Formation of a solid: AgCl
AgNO3(aq) + KCl(aq)  KNO3 (aq) + AgCl(s)

7. Pb2+ Cl1- K1+ I1- Pb2+ Cl1- K1+ I1- Cl2 Pb I K KCl PbI2 + +
lead (II) chloride potassium iodide
potassium chloride +
lead (II) iodide
Pb2+
Cl1-
K1+
I1-
Pb2+
Cl1-
K1+
I1-
Cl2 Pb + I K
KCl +
PbI2
(aq)
(ppt)
A reaction that yields an insoluble product, a precipitate, when two solutions are mixed
Are a subclass of exchange reactions that occur between ionic compounds when one of the products is insoluble
Used to isolate metals that have been extracted from their ores and to recover precious metals for recycling
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

8. Pb2+ Cl1- K1+ I1- Pb2+ Cl1- K1+ I1- Cl2 Pb I K KCl PbI2 + +
A reaction that yields an insoluble product, a precipitate, when two solutions are mixed
Are a subclass of exchange reactions that occur between ionic compounds when one of the products is insoluble
Used to isolate metals that have been extracted from their ores and to recover precious metals for recycling
lead (II) chloride potassium iodide
potassium chloride +
lead (II) iodide
Pb2+
Cl1-
K1+
I1-
Pb2+
Cl1-
K1+
I1-
Cl2 Pb + I K
KCl +
PbI2
(aq)
(ppt)
Copyright © 2007 Pearson Benjamin Cummings. All rights reserved.

9. Decomposition Reaction
2 H2O +
2 H2 O2 H H O O H + H O O H
General form: AB A + B
compound
two or more elements
or compounds

10. Single and Double Replacement Reactions
Single-replacement reaction
Mg CuSO4  MgSO Cu
General form:
A BC  AC B
Double-replacement reaction
CaCO HCl  CaCl H2CO3
General form:
AB CD  AD CB

11. Activity Series Element Reactivity Halogen Reactivity Foiled again –
Printable
Version of
Activity
Series
Element Reactivity Ca Li Rb K Ba Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb H2 Cu Hg Ag Pt Au
Foiled again –
Aluminum loses to Calcium
Halogen Reactivity F2
Cl2
Br2 I2
Objectives:
To explain the concept of an activity series for metals.
To predict whether a single-replacement reaction occurs by referring to the activity series.
The Activity Series
– Predicts the outcome of the single-displacement reactions
– Arranges metal and hydrogen in decreasing order of their tendency to be oxidized
– Active metals at the top of the series have the greatest tendency to lose electrons; these including alkali metals (Group 1), alkaline earths (Group 2), and Al (Group 13).
– Inert metals at the bottom of the series have the lowest tendency to be oxidized; these include precious metals or coinage metals located in the lower-right portion of the metals in the periodic table.
– Any element will reduce compounds of the elements below it in the series.
– Only those metals that lie above hydrogen in the activity series dissolve in acids to produce H2.

12. TABLE OF SOLUBILITIES IN WATER
aluminum ss s n i d
ammonium
barium
calcium
copper (II)
iron (II)
iron (III)
lead
magnesium
mercury (I)
mercury (II)
potassium
silver
sodium
zinc
acetate
bromide
carbonate
chloride
chromate
hydroxide
iodide
nitrate
phosphate
sulfate
sulfide
Legend
SOLID
AQUEOUS
i = insoluble
ss = slightly soluble
s = soluble
d = decomposes
n = not isolated

13. TABLE OF SOLUBILITIES IN WATER
aluminum s aq n d
ammonium
barium
calcium ss
copper (II)
iron (II)
iron (III)
lead
magnesium
mercury (I) si
mercury (II)
potassium
silver
sodium
zinc
acetate
bromide
carbonate
chloride
chromate
hydroxide
iodide
nitrate
phosphate
sulfate
sulfide
Legend
s = solid
aq = aqueous
d = decomposes
n = not isolated

14. TABLE OF SOLUBILITIES IN WATER
aluminum ss s n i d
ammonium
barium
calcium
copper (II)
iron (II)
iron (III)
lead
magnesium
mercury (I)
mercury (II)
potassium
silver
sodium
zinc
acetate
bromide
carbonate
chloride
chromate
hydroxide
iodide
nitrate
phosphate
sulfate
sulfide
Legend
s = solid
aq = aqueous
d = decomposes
n = not isolated

