Types of Chemical Reactions and Solution Stoichiometry

1
Types of Chemical Reactions and Solution
Stoichiometry

• Pg. 80-131 in text

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The Bottom LineChapter 4

• Aqueous reactions account for virtually all
  chemistry that takes place in living systems.

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Water, the Common Solvent

• Shape- bent 1050
• Electrons not distributed evenly
• Polar molecule
• Likes dissolve likes
• Polar and ionic
• Polar and polar
• Nonpolar and nonpolar

4
Likes Dissolve Likes

• State whether each pair of substances will mix.
  State Why or why not.
• NaNO3 and H2O
• C6H14 and H2O
• I2 and C6H14
• I2 and H2O

5
Polar Water Molecules Interact with the Positive
and Negative Ions of a Salt
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Water Also Dissolves Many Nonionic Substances

• Why is Ethanol (C2H5OH) very soluble in water?
• Ethanol contains a polar -OH bond like those in
  water

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Polar Bond
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Water Does Not Dissolve all Nonionic Substances

• Water Polar
• Fats Nonpolar
• Likes dissolve likes
• Hydrophilic / Hydrophobic

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Solute / Solvent

• Solute
• If it and the solvent are present in the same
  phase, it is the one in lesser amount.
• If it and the solvent are present in different
  phases, it is the one that changes phase.
• The one that dissolves into the solvent.
• Aqueous solution means that water is the solvent.

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Strong, Weak, Non-Electrolytes
Electrolyte Conductivity Degree of Dissociation Examples
Strong High Total Strong acids, many salts, strong bases
Weak Low to moderate Partial Weak organic acids, weak bases
Non None Close to zero Highly nonpolar organic compounds like sugar, insoluble salts like AgCl
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Acids

• Strong acids - dissociate completely to produce
  H in solution
• hydrochloric and sulfuric acid
• Weak acids - dissociate to a slight extent to
  give H in solution
• acetic and formic acid

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Bases

• Strong bases - react completely with water to
  give OH? ions.
• sodium hydroxide
• Weak bases - react only slightly with water to
  give OH? ions.
• ammonia

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• Strong Acids-
• Virtually every molecule dissociates into given
  ions
• HCl, hydrochloric
• HBr, hydrobromic
• HI, hydroiodic
• HClO3, chloric
• HClO4, perchloric
• HNO3, nitric
• H2SO4, sulfuric

• Strong Bases- soluble compounds containing the
  hydroxide (-OH-) ion
• Group IA metal hydroxides LiOH, NaOH, KOH,
  RbOH, CsOH
• Heavy group 2 metal hydroxides Ca(OH)2,
  Sr(OH)2, Ba(OH)2

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List whether each of the following is a strong,
weak, or nonelectrolyte

1. HClO4
2. C6H12
3. LiOH

• d. NH3
• CaCl2
• HC2H3O2

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Common Terms of Solution Concentration

• Stock - routinely used solutions prepared in
  concentrated form.
• Concentrated - relatively large ratio of solute
  to solvent. (5.0 M NaCl)
• Dilute - relatively small ratio of solute to
  solvent. (0.01 M NaCl)

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The Composition of Solutions

• Molarity moles of solute / liters of solution.
• Example 4.1 pg. 96
• Calculate the molarity of a solution prepared by
  bubbling 1.56 g of gaseous HCl into enough water
  to make 26.8 mL of solution.

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Standard Solutions- a solution whose
concentration is accurately known.

• Example 4.4 pg. 97
• To analyze the alcohol content of a certain wine,
  a chemist needs 1.00 L of an aqueous 0.200 M
  K2Cr2O7 (potassium dichromate) solution.
  Describe how to prepare this solution.

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Ion Concentration in Solutions

• Determine the molarity of Fe3 ions and SO42-
  ions in a solution prepared by dissolving 48.05 g
  of Fe2(SO4)3 in enough water to make 800. mL of
  solution.

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Dilution- water is added to a stock concentrated
stock solution to achieve desired molarity.

• Concept
• Moles of solute after dilution moles of solute
  before dilution
• M1V1 M2V2
• Example
• What volume of 12 M hydrochloric acid must be
  used to prepare 600. mL of a 0.30 M HCl solution?

