1. III. Types of Chemical Reactions
Ch. 8 – Chemical Reactions
III. Types of Chemical Reactions
(p )

2. Types of Reactions Synthesis reactions Decomposition reactions
There are five types of chemical reactions we will talk about:
Synthesis reactions
Decomposition reactions
Single displacement reactions
Double displacement reactions
Combustion reactions
You need to be able to identify the type of reaction and predict the product(s)

3. Steps to Writing Reactions
Identify the type of reaction
Predict the product(s) using the type of reaction as a model
Balance it
Don’t forget about the diatomic elements! (BrINClHOF)
In a compound, it can’t be a diatomic element because it’s not an element anymore, it’s a compound!

4. A. Synthesis
the combination of 2 or more substances (generally elements) to form a compound
only one product
A + B  AB

5. A. Synthesis
H2(g) + Cl2(g)  2 HCl(g)

6. A. Synthesis
Here is another example of a synthesis reaction

7. Al(s)+ Cl2(g)  2 3 2 AlCl3(s) A. Synthesis
reactant + reactant  1 product
Al(s)+ Cl2(g) 
AlCl3(s)

8. Practice Predict the productsand balance the following reactions.
Sodium metal reacts with chlorine gas
Na(s) + Cl2(g) 
Solid Magnesium reacts with fluorine gas
Mg(s) + F2(g) 
Aluminum metal reacts with fluorine gas
Al(s) + F2(g)  2 2
NaCl(s)
MgF2(s) 2 3 2
AlF3(s)

9. AB  A + B B. Decomposition
a compound breaks down into 2 or more simpler substances
only one reactant
AB  A + B

10. B. Decomposition
2 H2O(l)  2 H2(g) + O2(g)

11. B. Decomposition
Another view of a decomposition reaction:

12. B. Decomposition delta (Δ) - heat
catalyst - a substance that speeds up a chemical reaction without being permanently altered, lowers activation energy. (Therefore catalysts are NOT written in the equation.)
five subcategories to identify!!!!!

13. B. Decomposition Type Reactant Products Example oxy-acids HXO
water + gas (from ion)
H2CO3 --> H2O + CO2
metallic hydroxides
MOH
metal oxide + water
Ca(OH)2 --> CaO + H2O
metallic carbonates
MCO3
metal oxide + CO2
Li2CO3 --> Li2O + CO2
metallic chlorates
MClO3
metal chloride + O2
2KClO3 --> 2KCl +3O2
metallic oxides MO
metal + O2
2HgO --> 2Hg + O2
other binary
XY electricity
X + Y
2NaCl --> 2Na + Cl2

14. Practice N2(g) + O2(g)  BaCO3(s)  Co(s)+ S(s)  NH3(g) + H2CO3(aq)
Identify the type of reaction and write the balanced equation:
N2(g) + O2(g) 
BaCO3(s) 
Co(s)+ S(s) 
NH3(g) + H2CO3(aq)
NI3(s) 
2 NO (g)
BaO(s) + CO2 (g) 2 3
Co2S3 (s) 2
(NH4)2CO3(s) 2
N2 (g) I2 (s) 3

15. A + BC  B + AC C. Single Replacement
one element replaces another in a compound
metal replaces metal (+)
nonmetal replaces nonmetal (-)
A + BC  B + AC

16. C. Single Replacement element + compound product + product
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!)
When H2O splits into ions, it splits into H+ and OH- (not 2H+ and O-2 !!)

17. Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)
C. Single Replacement
Cu(s) + 2AgNO3(aq)  Cu(NO3)2(aq) + 2Ag(s)

18. C. Single Replacement
Another view:

19. C. Single Replacement Fe(s)+ CuSO4(aq)  Cu(s)+ FeSO4(aq)
Products:
metal  metal (+)
nonmetal  nonmetal (-)
free element must be more active (check activity series)
Fe(s)+ CuSO4(aq) 
Cu(s)+ FeSO4(aq)
Br2(l)+ NaCl(aq) 
N.R.

20. C. Single Replacement
Sodium chloride solid reacts with fluorine gas
NaCl(s) + F2(g)  NaF(s) + Cl2(g)
Aluminum metal reacts with aqueous copper(II)nitrate
Al(s)+ Cu(NO3)2(aq) Cu(s) + Al(NO3)3(aq)
Zinc metal reacts with aqueous hydrochloric acid
Zn(s) HCl(aq)  ZnCl2 + H2(g) 2 2 3 2 3 2 2

21. AB + CD  AD + CB D. Double Replacement
ions in two compounds “change partners”
cation of one compound combines with anion of the other
AB + CD  AD + CB

22. Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)
D. Double Replacement
Pb(NO3)2(aq) + K2CrO4(aq)  PbCrO4(s) + 2KNO3(aq)

23. D. Double Replacement Pb(NO3)2(aq)+ KI(aq)  2 2 PbI2(s)+ KNO3(aq)
Products:
switch negative ions
one product must be insoluble/ppt (check solubility table), gas or water
Pb(NO3)2(aq)+ KI(aq) 
PbI2(s)+ KNO3(aq)
NaNO3(aq)+ KI(aq) 
N.R.