15. Solubility Rules 1. Most nitrates are soluble.
Most salts containing Group I ion and ammonium ion, NH4+, are soluble.
Most chloride, bromide, and iodide salts are soluble, except Ag+, Pb2+
and Hg22+.
Most sulfate salts are soluble, except BaSO4, PbSO4, Hg2SO4, and CaSO4.
Most hydroxides except Group 1 and Ba(OH)2, Sr(OH)2, and Ca(OH)2 are
only slightly soluble.
Most sulfides, carbonates, chromates, and phosphates are only slightly
soluble.
Ohn-Sabatello, Morlan, Knoespel, Fast Track to a 5 Preparing for the AP Chemistry Examination 2006, page 91

16. Potassium reacts with Water
P O W !

17. Double Replacement Reaction
K2CO3 (aq)
Potassium carbonate +
BaCl2 (aq)
Barium chloride
2 KCl (aq)
Potassium chloride +
BaCO3 (s)
Barium carbonate

18. Synthesis Reactions Photosynthesis 6 CO2 + H2O C6H12O6 + O2 6 6
Formation of water 2
H O H2O 2
Formation of salt
A photomicrograph of common table salt © Carl J. Bowser, Department of Geology and Geophysics, University of Wisconsin, Madison 2
Na Cl NaCl 2
General Form
A B C

19. Decomposition Reactions
Hydrogen Peroxide 2
H2O H2O O2 2
Electrolysis of water
H2O H O2
electricity 2 2
Nitrogen triiodide 2
NI N I2 3
General Form
AB A B

20. Predict if these reactions will occur3
Mg AlCl3 2 2
Al MgCl2 3
Can magnesium replace aluminum?
YES, magnesium is more reactive than aluminum.
Activity Series
Al MgCl2
No reaction
Can aluminum replace magnesium?
NO, aluminum is less reactive than magnesium.
Therefore, no reaction will occur.
Activity Series
Order of reactants
DOES NOT
determine how
they react.
MgCl Al
No reaction
The question we must ask is can the single element replace its counterpart?
metal replaces metal or nonmetal replaces nonmetal.

21. Single-Replacement Reactions
Activity Series Li Rb K Ba Ca Na Mg Al Mn Zn Cr Fe Ni Sn Pb H2 Cu Hg Ag Pt Au
“Magic blue-earth”
Fe CuCl2
FeCl Cu
Can Fe replace Cu?
Yes
Zinc in nitric acid F2
Cl2
Br2 I2
Zn HNO3
Zn(NO3) H2 2
Can Zn replace H?
Yes
NO REACTION
MgBr Cl2
MgCl Br2
Can Br replace Cl? No
General Form
A BC AC B

22. KOH(aq) + HNO3(aq) KNO3(aq) + H2O ?
How would you prepare potassium nitrate
(using a double replacement reaction)?
potassium nitrate
formation of water
is a driving force. K K OH
KOH HNO KNO H2O + H H
NO3
KNO3 +
H2O
_________
_________
_________
Ca(NO3)2
Both potassium nitrate and
calcium chloride are soluble
(no driving force – no reaction!) 2
KOH +
Ca(NO3)2 2
KNO3 +
Ca(OH)2
Combine a potassium hydroxide solution with nitric acid
to yield soluble potassium nitrate.
KOH(aq) HNO3(aq) KNO3(aq) +
H2O ?
The water could then be removed by distillation to recover solid potassium nitrate.

23. Fe2+ Cl1- Na1+ CO32- Fe2+ Cl1- Na1+ CO32- Cl2 Fe CO3 Na2 NaCl FeCO3
Predict if a reaction will occur when you combine aqueous solutions
of iron (II) chloride with aqueous sodium carbonate solution.
If the reaction does occur, write a balanced chemical equation showing it.
(be sure to include phase notation)
Balanced chemical equation
iron (II) chloride sodium carbonate
sodium chloride +
iron (II) carbonate
Fe2+
Cl1-
Na1+
CO32-
Fe2+
Cl1-
Na1+
CO32-
Cl2 Fe
CO3
Na2
NaCl
FeCO3
(aq)
(ppt)
Using a SOLUBILITY TABLE:
sodium chloride is soluble
iron (II) carbonate is insoluble
FeCl2
Na2CO3
NaCl
FeCO3
(aq)
(ppt) + 2
(aq)
(aq)
Complete Ionic Equation
Fe2+(aq) + 2Cl1-(aq) + 2Na1+(aq) + CO32-(aq) Na1+(aq) + 2Cl1-(aq) + FeCO3(s)