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Types of Chemical Reactions

• Single Replacement
• Double Replacement
• Synthesis or Combination
• Decomposition
• Combustion
• Acid- Base
• Oxidation-Reduction (Redox)

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Single Replacement Reactions

• In these reactions, a free element reacts with a
  compound to form another compound and release one
  of the elements of the original compound in the
  elemental state. There are two different
  possibilities
• One metal replaces another metal.
• One nonmetal replacement another nonmetal.

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General Format

• A BC ? B AC (when A is a metal)
• K (s) LiCl (aq) ? Li (s) KCl (aq)
• A BC ? C BA ( when A is a nonmetal)
• F2 (g) 2NaBr (aq) ? 2 NaF(aq) Br2 (l)

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SR Reaction Guidelines

1. Not every metal can react and replace or displace
   a metal out of solution.

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SR Reaction Guidelines

• 2. A nonmetal can also replace another nonmetal
  to form a compound. This replacement is usually
  limited to the halogens (F2 , Cl2 , Br2 , I2 ).
  The activity of the halogens decreases as you go
  down the column on the periodic table.

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SRR Examples

• Mg Zn(NO3 )2 ?
• Mg LiNO3 ?
• Cl2 2NaBr ?

26
SRR Net Ionic Equations
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Double Replacement Reactions

• During double replacement, the cations and anions
  of two different compounds switch places, if and
  only if an insoluble product is formed.

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General Format

• AB CD ? AD CB

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DR Reaction Guidelines

• All formulas of the products must be written
  correctly.
• These reactions never involve a change in ionic
  charge.
• A reaction takes place if
• A gas is formed (H2S, CO2 , SO2)
• A covalent substance is formed (H2O, NH3)
• A precipitate is formed (solubility guidelines!)

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Solubility Guidelines
Always Soluble Generally Soluble Generally Insoluble
Alkali metals ions (Li, Na, K, Rb, Cs), NH4, NO3-, ClO3-, ClO4-, C2H3O2- Cl-, Br-, I- Soluble except Ag, Pb2, Hg22 (AP/H). O2-, OH- insoluble except with alkali metals, and NH4, Ca2, Sr2, Ba2 (CBS) somewhat soluble.
F- Soluble except Ca2, Sr2, Ba2, Pb2, Mg2 (CBS-PM). CO32-, PO43-, S2-, SO32-, C2O42-, CrO42- insoluble except with alkali metals or NH4.
SO42- Soluble except Ca2, Sr2, Ba2, Pb2, (CBS/PBS).
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DRR Examples

• AgNO3 (aq) NaCl (aq)?
• CaCO3 (aq) HCl (aq) ?
• Pb(NO3 ) (aq) CuSO4 (aq) ?

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DRR Net Ionic Equations
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The Reaction of K2CrO4(aq) and Ba(NO3)2(aq)
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Describing Reactions in Solution
The Reaction of KCI(aq) and AgNO3(aq)
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Describing Reactions in Solution

• 1. Molecular equation (reactants and products
  as compounds)
• AgNO3(aq) NaCl(aq) ? AgCl(s) NaNO3(aq)
• 2. Complete ionic equation (all strong
  electrolytes shown as ions)
• Ag(aq) NO3?(aq) Na(aq) Cl?(aq) ?
  AgCl(s) Na(aq) NO3?(aq)

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Describing Reactions in Solution (continued)

• 3. Net ionic equation (show only components
  that actually react)
• Ag(aq) Cl?(aq) ? AgCl(s)
• Na and NO3? are spectator ions.

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Write the molecular, complete ionic, and net
ionic forms

• Aqueous nickel (II) chloride reacts with aqueous
  sodium hydroxide

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Selective Precipitation

• Precipitation reactions allow us to target
  specific substances, and separate and recover
  them from a solution.
• Example
• A solution contains Ca2, Cu2, and Pb2. What
  anions can we add, and in what order , to
  separate and recover each cation?

