Transcript Chapter 4 Aqueous Reactions and Solution Stoichiometry

CHM 1045: General Chemistry and
Qualitative Analysis
Unit # 4:
Aqueous Reactions and
Solution Stoichiometry
Dr. Jorge L. Alonso
Miami-Dade College –
Kendall Campus
Miami, FL
Textbook Reference:
•Module # 6 and 4 (V-VII)
•Chapt 4 (Brown & LeMay)
Aqueous
•Chapter # 3-6 to 3-8,Reactions
4&
11-1 to 11-3
Solutions
• Solutions (soln) are homogeneous
mixtures of two or more pure substances.
• The solvent (solv) is present in greatest
abundance.
• All other substances are solutes (solu).
Most useful measure of concentration
of solutions:
Moles of solute
Molarity (M) =
Liters of solution
H2O
Volumetric flask
Aqueous
Reactions
Cu(NO3)2
{PrepASolu}
Solubility of Chemical Substances
Elements: mostly insoluble solids, liquids & gases.
Covalent Compounds: mostly insoluble gases, except
O & N containing organic (C) liquids (polar: acids, bases,
alcohols, etc.)
Ionic Compounds: many are soluble.
Except Except
PBS HAP
SOLUBILITY RULES: for Ionic Compounds (Salts)
1. All salts of alkali metals (IA) are soluble.
2. All NH4+ salts are soluble.
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
6. All O2- are insoluble except for IA metals Ca2+, Ba2+, and Sr2+ salts.
HO
{Soluble metal oxides form hydroxides: CaO 2
Ca 2+ + 2OH-}
7.
8.
9.
All OH- are insoluble except for IA metals, NH4+ & slightly soluble Ca 2+ Ba2+ & Sr2+
All salts containing the anions: CO32-, PO43-, AsO43-, S2- and SO32- are insoluble
except fro IA metals and NH4+ salts.
For salts containing the anions not mentioned above (e.g., CrO42-, Cr2O72-Aqueous
, P3-,
C2O42- etc.) assume that they are insoluble except for IA metals and NH4+ Reactions
salts,
unless, otherwise informed.
The Solution Process:
Ionic vs. Molecular
(1) Ionic Compounds: undergo dissociation - process by
which many ionic substances dissolve in water, the solvent
pulls the individual ions from the crystal and solvates them.
_
{*NaCl + H2O }
Polar water molecule
+
+
H2O
Dissociation
NaCl(s)
H2O
Aqueous
Na+(aq) + Reactions
Cl
(aq)
Electrolytes vs Nonelectrolytes?
Do soluble substances conduct electricity in water?
• Electrolytes substances that
dissociate in water and conduct
electricity (many ionic salts)
{ElectrVsNonE}
H2O
NaCl(s)
H2O(l)
{DoesWaterConduct?}
H2O
Na+(aq) + Cl-(aq)
H+(aq) + OH-(aq)
A nonelectrolytes may dissolve in water, but they do not
dissociate into ions, thus do not conduct electricity. These are
most commonly polar molecular (covalent) compounds.
C6H12O6(s)
Glucose molecules
H2O
C6H12O6 (aq)
Dissolved Glucose molecules
Aqueous
Reactions
Electrolytes: Strong and Weak
A strong electrolyte dissociates
completely when dissolved in water.
HCl (g)
H2O
0
1
{Strong&WeakElectrolytes}
H+(aq) + Cl-(aq)
1
1
A weak electrolyte only dissociates
partially when dissolved in water.
NH4OH(aq)
H2O
.75
1 
Aceitic Acid,
HC2H3O2 (aq)
.80
1 
molecules
NH4+(aq) + OH-(aq)
.25 
H2O
.25 
H+(aq) + C2H3O2-(aq)
.20 
.20 
ions
Aqueous
Reactions
Strong Electrolytes & Ion
Concentration
NaCl (s)
H2O
+
Na + Cl
1
CaCl2 (s)
1
H2O
1
Na3PO4 (s)
1
2+
Ca
1
H2O
1
1M Na3PO4
-
-
+ 2Cl
2
+
3Na + PO4
3
What Molarity of Ions?
(=2)
1
4M in Ions
(=3)
3-
(=4)
Aqueous
Reactions
The Solution Process:
Ionic vs. Molecular
(2) Molecular (covalent) Compounds: mostly insoluble
gases, except polar organic (C) liquids containing O & N
(polar: acids, bases, alcohols, etc.)
Insoluble gases: NO2, CH4, CO2, O2, P2O5, N2, CO, etc.
Polar Covalent {carbon (C) chains containing H,O or N}:
CH3OH, C6H12O6, C6H5OH, etc.
C6H12O6(s)
C3H5OH(l)
H2O
H2O
C6H12O6 (aq)
C3H5OH (aq)
Dissolve without dissociating into ions
Aqueous
Reactions
{Ethanol+Water}
Molecular Compounds
Molecular compounds tend to be nonelectrolytes, except
for strong acids (and weak acids & bases).
Covalent Compounds: HCl, CO2, O2, P2O5, C6H14, C6H12O6, etc.
Solubility Rules:
3. All salts containing the anions:
NO3-, ClO3-, ClO4-, (C2H3O2-) are
soluble.
4. All Cl-, Br-, and I- are soluble
except for Hg22+, Ag+, and Pb2+
salts.
