1. Chemical Equations Preparation for College Chemistry
Columbia University
Department of Chemistry

2. Chapter Outline The Chemical Equation Writing and Balancing Equations
Information in an Equation
Types of Chemical Equations
Heat in Chemical Equations
The Greenhouse Effect

3. The Chemical Equation Al + Fe2O3 Fe + Al2O3 2 (s) (l)
Shorthand Expression for a Chemical Change
Al Fe2O3
Reactants
Fe Al2O3
Products 2
Stoichiometric Coefficients
Conditions
(s)
(l)
Physical State

4. Writing Chemical Equations
Identify the Reaction
magnesium hydroxide + phosphoric acid
magnesium phosphate + water
Mg3(PO4) H2O
Write the skeleton equation
Mg(OH) H3PO4
Find the Stoichiometric Coefficients (Balance) 2 6 3
3Mg
PO4 Mg R P
12 H
14 O 2

5. Types of Chemical Equations
Combination:
A + B AB
Decomposition
AB A + B
Single -Displacement
A + BC AB + C
Double -Displacement
AB + CD AD + CB

6. Combination Reactions
metal + Oxygen
metal oxide
2Mg(s) + O2(g) MgO(s)
nonmetal + Oxygen
non metal oxide
2S (s) + 3O2(g) SO3 (g)
metal + nonmetal
Salt
2Na (s) + Cl2(g) NaCl(s)
metal oxide + water
Metal Hydroxide
MgO (s) + H2O(l) Mg(OH)2(s)
nonmetal oxide + water
Oxy-acid
SO3 (g) + H2O(g) H2SO4(s)

7. Decomposition Reactions
Metal oxides
2HgO(s) Hg (l) + O2(g)
2PbO2(g) PbO (g) + O2(g)
Carbonates and Hydrogen carbonates
CaCO3 (s)
CaO (s) + CO2(g)
2NaHCO3 (s)
Na2CO3 (s) + H2O(l) + CO2(g)
Other decomposition reactions
2KClO3 (s)
2KCl (s) + 3O2(g)
2NaNO3 (s)
2NaNO2 (s) + O2(g)
2H2O2 (l)
2H2O (l) + O2(g)
2NaN3 (s)
2Na (s) + 3N2(g)

8. Single-Displacement Reactions
metal + acid
Hydrogen + Salt
Zn(s) + 2HCl(g) H2(g) ZnCl2(s)
metal + water
Hydrogen + metal hydroxide or oxide
2Na(s) + 2H2O(l) H2(g) NaOH(aq)
metal + Salt
Salt + metal
Zn(s) + CuSO4(aq) ZnSO4(aq) + Cu(s)
halogen + halide salt
Halide salt + Halogen
Cl2 (g) + 2NaBr(aq) NaCl(aq) + Br2(l)

9. Double-Displacement Reactions
AB + CD AD + CB
NaCl(aq) + KNO3(aq) NaNO3(aq) + KCl(aq)
Physical Evidences for double-displacement
Formation of an Insoluble precipitate
Evolution of Heat (Neutralization Reactions)
Gas Formation

10. Precipitation Reactions
Appendix V Solubility Table
NO3-
All nitrates are soluble
Cl-
AgCl, Hg2Cl2, PbCl2
All chlorides are soluble, except
SO42-
Most sulfates are soluble, except SrSO4, PbSO4 and BaSO4
CaSO4 is slightly soluble
CO32-
All carbonates are insoluble, except Group I and NH4+
OH-
All hydroxides are insoluble, except group I Sr(OH) 2
and Ba(OH)2. Ca(OH) 2 is slightly soluble
S2-
All sulfides except Groups I and II and NH4+ are insoluble

11. Solubility Rules Used to predict results of precipitation reactions
Example 1
What happens when solutions of Ba(NO3)2 and Na2CO3 are mixed?
Ions present: Ba2+ (aq), NO3-(aq), Na+(aq), CO32-(aq)
Possible precipitates: BaCO3, NaNO3
According to solubility rules, BaCO3 is insoluble
Ba2+(aq) + CO32-(aq) BaCO3(s)

12. Solubility Rules Example 2 Mix solutions of BaCl2, NaOH
ions present: Ba2+(aq) , Cl-(aq), Na+(aq), OH-(aq)
possible precipitates: Ba(OH)2, NaCl
both are soluble; no reaction

13. Net Ionic Equations (Spectator ions do not appear)
Example
Mix solutions of Cu(NO3)2, NaOH
ions present: Cu2+(aq), NO3 -(aq), Na+(aq), OH-(aq)
possible precipitates: Cu(OH)2, NaNO3
NaNO3 is soluble; Cu(OH)2 is not.
Spectator ions: Na+(aq), NO3 -(aq)
Net Ionic Equation:
Cu2+ (aq) + 2 OH- (aq) Cu(OH)2 (s)

14. Heat in Chemical Reactions
Endothermic Reaction
Potential Energy
Activation
Energy
Products
Net Energy absorbed
Reactants
Time

15. Heat in Chemical Reactions
Exothermic Reaction
Potential Energy
Activation
Energy
Reactants
Net Energy released
Products
Time