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Stoichiometry of Precipitation Reactions

• Solving problems involving precipitates from
  solutions makes use of molarity, solubility
  rules, balancing equations, and limiting reactant
  calculations.
• Take a systematic approach!
• Example
• What mass of precipitate is produced when 35.mL
  of a 0.250 M Fe(NO3)3 solution is mixed with 55
  mL of a 0.180 M KOH solution?

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Synthesis or Combination Reactions

• In these reactions, two different molecules or
  atoms unite to usually form a single substance.
• Format
• A B ? AB

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Synthesis Reaction Guidelines

1. Direct union of two elements will produce a
   binary compound.
2. Metallic oxides and carbon dioxide react to
   produce carbonates.
3. Binary salts and oxygen react to produce
   chlorate.
4. Metallic oxides and water react to produce a
   base.
5. Nonmetallic oxides and water react to produce an
   acid.

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Synthesis Examples

1. 2 Mg (s) O2 (g) ?
2. Na2 O (s) CO2 (g) ?
3. KCl (s) O2 (g) ?
4. Na2 O (s) H2O (l) ?
5. N2 O5 (s) H2O(l) ?

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Now You Try- Synthesis Reactions

1. K (s) O2 (g)?
2. SO2 (g) H2O (l) ?
3. LiCl (s) O2 (g) ?
4. K2 O (s) H2 O (l) ?

44
Decomposition Reaction

• During decomposition, one compound splits apart
  into two or more substances. These substances
  can be elements or simpler compounds.
• Reaction Format
• AB ? A B

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Decomposition Reaction Guidelines

1. Binary compounds decompose into their component
   elements.
2. Carbonates decompose into an oxide and carbon
   dioxide.
3. Chlorates decompose into binary salt and oxygen.
4. Bases decompose into oxide of their metal and
   water.
5. Acids decompose into the oxide of the nonmetal
   plus water.
6. Most require a large investment of energy in
   either the form of heat or electricity.

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Decomposition Examples

1. 2 NaCl (s) ?
2. 2 Na2 CO3 (s) ?
3. Ba(ClO3 )2 (s) ?
4. Ca(OH)2 (s) ?
5. 2H3PO4 (aq) ?

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Now You Try- Decomposition Reacations

1. Ba(OH)2 (s) ?
2. Ca(ClO3 )2 (s) ?
3. K2CO3 (s) ?
4. H2 SO3 (aq) ?

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Combustion Reactions

• There are 2 types of combustion reactions,
• During a complete combustion reaction, a
  hydrocarbon reacts with pure oxygen to produce
  carbon dioxide and water as the only products.
• During a partial or incomplete combustion
  reactions, a hydrocarbon reacts with atmospheric
  oxygen to produce carbon dioxide, water, carbon
  monoxide, and carbon in the form of soot, smoke
  or ash.

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Combustion Reaction Format

1. Cx Hy (O) O2 (g) ? CO2 (g) H2 O(g)
   (complete)
2. Cx Hy (O) O2 (g) ? CO2 (g) H2 O(g)
   CO C (incomplete)

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Combustion Reaction Example

1. 2C6 H6 (g) 15 O2 (g) ?

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Now You Try- Combustion Reactions

• Assume all are complete combustion unless
  otherwise stated.
• 1. C3 H6 O2 ?
• 2. C3 H8 O2 ?
• 3. C6 H12 O6 O2 ?
• 4. C9 H20 O2 ?

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Acid / Base Reactions

• Acid / Base reactions are a specialized double
  replacement reaction. Where the cation from the
  base hooks up with the anions from the acid to
  form a salt and water.
• Acids are usually compounds that contain loosely
  held hydrogen ions. Follow the pattern HX, where
  X is an anion.
• Bases are compounds that contain loosely held
  hydroxide ions. They follow the pattern YOH,
  where Y is a metal cation.

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Acid / Base Reaction Examples

• 1. HCl (aq) NaOH (aq) ?
• 2. H2 C3H3O2 (aq) Mg(OH)2 (aq)?

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Acids and Bases Neutralize Each Other

• When acids and bases react, they form water and a
  salt. The acid is neutralized by the base and
  the base is neutralized by the acid.
• The concentration of an acid can be determined by
  reacting a set volume with a known concentration
  of base through a process called titration.