5. All SO42- are soluble except for
Pb2+, Ba2+, and Sr2+.
Strong Electrolytes: HCl.
100% ions
Weak Electrolytes: HF, Ammonia NH3,
Acetic acid HC2H3O2
Non-Electrolyte: H2O, Ethanol C2H5OH
Some ions
Dissolve,
but no ions
Aqueous
Reactions
Chemical Reactions
Occurring in Aqueous
Environments
(1) Precipitation
(2) Gas-Forming
(3) Acid-Base Neutralization
(4) Oxidation-Reduction
(Redox)
Mostly Single
& Double–
Replacement
Reactions
Aqueous
Reactions
(1) Precipitation
Reactions
A special category of Metathesis
(Double Replacement, Exchange)
Reactions
Aqueous
Reactions
Precipitation Reactions
Aqueous solutions, reacting to produce a precipitate
(an insoluble compound). Example: KI(aq) + Pb(NO3)2 (aq)
Predict the solubility of compounds in reaction:
Pb(NO3)2 (aq) + 2KI (aq) 2 KNO3 (aq)+ PbI2 (s)
Precipitate
(ppt)
Pb(NO3)2
Solubility Rules:
3. All salts containing the anions: NO3-, ClO3-, ClO4-,
(C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+,
and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and
Sr2+.
KI
{Movie}
PbI2
Aqueous
Reactions
Precipitation Reactions are Double
Displacement (Replacement)
Does a reaction occur? Does the activity series apply to
double displacement reactions?
AgNO3(aq)+ KCl(aq)
AgCl (s) + KNO3(aq)
• It appears the ions in the reactant compounds exchange ion
• Reaction occurs only if a precipitate is formed!
{*AgNO3+NaCl&NaI}
Solubility Rules:
3. All salts containing the anions: NO3-, ClO3-, ClO4-,
(C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+,
and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
Aqueous
Reactions
Ways of Expressing
Precipitation Reactions
There are three different:
(1) Molecular Equations
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
(2) Ionic Equations
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq) 
AgCl (s) + K+ (aq) + NO3- (aq)
(3) Net Ionic Equations
Ag+(aq) + Cl-(aq)  AgCl (s)
Aqueous
Reactions
Molecular Equation
The molecular equation lists the reactants and
products in their molecular (formula unit) form.
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
Ionic Equation
• In the ionic equation all strong electrolytes (strong
acids, strong bases, and soluble ionic salts) are
dissociated into their ions.
• This more accurately reflects the species that are
found in the reaction mixture.
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq)  Aqueous
Reactions
AgCl (s) + K+ (aq) + NO3- (aq)
Net Ionic Equation
• To form the net ionic equation, cross out anything
that does not change from the left side of the
equation to the right.
Ag+(aq) + NO3-(aq) + K+(aq) + Cl-(aq) 
AgCl (s) + K+(aq) + NO3-(aq)
• The only things left in the equation are those things
that change (i.e., react) during the course of the
reaction. Ag+(aq) + Cl-(aq)  AgCl (s)
• Those things that didn’t change (and were deleted
Aqueous
from the net ionic equation) are called spectator Reactions
ions.
Solution Chemistry
There are three different ways of expressing
precipitation reactions:
(1) Molecular Equations
AgNO3 (aq) + KCl (aq)  AgCl (s) + KNO3 (aq)
(1) Ionic Equations
Ag+ (aq) + NO3- (aq) + K+ (aq) + Cl- (aq) 
AgCl (s) + K+ (aq) + NO3- (aq)
(2) Net Ionic Equations
Ag+(aq) + Cl-(aq)  AgCl (s)
Aqueous
Reactions
Writing Net Ionic Equations
1. Write a balanced molecular equation.
2. Dissociate all strong electrolytes (ionic
equation).
3. Cancel-out ions that remains unchanged
from the left side to the right side of the
equation (spectator ions).
4. Write the net ionic equation with the
species that remain.
Aqueous
Reactions
Writing Net Ionic Equations
+ + SO 2- + Ba2+ + 2NO -  2NH + + 2NO - + BaSO
2NH4(aq)
4(aq)
3(aq)
4(aq)
3(aq)
4 (s)
(aq)
Ba2+(aq)+ SO42-(aq)  BaSO4(s)
Solubility Rules:
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
Aqueous
Reactions
Precipitation Reactions
SOLUBILITY RULES: for Ionic Compounds (Salts)
1. All salts of alkali metals (IA) are soluble.
2. All NH4+ salts are soluble.
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
