Chemical Equations
Preparation for College Chemistry
Columbia University
Department of Chemistry
Chapter Outline
The Chemical Equation
Writing and Balancing Equations
Information in an Equation
Types of Chemical Equations
Heat in Chemical Equations
The Greenhouse Effect
The Chemical Equation
Shorthand Expression for a Chemical Change
 Reactants
 Products
 Stoichiometric Coefficients
 Conditions
 Physical State
2 Al(s) +
Fe2O3 (s)
2 Fe (l) + Al2O3 (s)
Writing Chemical Equations
 Identify the Reaction
magnesium hydroxide + phosphoric acid
magnesium phosphate + water
2 Write the skeleton equation
3 Find the Stoichiometric Coefficients (Balance)
3 Mg(OH)2 + 2 H3PO4
Mg3(PO4)2 + 6 H2O
R
3 Mg
2 PO4
14 O
12 H
P
3Mg
2 PO4
14 O
12 H
Types of Chemical Equations
 Combination:
A + B
AB
 Decomposition
AB
A + B
 Single -Displacement
 Double -Displacement
A + BC
AB + CD
AB + C
AD + CB
Combination Reactions
 metal + Oxygen
metal oxide
2Mg(s) + O2(g)
 nonmetal + Oxygen
2MgO(s)
non metal oxide
2S (s) + 3O2(g)
 metal + nonmetal
2Na (s) + Cl2(g)
 metal oxide + water
MgO (s) + H2O(l)
 nonmetal oxide + water
SO3 (g) + H2O(g)
2SO3 (g)
Salt
2NaCl(s)
Metal Hydroxide
Mg(OH)2(s)
Oxy-acid
H2SO4(s)
Decomposition Reactions
 Metal oxides
2HgO(s)
2Hg (l) + O2(g)
2PbO2(g)
2PbO (g) + O2(g)
 Carbonates and Hydrogen carbonates
CaCO3 (s)
CaO (s) + CO2(g)
2NaHCO3 (s)
Na2CO3 (s) + H2O(l) + CO2(g)
 Other decomposition reactions
KClO3 (s)
2KCl (s) + 3O2(g)
NaNO3 (s)
NaNO2 (s) + O2(g)
2H2O2 (l)
2H2O (l) + O2(g)
2NaN3 (s)
2Na (s) + 3N2(g)
Single-Displacement Reactions
 metal + acid
Hydrogen + Salt
Zn(s) + 2HCl(g)
 metal + water
H2(g) + ZnCl2(s)
Hydrogen + metal hydroxide or oxide
2Na(s) + 2H2O(l)
 metal + Salt
H2(g) + 2NaOH(aq)
Salt + metal
Zn(s) + CuSO4(aq)
 halogen + halide salt
Cl2 (g) + 2NaBr(aq)
ZnSO4(aq) + Cu(s)
Halide salt + Halogen
2NaCl(aq) + Br2(l)
Double-Displacement Reactions
AB + CD
NaCl(aq) + KNO3(aq)
AD + CB
NaNO3(aq) + KCl(aq)
Physical Evidences for double-displacement
 Formation of an Insoluble precipitate
 Evolution of Heat (Neutralization Reactions)
 Gas Formation
Ionic Dissolution
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+
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+
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+
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+
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+
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Precipitation Reactions
Appendix V p. A19
Solubility Rules
NO3-
All nitrates are soluble
Cl-
All chlorides are soluble, except AgCl, Hg2Cl2, Pb2Cl2
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
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)
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
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)
Heat in Chemical Reactions
Endothermic Reaction
Activation
Energy
Products
Net Energy absorbed
Reactants
Time
Heat in Chemical Reactions
Exothermic Reaction
Activation
Energy
Reactants
Net Energy released
Products
Time
Greenhouse Effect
http://web1.infotrac-college.com/wadsworth/
session/61/39/3567398/27!xrn_2_0_A17279460
session/61/39/3567398/12!xrn_39_0_A20080477
session/61/39/3567398/25!xrn_4_0_A15273396
session/61/39/3567398/8!xrn_29_0_A20571782&bkm_8_29
Redox Reactions
(electron-transfer reactions)
Oxidation Number
Oxidation & Reduction
Balancing Redox Reactions
Oxidation number(oxidation state)
# of e- lost, gained or unequally shared by the atom
“pseudocharge” assigned according to arbitrary rules . (rules p.436)
1. ON of an element in an elementary substance is zero
2. H ON = +1, except in metal hydrides NaH, CaH2 What is it?
3. O ON = -2 in most compounds, -1 in peroxides Na2O2 , +2 in OF2
4. ON of metallic elements in ionic compounds is positive.
5. Negative ON is assigned to the most electronegative element
in a covalent compound.
Oxidation number. Calculation
Determine the ON of As in K3AsO4
?
+1
6. In a compound:  ON i  0
-2
K3AsO4
i
As + (-2)x4 +(+1)x3 = 0
As = +5
?