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Acid / Base Titration
Ma Va Mb Vb
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Key Titration Terms

• Titrant - solution of known concentration used
  in titration
• Analyte - substance being analyzed
• Equivalence point - enough titrant added to
  react exactly with the analyte
• Endpoint - the indicator changes color so you
  can tell the equivalence point has been reached.

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Acid-Base Neutralization Reactions

• Example 1
• How many mL of a 0.800M NaOH solution is needed
  to just neutralize 40.00 mL of a 0.600M HCl
  solution?

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• Example 2
• You wish to determine the molarity of a solution
  of sodium hydroxide. To do this, you titrate a
  25.00 mL aliquot of your sample, which has had 3
  drops of phenolphthalein indicator added so that
  it is pink, with 0.1067 M HCl. The sample turns
  clear (indicating that the NaOH (aq) has been
  precisely neutralized by the HCl solution) after
  the addition of 42.95 mL of the HCl. Calculate
  the molarity of your NaOH solution.

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Now You Try- Titration Problems

1. A volume of 90. mL of 0.20 M HBr neutralizes a
   60. mL sample of NaOH solution. What is the
   concentration of the NaOH solution?
2. A volume of 46 mL of 0.40 M NaOH neutralizes an
   80. mL sample of HCN solution. What is the
   concentration of the HCN?
3. A volume of 60. mL of 0.60 M HBr solution
   neutralizes a 80. mL sample of Ca(OH)2 solution.
   What is the concentration of Ca(OH)2 ?

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Oxidation /Reduction ReactionREDOX

1. REDOX reactions primarily involve the transfer of
   electrons between 2 chemical species.
2. In these reactions, the oxidation numbers of the
   reactants change.

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Oxidation-Reduction Reactions

• Redox reactions involve a transfer of electrons.
  In order to determine if an electron has been
  transferred, one must be able to assign oxidation
  states.
• Assigning oxidation states to an element in a
  molecule requires knowledge of a set of rules.
  These rules are outlined on page 120 of your text.

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Rules for Assigning Oxidation States

• OS of an atom in an element is 0.
• Na (s), O2 (g)
• 2. OS of a monatomic ion is the same as its
  charge.
• Na OS 1, Cl- OS -1
• 3. In its covalent compounds with nonmetals,
  hydrogen is assigned an OS of 1.
• HCl, NH3, H2O.
• 4. Oxygen is assigned an OS of -2 in its
  covalent compounds.
• CO, CO2, SO2, SO3
• The exception to this rules occurs in peroxides
  (compounds contains the O22- group), where each
  oxygen is assigned an OS of -1.
• H2O2

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• 5. In binary compounds the element with the
  greater attraction for the electrons in the bond
  is assigned a negative OS equal to its charge in
  its ionic compounds.
• HF, NH3, H2S, HI
• 6. The sum of the oxidation states must be zero
  for an electrically neutral compound and must be
  equal to the overall charge for an ionic species.
• NH4, CO32-

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Determine the Oxidation States

• Fe2O3 2Al --gt Al2O3 2Fe
• Iron (III) gains 3 electrons to become elemental
  iron.
• Elemental aluminum lost 3 electrons to become the
  aluminum ion.

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LEO the Lion Says GEROIL RIG

• Loss of Electrons is Oxidation
• Gain of Electrons is Reduction
• Oxidation Involves Loss
• Reduction Involves Gain

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A Summary of an Oxidation-Reduction Process
LEO the lion says GER or OIL RIG
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Fe2O3 2Al --gt Al2O3 2Fe

• Iron (III) ion gained electrons. It has been
  reduced.
• The aluminum lost electrons. It has been
  oxidized.
• The oxidizing agent is the species that is
  reduced (Iron (III)).
• The reducing agent is the species that is
  oxidized (aluminum).

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Which Atoms Undergo Redox?