SOLUBILITY RULES: for Ionic Compounds (Salts)
1. All salts of alkali metals (IA) are soluble.
2. All NH4+ salts are soluble.
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.Aqueous
Reactions
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
SOLUBILITY RULES: for Ionic Compounds (Salts)
1. All salts of alkali metals (IA) are soluble.
2. All NH4+ salts are soluble.
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
SOLUBILITY RULES: for Ionic Compounds (Salts)
1. All salts of alkali metals (IA) are soluble.
2. All NH4+ salts are soluble.
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.Aqueous
Reactions
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
(2) Gas-Forming
Reactions
A special category of Metathesis
(Double Replacement, Exchange)
Reactions
Aqueous
Reactions
Gas-Forming Reactions
These metathesis reactions do not give the expected products.
CaCO3 (s) + 2HCl (aq)  CaCl2 (aq) + CO2 (g)
H2+CO
H23O (l)
NaHCO3 (aq) + HBr (aq)  NaBr (aq) + CO2 (g)
H2+CO
H23O (l)
SrSO3 (s) + 2 HI (aq) 
SrI2 (aq) + SO2 (g)
H2+SO
H23O (l)
• The expected products decompose to give a gaseous
products
• Carbonate + Acid produce H2CO3  CO2 + H2O
• Sulfites + Acids produce H2SO3  SO2+ H2O.
Aqueous
Reactions
{CaCO3 + HCl*}
Gas-Forming Reactions
Aspirin
Alka Seltzer: aspirin + baking soda
C9H8O4
NaHCO3
C6H4(OCOCH3)COOH
Aspirin: 2-(acetyloxy)benzoic acid
or acetyl-salicylic acid
H
CH3
1) C6H4(OCOCH3)COOH + NaHCO3
H2O
C6H4(OCOCH3)COONa(aq) + {H2CO3 (aq) }
2) { H2CO3(aq) }
CO2 + H2O
Aqueous
Reactions
{Alka Seltzer Movie}
Gas-forming Reactions
1.
2.
3.
4.
5.
Aqueous
Reactions
2003 A
Aqueous
Reactions
(3) Acid-Base
Neutralization
Reactions
Another special category of
Metathesis (Double Replacement,
Exchange) Reactions
Aqueous
Reactions
Acids
• Arrhenius: Substances that
release their H+ when
dissolved in water.
• Examples:
Strong:
HCl(g)
100%
molecule
H+(aq) + Cl-(aq)
ions
Weak:
HC2H3O2 (aq).
molecule
15%
H+ + C2H3O2- (aq)
ions
Aqueous
Reactions
Acids
There are only seven
strong acids:
•
•
•
•
•
•
•
Hydrochloric (HCl)
Hydrobromic (HBr)
Hydroiodic (HI)
Nitric (HNO3)
Sulfuric (H2SO4)
Chloric (HClO3)
Perchloric (HClO4)
Solubility Rules:
3. All salts containing the anions: NO3-, ClO3-, ClO4-, (C2H3O2-) are soluble.
4. All Cl-, Br-, and I- are soluble except for Hg22+, Ag+, and Pb2+ salts.
5. All SO42- are soluble except for Pb2+, Ba2+, and Sr2+.
Aqueous
Reactions
Bases
• Arrhenius: Substances
that release their OH−
when dissolved in water.
Strong:
Ba(OH)2 (s)
100%
Weak:
Mg(OH)2 (s)
Ba2+(aq) + 2OH-(aq)
5%
Mg2+(aq) + 2OH-(aq)
{IntroBases}
Aqueous
Reactions
Bases
The strong bases are :
• Alkali metals (IA)
Hydroxides
• Barium Hydroxide
• Strontium Hydroxide
• (weaker: Ammonium,
Calcium Hydroxides)
SOLUBILITY RULES: for Ionic Compounds (Salts)
7.
-
+
All OH are insoluble except for IA metals, NH4 ,
Ba2+,
and
Sr2+.
Aqueous
Reactions
Neutralization Reactions (Arrhenius).
Generally, when solutions of an acid and a base are
combined, the products are a salt and water.
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
H+ (aq) + Cl- (aq) + Na+ (aq) + OH-(aq) 
Na+ (aq) + Cl- (aq) + H2O (l)
H+ (aq) + OH- (aq) H2O (l)
Aqueous
Reactions
Neutralization Reactions
Observe the reaction between a weak
base, Milk of Magnesia, Mg(OH)2 (s)
and a strong acid HCl (aq).
How would you write the net ionic
equation for such a reaction?
{Movie}
Mg(OH)2(s) + 2HCl (aq)  MgCl2(aq) + 2H2O (l)
Mg(OH)2(s) + 2H+(aq) + 2Cl-(aq)  Mg2+(aq) + 2Cl-(aq) + 2H2O (l)
Mg(OH)2(s) + 2H+(aq)  Mg2+(aq) + 2H2O (l)
Aqueous
Reactions
Acid-Base Neutralization Rxs
Aqueous
Reactions
Solution
Stoichiometry
Quantitative aspects of chemical
reactions occurring in aqueous
environments.
Aqueous
Reactions
Preparing Solutions: Molarity
• Most useful way to measure the concentration of a
solution.
moles of solute
Molarity (M) =
volume of solution in liters
{Prep1MSoln}
Example 1: How would you prepare a 1M
solution of CuSO4. 5H2O (FW= 249.7 g/mol)
in a 250 mL volumetric flask ?
?g   1 mol 250mL 
 L 
1L 