24. D. Double Replacement Examples:
AgNO3(aq) + NaCl(s)  AgCl(s) + NaNO3(aq)
K2SO4(aq) + Ba(NO3)2(aq)  KNO3(aq) + BaSO4(s) 2

25. Practice Fe(OH)3(s) + NaCl(aq)
Predict the products. Balance the equation
HCl(aq) + AgNO3(aq) 
CaCl2(aq) + Na3PO4(aq) 
Pb(NO3)2(aq) + BaCl2(aq) 
FeCl3(aq) + NaOH(aq) 
H2SO4(aq) + NaOH(aq) 
KOH(aq) + CuSO4(aq) 
HNO3(aq) + AgCl(s) 3 2
Ca3(PO4)2(s) + NaCl(aq) 6
PbCl2(s) + Ba(NO3)2(aq) 3
Fe(OH)3(s) + NaCl(aq) 3 2 2
H2O(l) + Na2SO4(aq) 2
K2SO4(aq) + Cu(OH)2(s)

26. Net Ionic Equations
Once you write the molecular equation (synthesis, decomposition, etc.), you should check for reactants and products that are soluble or insoluble.
We usually assume the reaction is in water
We can use a solubility table to tell us what compounds dissolve in water.
If the compound is soluble (does dissolve in water), then splits the compound into its component ions
If the compound is insoluble (does NOT dissolve in water), then it remains as a compound

27. Solubility Table
You will be using the rules on your green card!!!!

28. Other Solubilities! Gases only slightly dissolve in water
Strong acids and bases dissolve in water
Hydrochloric, Hydrobromic, Hydroiodic, Nitric, Sulfuric, Perchloric Acids
Group I hydroxides (should be on your chart anyway)
NH4OH goes further to NH3 and H2O
Water slightly dissolves in water! (H+ and OH-)
There are other tables and rules that cover more compounds than your sheet!

29. Total Ionic Equations Molecular Equation:
K2CrO4 + Pb(NO3)2  PbCrO KNO3
Soluble Soluble Insoluble Soluble
Total Ionic Equation:
2K+ + CrO Pb+2 + 2NO3- 
PbCrO4(s) + 2K+ + 2NO3-

30. Net Ionic Equations
These are the same as total ionic equations, but you should cancel out ions that appear on BOTH sides of the equation
Total Ionic Equation:
2K+ + CrO Pb+2 + 2NO3- 
PbCrO4 (s) + 2K+ + 2NO3-
Net Ionic Equation:
CrO Pb+2  PbCrO4 (s)

31. Net Ionic Equations
Write the molecular, total ionic, and net ionic equations Silver nitrate reacts with Lead (II) Chloride in hot water
AgNO3 + PbCl2 
Molecular:
2 AgNO3 + PbCl2  2 AgCl + Pb(NO3)2
Total Ionic:
2Ag+ + 2NO3- + Pb+2 + 2Cl-  2AgCl (s) + Pb+2 + 2NO3-
Net Ionic:
Ag+ + Cl-  AgCl (s)

32. E. Combustion CxHy + O2  CO2 + H2O
the burning of a hydrocarbon in O2 to produce heat, carbon dioxide and water
theory assumes a “complete burn” (blue flame, NOT luminous)
CxHy + O2  CO2 + H2O

33. CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
E. Combustion
Products:
ALWAYS form CO2 + H2O
CH4(g) + 2O2(g)  CO2(g) + 2H2O(g)
C3H8(g)+ O2(g) 
CO2(g)+ H2O(g)

34. E. Combustion Reactions
In order to combust/burn something you need the 3 things in the “fire triangle”:
1) A Fuel (hydrocarbon) 2) Oxygen to burn it with 3) Something to ignite the rxn (spark)

35. Combustion Reactions
Edgar Allen Poe’s drooping eyes and mouth are potential signs of CO poisoning.

36. E. Combustion Examples: C5H12 + O2  C10H22 + O2  8 CO2 + H2O 5 6 231
CO H2O 10 11

37. Mixed Practice
State the type, predict the products, and balance the following reactions:
BaCl2 + H2SO4 
C6H12 + O2 
Zn + CuSO4 
Cs + Br2 
FeCO3 
BaSO4 + HCl 2 9 6
CO2 + H2O 6
ZnSO4 + Cu 2 2
CsBr
FeO + CO2

38. LET’S WATCH!!!!

39. IV. Reaction Energy (p. 514 - 517)
Chapter 17 – Chemical Reactions
IV. Reaction Energy (p )

40. A. Reaction Pathway
Shows the change in energy during a chemical reaction

41. 2H2(l) + O2(l)  2H2O(g) + energy
B. Exothermic Reaction
reaction that releases energy
products have lower PE than reactants
energy released
2H2(l) + O2(l)  2H2O(g) + energy

42. C. Endothermic Reaction
reaction that absorbs energy
reactants have lower PE than products
energy absorbed
2Al2O3 + energy  4Al + 3O2