Determine the ON of Cr in Cr2O72-
-2
Cr2O72-
7. In a PAI:  ON i  charge PAI
i
2Cr + (-2)x7 = -2
2Cr = +12
Cr = +6
Oxidation & Reduction
Oxidation (lost of electrons)
ON
-7 -6 -5 -4 -3 -2 -1 0 1 2 3 4 5 6 7
Reduction (gain of electrons)
oxid. # H increases from 0 to +1 (oxidizes)
REDUCING AGENT
O2(g) + H2(g)
OXIDIZING AGENT
2H2O(l)
oxid. # O decreases from 0 to -2 (reduces)
Balancing Redox Equations
 Oxidation number method
(Molecular redox equations)
Two Methods

Ion-electron method
(Ionic redox equations)
Oxidation number method
Oxidation
2KMnO4 + 6 HCl + 5H2S
+1
+7
-2
+1 -1
+1 -2
2 KCl +2 MnCl2 +5 S + 8 H2O
+1 -1
+2
-1
0
Reduction
Reduction:
Mn+7 +5e-
Oxidation:
S-2
2Mn+7 + 5S-2 + 10e2Mn+7 + 5 S-2
Mn+2 x 2
S0 + 2e- x 5
2Mn+2 + 5 S0 + 10e2Mn+2 + 5 S0
Ion-electron method (rules p. 443-444)
Mass and charge must balance
 Acidic Medium
H+(aq)
 Basic Medium
OH-(aq)
Neutralization:
H+(aq) + OH-(aq)
H2O(l)
Ion-electron method (Acidic Medium)
KMnO4 + HCl +
H2S
KCl + MnCl2 + S + H2O
write the molecular equation in ionic form
K+ (aq) + MnO4 - (aq) + H+ (aq) + Cl - (aq ) + 2H+ (aq) + S2-(aq) =
K+ (aq) + Cl - (aq ) + Mn2+ (aq) + 2Cl - (aq ) ) + S0(s) + H2O
Eliminating spectator ions (appear in both sides of the equation)
Net ionic Equation
MnO4 - (aq) + H+ (aq) + S2-(aq)
Reduction
Oxidation
Mn 2+ (aq) + S0 (s)
Write the two half reactions
Reduction:
MnO4- + 8H+
Oxidation:
+ 5eS-2
2MnO4- + 16H+ + 5 S-2
Mn+2 + 4H2O x 2
S0 + 2e-
2Mn+2 + 5S0
x5
+ 8H2O
 Balance elements other than O and H
 Balance O and H, acidic medium:
 Balance each half reaction electrically with electrons:
 Equalize loss and gain of e Add the half equations
Ion-electron method (Basic Medium)
Oxidation
SbO2 - (aq) + ClO2 (aq)
Sb(OH)6 - (aq) + ClO2 - (aq)
Reduction
Write the two half reactions
Oxidation:
SbO2
Sb(OH)6 -
Reduction:
ClO2
ClO2 -
 Balance elements other than O and H
 Balance O and H, ACIDIC medium,
 NEUTRALIZE: add OH- in both sides of the equation
 Balance each half reaction electrically with electrons:
 Equalize loss and gain of e-
Oxidation: SbO2 - + 4H2O + 2OHSbO2 - + 4H2O + 2OHSbO2 - + 2H2O
Reduction:
ClO2 + e-
+ 2OH-
Sb(OH)6 - + 2OH- + 2H+
Sb(OH)6 - + 2H2O
Sb(OH)6 - + 2e-
ClO2 - x 2
SbO2 - + 2OH - + 2H2O + 2ClO2
2ClO2 - + Sb(OH)6 -
Activity Series
of Metals
(table 17.3)
K
Ba
Ca
Na
Mg
Al
Zn
Cr
Fe
Ni
Sn
Pb
H2
Cu
As
Ag
Hg
Au
K+ + eBa+2 + 2eCa2+ + 2eNa+ + eMg2+ + 2eAl3+ + 3eZn2+ + 2eCr3+ + 3eFe2+ + 2eNi2+ + 2eSn2+ + 2ePb2+ + 2e2H + + 2eCu2+ + 2eAs3++ 3eAg + + eHg2++ 2eAu3+ + 3e-
Activity Series of Metals
Useful to Predict the Course of Chemical Reactions
� Na(s) + HCl(aq)
NaCl(aq)
?
+ H2
Net Ionic Reaction:
Na(s) + 2H+(aq)
� Cr(s) + Sn(SO4 )(aq)
2Na+(aq) + H2
?Sn + Cr2 (SO4)3
Net Ionic Reaction:
Cr(s) + 3Sn2+(aq)
� Hg + AgNO3
2Cr3+(aq) + Sn
?No Reaction
Applications
 Electrolytic Cells
Use electrical energy to produce a chemical reaction
 Voltaic (Galvanic) Cells
Use chemical reactions to produce electrical energy
Anode: the OXIDATION SITE
Cathode: the REDUCTION SITE
 Corrosion