1000
mL


 249.7 g   62.42g


1
mol


Example 2: How would you prepare a 0.5M
solution of CuSO4. 5H2O (FW= 249.7 g/mol)
in a 50 mL volumetric flask?
 0.5 mol
?g  
 50 mL
 L 
 1 L   249.7 g   Aqueous
6.242g


 Reactions
 1000mL   1 mol 
Calculations using Molarity
Molarity (M) =
moles of solute ( )
volume of solution in liters (L)
moles of solute ( ) = Molarity ( /L) x volume of solution (L)
mol () = (mol/L) x L = M x V
Aqueous
Reactions
Calculations using Molarity
 1 
# g  

g - MM 



Molarity (M) 
L
L
Problem: What is the molecular weigh (g-MM) of an acid of which
it takes 18.25 g to make 250. mL of a 2.00M concentration?
g
 1 L   1   1000m L 
g

 18.25 g 

  36.5

250
mL

2.00

  1L 


Aqueous
Reactions
Preparing Solutions: Dilution
Moles of
chemical from
Solution 1
M1 x V1
mol1 = mol2
Moles of
chemical in
Solution 2
(mol/L)1 x L1 = (mol/L)2 x L2
Solution 1
Concentrated
M1 x V1 = M2 x V2
M2 x V 2
Solution 2
Diluted
Aqueous
Reactions
Preparing Solutions by Dilution
M1 x V1 = M2 x V2
How would you prepare 500 mL of a 1.0 M solution from a
2.0 M solution?
M1 x V1
mol/L1 x L1 = mol/L2 x L2
mol1 = mol2
M2 x V 2
M 2 V2
V1 