1. 2H2 (g) O2 (g) --gt 2H2O (g)
2. Zn (s) Cu2(aq) --gt Zn2 (aq) Cu(s)
3. 2AgCl (s) H2 (g) --gt 2H (aq) 2Ag(s) 2Cl-
   (aq)
4. 2MnO4- (aq) 16H (aq) 5C2O42- (aq) --gt
   2Mn2(aq) 10 CO2 (g) 8 H2O (l)

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Now You Try- Redox Reactions

• What is the oxidation number of nitrogen in the
  following molecules?
• NO2 b. NO3- c. N2O5 d. NH4

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Now You Try- Redox Reactions

• K (s) MgCl2 (aq) ?
• 2. In the reaction above determine the following
• The products of the reaction.
• The substance oxidized.
• The substance reduced.
• The oxidizing agent.
• The reducing agent.

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The Half-Reaction Method for Balancing Redox
Reactions

• Half Reactions the two parts of an oxidation
  reduction reaction, one representing oxidation,
  the other reduction.
• Balance the following equation in acid solution
  using the half reaction method.
• Cu(s) HNO3 (aq) --gt Cu2 (aq) NO(g)

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Cu(s) HNO3 (aq) --gt Cu2 (aq) NO(g)

• Identify and write equations for the half
  reactions.
• Copper is being oxidized Cu --gt Cu2
• Nitrogen is being reduced HNO3 --gt NO

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• 2. Balance each half reaction.
• (oxidation) Cu --gt Cu2 2e-
• (reduction) HNO3 --gt NO
• Balance all atoms that are neither oxygen nor
  hydrogen.
• Balance oxygens by adding water to the side that
  needs oxygen.
• HNO3 --gt NO 2H2O
• c. Balance hydrogens by adding H to the side
  that needs hydrogen.
• HNO3 3H --gt NO 2H2O
• d. Balance charges by adding electrons to the
  side that is more positive
• HNO3 3H 3 e - --gt NO 2H2O

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• 3. Equalize the electron transfer. The same
  number of electrons must be gained as are lost in
  the reaction
• 3 Cu --gt 3Cu2 6 e-
• 2HNO3 6H 6e- --gt 2NO 4H2O
• 4. Add the half reactions and cancel
  appropriately to get a complete redox reaction.
• 3 Cu --gt 3Cu2 6 e-
• 2HNO3 6H 6e- --gt 2NO 4H2O
• 3Cu 2HNO3 6H 6e- --gt 3Cu2 2NO 4 H2O
  6e-
• Cancel electrons on both sides and double check.
  Do we have the correct number of electrons on
  both sides?

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Cr2O72- (aq) NO (g) --gt Cr3 (aq) NO3- (aq)
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Balancing Redox Equations in Basic Solutions

• Balance the following equation (it is already
  balanced in acid) assuming it is now in basic
  solution.
• Cr2O72-(aq) 2NO (g) 6H (aq) --gt 2Cr3 (aq)
  2NO3- (aq) 3H2O (l)
• We need to get rid of excess H, because OH- is
  the dominate acid base related species.
• Solution add 6 OH- to both sides of the
  equation.
• Cr2O72-(aq) 2NO (g) 6H (aq) 6OH- --gt
  2Cr3 (aq) 2NO3- (aq) 3H2O (l) 6OH-
• Cr2O72-(aq) 2NO (g) 6 H2O --gt 2Cr3 (aq)
  2NO3- (aq) 3H2O (l) 6OH-
• Cancel waters on both sides
• Cr2O72-(aq) 2NO (g) 3 H2O --gt 2Cr3 (aq)
  2NO3- (aq) 6OH- (aq)

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Simple Oxidation-Reduction Titrations

1. Balance the redox equation.
2. Determine the moles of titrant.
3. Use the balanced redox equation to determine the
   number of moles of unknown.
4. Convert from moles of unknown to grams, percent,
   molarity, or whatever.

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• The use of potassium permanganate (KMnO4) as an
  oxidizing agent is described in your text. A
  0.0483M KMnO4 solution was used to titrate a
  solution containing 0.8329 g of impure calcium
  oxalate, CaC2O4. If 30.25 mL of the KMnO4
  solution was required to reach the titration
  endpoint, calculate the percent purity of the
  CaC2O4.
• MnO4- (aq) C2O42- (aq) --gt Mn2 (aq) CO2 (g)