M1
(1.0 mol/L)(0.500 L)
 0.25L
(2.0 mol/L)
Aqueous
Reactions
Preparing Solutions by Dilution
M1 x V1 = M2 x V2
How would you prepare 500 mL of a 1.6 x 10-4 M solution
from a 4.0 x 10-2 M solution?
M1 x V1
mol/L1 x L1 = mol/L2 x L2
mol1 = mol2
M2 x V 2
{*SolnByDilution}
M 2 V2
(1.6 x 104 mol/L)(0.500 L)
V1 

 0.002L
-2
(4.0 x 10 mol/L)
M1
Aqueous
Reactions
2003A #5
Aqueous
Reactions
2005 B
Aqueous
Reactions
Aqueous
Reactions
Aqueous
Reactions
2006 (A)
From:
3.0M NaOH
Prepare:
100 mL
1.0M NaOH
Aqueous
Reactions
100 mL
Vol Flask
Determining the Concentration of
Solutions by Titration
A volumetric analytical technique in which one can
determine the concentration of a solute in a
solution, by making it react with another solution of
known concentration (standard).
Neutralization:
# moles(acid) = # moles(base)
(MxV)acid = (MxV)base
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Solution of
unknown
concentration (MA?)
React a known volume (VA)
Solution of known
concentration (MB)
(Standard)
Measure reacting volume (VB)
Aqueous
Reactions
Titrations
moles base () = MBVB
moles acid () =
moles base ()
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
moles acid () = MAVA
Does the acid need to be in solution?
Can you titrate a solid acid to
determine number of moles of acid?

 Aqueous
1 

acid  # g   gReactions
 - MM 
{A-B w/o Ind}
{*A-B w Ind}
Titration
Phenolphthalein indicator
ACID (clear)
Neutralization:
↔
#moles1(acid) = #moles2(base)
BASE (red)
Aqueous
Reactions
Titration: measuring the
equivalence point
Phenolphthalein
in base
Methyl orange
in acid
A pH meter or indicators are used to determine when
the solution has reached the equivalence point, at
Aqueous
Reactions
which the stoichiometric amount of acid equals that
of
base.
Titration: pH vs. Volume Graph
Excess base
Acid = Base
Excess acid
Aqueous
Reactions
{Titration2}
Titration Calculations:
Stoichiometry using Molarities
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
Neutralization: #moles(acid) = #moles(base)
MAVA = MBVB
Problem: When 20.0 mL of an HCl solution of unknown
concentration was titrated with a STANDARD 0.1M NaOH
solution, a volume of 10.3 mL of NaOH was required to
neutralize the acid. What is the concentration of the HCl
solution?
M B VB
(0.100mol/L)(0.0103L)
MA 

VA
(0.020L)
 0.052 M HCl
Aqueous
Reactions
Using Molarities in
Stoichiometric Calculations
g acid 



η acid
(M x V)acid
g base



η base
(M x V)base
1 mole
 g - MM

1 mole
 g - MM
HN + MOH  MN + HOH
(M x V)acid = (M x V)base
Aqueous
Reactions
Molarity (M)
=
mol of solute
L of solution
 1 mole 
mole s # g Solute

 g - MM 
2H3PO4 + 3 Ca(OH)2  6 HOH + Na3PO4
For titrations:
MA x VA
xA
Where
xA or B
=
MB x VB
xB
= coefficients from balanced equations
Aqueous
Reactions
Solution Stoichiometry Problems:
Molarity
Problem: A volume of 16.3 mL of a 0.30M Ca(OH)2 solution was
used to titrate 25.00 mL H3PO4. What is the concentration of
H3PO4 in the solution of unknown concentration?
2H3PO4 + 3 Ca(OH)2  6 HOH + Na3PO4
?acid  base  2 Acid 
 3 Base 
OR
 acid
base

2 acid 3base
MA x VA = MB x VB
x2A
2M B VB
MA 
3 VA
Where
xA or B
x3B
= coefficients from balanced equations
2 0.30M 16.3 mL  = 0.13 M H PO

3
4
3 25.00 mL 
Aqueous
Reactions
Titrations of Polyprotic Acids
In these
cases there is
an
equivalence
point for each
dissociation.
Aqueous
Reactions
Aqueous
Reactions
(4) OxidationReduction (Redox)
Reactions
These Reactions fall in the
categories of either Double or
Single Replacement Reactions
Aqueous
Reactions
Oxidation-Reduction Reactions
Oxidation occurs when an atom or ion loses electrons.
2 Zn (s) + O2 (g)  2 ZnO (s)
2
2
O2- O2-
Zn
Zn
Zn
Zn2+
Zn
Zn2+
Zn
Zn
Zn
Aqueous
Reduction occurs when an atom or ion gains electrons.
Reactions
Oxidation-Reduction Reactions
One cannot occur
without the other.
2 Zn (s) + O2 (g)  2 ZnO (s)
Aqueous
{oxy-red1} Reactions
How can we determine when an
oxidation-reduction reaction has
occurred?
Na2S (aq) + H2SO4 (aq)  Na2SO4 (aq) + H2S (g)
Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)
Zn (s) + 2 CuNO3 (aq)  2 Cu (s) + Zn(NO3)2 (aq)
Cu (s) + 2 AgNO3 (aq)  2 Ag (s) + Cu(NO3)2 (aq)
HCl (aq) + NaOH (aq)  NaCl (aq) + H2O (l)
To determine if an oxidation-reduction reaction
has occurred, we assign an oxidation number
(charge) to each element in a neutral
Aqueous
Reactions
compound or charged entity.
Rules for Determination of
Oxidation Numbers
1. Elements in their natural elemental form have
an oxidation number of 0.
2 Fe (s) + O2 (g)  2 FeO (s)
Feo
O 2o
• The oxidation number of a monatomic ion is
the same as its charge.
{Fe+O2}
Aqueous
Reactions
Oxidation Numbers
•
Nonmetals in copounds tend to have negative
oxidation numbers, although some are positive in
certain compounds or ions.
2. Oxygen has an oxidation number of −2, except
in the peroxide ion in which it has an oxidation
number of −1.
Peroxide ion O223. Hydrogen is −1 when bonded to a metal, +1
when bonded to a nonmetal.
HCl NaH
2 Fe (s) + O2 (g)  2 FeO (s)
Feo
O 2o
O2-
Aqueous
Reactions
Oxidation Numbers
• Nonmetals tend to have negative oxidation
numbers, although some are positive in certain
compounds or ions.
4. Fluorine always has an oxidation number of −1.
5. The other halogens have an oxidation number of
−1 when they are negative (they can have positive
oxidation numbers, however, most notably in the
polyatomic oxyanions).
hypobromite
BrO
-
Br + (-2) = -1
Br = +1
chlorate
ClO3- Cl + 3(-2) = -1
Aqueous
Reactions
Cl = +5
Oxidation Numbers
6. The sum of the oxidation numbers in a
neutral compound is 0.
7. The sum of the oxidation numbers in a
polyatomic ion is the charge on the ion.
2 Fe (s) + O2 (g)  2 FeO (s)
Feo
O 2o
[Fex + O2- ] = 0
PO43- SO42P + 4(-2) = -3
P = +5
S + 4(-2) = -2
S = +6
Aqueous
Reactions
Oxidation Reduction Reactions
Combination (Synthesis) Reactions
2 Mg (s) + O2 (g)  2 MgO (s)
2 Fe (s) + O2 (g)  2 FeO (s)
2 Fe (s) + 3 Cl2 (g)  2 FeCl3(g)
2 NO (g) + O2 (g)  2 NO2 (g)
{Mg+O2}
{Fe+O2}
{*Fe+Cl2}
{NO+O2}
Displacement Reactions:
Zn (s) + 2 HCl (aq)  ZnCl2 (aq) + H2 (g)
Zn (s) + 2 Cu(NO3) (aq) 2 Cu (s) + Zn(NO3)2 (aq)
Cu (s) + 2 AgNO3 (aq)  2 Ag (s) + Cu(NO3)2 (aq)
Cu (s) + 2 HNO3 (aq)  Cu2+ (aq) + NO2 (g) + H2O
Zn (s) + 2 HNO3 (aq)  Zn2+ (aq) + NO2 (aq) + H2O
Combustion Reaction:
CH4 (g) + 2 O2 (g)  CO2 (g) + 2 H2O (g)
C + 4(+1) = 0
(0)
C + 2(-2) = 0
(+1) + (-2) = 0
{OxyRed}
{Cu+AgNO3}
{*Cu+HNO3}
Cu in Brass Lab
{*Zn+HNO3}
Aqueous
Reactions
Oxidation Reduction Reactions
Zn (s) + SnCl2 (aq)  Sn (s) + ZnCl2 (aq)
{*Zn+SnCl2}
Aqueous
• In displacement reactions, ions oxidize an element.
Reactions
• The ions, then, are reduced.
Oxidation Reduction Reactions
In this reaction, silver ions oxidize copper metal. Aqueous
Cu(s) + 2 Ag+(aq) + 2 NO3-(aq)  Cu2+(aq) + 2 NO3-(aq) + 2 Ag (s)
Reactions
{*Cu+AgNO3}
Oxidation Reduction Reactions
x
The reverse reaction, however, does not occur.
Cu2+ (aq) + 2 Ag (s) 
x Cu (s) + 2 Ag+ (aq)
Aqueous
Reactions
Activity Series
Aqueous
Reactions
Most
active
Non-Metal
Most
active
Metal
Aqueous
Reactions
Table
Continues
Continuation
Most
active
Non-Metal
Most
active
Metal
Aqueous
Reactions
Redox Reactions
1.
2.
Cl2 + 2 KBr
 2 KCl + Br2
Aqueous
Reactions
2006 (B)
Aqueous
Reactions
Aqueous
Reactions
Aqueous
Reactions
Difficult Questions
Aqueous
Reactions
Expressing
Concentrations of
Solutions:
Molarity (& Normality*)
* For MDC students only!
Aqueous
Reactions
Molarity (M)
=
moles of solute
Liters of solution
xA HN + xB MOH  MN + HOH
mol
A
xA
=
mol
B
xB
(mol/L)A x LA
xA
=
(mol/LB) x LB
xB
MA x VA = MB x VB
xA
Where
xA or B
xB
= coefficients for the acid (A) and the base (B)
balanced neutralization equations
Aqueous
Reactions
from
the
mol
A
xA
(mol/L)A x LA
xA
MA x VA
molesA
=
=
=
=
mol
B
xB
(mol/LB) x LB
xB
xAM
B
x VB
B
x VB
xB
xAM
xB
Aqueous
Reactions
xA HN + xB MOH  MN + HOH
For titrations:
Since
 1 mole
mole sA  # g soluteA 
 g - MM A
xA
 # g solute


 g - MM 



=
xB
MB x VB
Aqueous
Reactions
Lesson for MDC students only:
Molarity (M) vs. Normality (N)
M=
mol of solute
L of solution
 1 mole 
mole s # g Solute

 g - MM 
equiv of solute
N=
L of solution
 1 mole 
e quivale nt
s  # g Solute

 g - EW 
g - EW 
M=N
When n = 1
That is when using HCl, KHP NaOH
But not when using H2SO4, Ca(OH)2
g - MM
n
Where:
n
+
=
#
H
or
#OH
A/B
n
Aqueous
=
#e
involved
in
balanced
Redox
Reactions
redox equation.
Molarity (M) vs. Normality (N)
Acid
g-MM (g/)
Molarity (/L)
+ H20 to
Normality (eq/L)
g-EW
HCl
36 g
+ 1L = 1M
=
H2SO4
98 g
+ 1L = 1M
=
H3PO4
98 g
+ 1L = 1M
=
Eq(g/gEW)
1N 36/1 36/36
=36
2N 98/2 98/49
=49
3N 98/3 98/33
=32.7 Aqueous
Reactions
2H3PO4 + 3 Ca(OH)2  6 HOH + Na3PO4
2M H3PO4 3M Ca(OH)2 Using Molarity
1N H3PO4 1N Ca(OH)2
Using Normality
N = nM or M = N
n
Using Normality for titrations:
NA x VA = NB x VB
Aqueous
Reactions