Download Section 4.8: Acid-Base Reactions

AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
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Section 4.1: Water, the Common Solvent
Properties of water: called the “universal solvent”
 Tetrahedral electron –pair geometry; bent molecular shape (angle ≈ 105)
 O-H bonds are polar covalent
 Polar covalent bonds have unequal sharing of electrons in the bond
o The oxygen is more electronegative, so it has a greater attraction between the oxygen and the
electrons in the bonds, giving oxygen a partial negative charge and hydrogen a partial positive
charge - results in water being a polar molecule
 Hydration – positive end of water attracted to negative ions in salt and negative end of water attracted to
positive ions in salt, causing the ionic salt to dissolve and dissociate
 nonionic (covalent), polar compounds also dissolve in water
 nonpolar substances generally not soluble in water
Polar vs. nonpolar: Polar molecules have differently charged ends; they are “asymmetrical” in shape. Nonpolar
molecules have even charge distribution; they are “symmetrical” in shape.
“Like dissolves like”:
Section 4.2: Electrolytes
 compounds that conduct an electrical current, as measured by a conductivity meter
 solutions of ionic compounds conduct electricity, and are strong electrolytes
 solutions of polar covalent compounds and covalent compounds are either weak electrolytes or nonelectrolytes
strong electrolytes
weak electrolytes
nonelectrolytes
• completely dissociate into ions
• incompletely dissociate into ions
• no dissociation
• solutions strongly conduct electricity
• solutions weakly conduct
electricity
• solutions don't conduct
electricity
• typical compounds:weak acids &
insoluble salts
• typical compounds: molecular
compounds
• typical compounds: soluble ionic
compounds strong acids & bases
A conductivity meter is used in a solution to gauge the ability of a solution to conduct electrical current. A light source is
usually connected to the meter. If the light bulb burns brightly, the current is strong, and the solute is a strong
electrolyte.
(Zumdahl, Pg. 180 #11) Show how each of the following strong electrolytes dissociates into its component ions upon
dissolving in water.
a. NaBr
b) MgCl2
c) Al(NO3)3
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
(Zumdahl, Pg. 180 #17) Calculate the concentration of all ions present in each of the following solutions of strong
electrolytes.
a. 0.15 M CaCl2
0.26 M Al(NO3)3
Section 4.3: Molarity, Preparation of Standard Solutions, and Dilution
Molarity, or M =
molesofsolute
volumeofsolutioninli ters
The units of M are mol/L or moles/L
# Moles of solute = Volume of solution, in L  Molarity of solution, in M
For a dilution, M1V1 = M2V2
Technique: Preparation of a Standard Solution
ex. Describe each step of the preparation of a standard solution of 500 mL of 0.15 M KMnO4, and include the necessary
calculations as needed.
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
ex. Calculate the mass of NaCl needed to prepare 250 mL of a 0.85 M NaCl solution.
ex. You need to dilute 100. mL of 0.100 M sodium chloride by a factor of 2. How much water do you need to
add? (This is a great thinking question and a great shortcut on the AP exam.)
ex. Determine the molarity of the iron (III) ion and the sulfate ion in a solution prepared by dissolving 48.05 g of
Fe2(SO4)3 in enough water to make 800 mL of solution.
ex. What volume of 12 M hydrochloric acid must be used to prepare 600. mL of a 0.30 M HCl solution?
Describe the steps and equipment necessary to make this solution.
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
Section 4.5: Precipitation Reactions
Two compounds react to form two new compounds. All double replacement reactions must have a "driving force" that
removes a pair of ions from solution. Ions in a precipitation reaction will keep their same charges as reactants and
products.
Formation of a precipitate: A precipitate is an insoluble substance formed by the reaction of two aqueous substances.
Two ions bond together so strongly that water cannot pull them apart, so precipitation occurs almost always in aqueous
solution.
To predict the formation of a precipitate, you must memorize the following solubility rules:
Simple Rules for Solubility
1. Most nitrate (NO3−) salts are soluble.
2. Most alkali (group 1A) salts and NH4+ are soluble.
3. Most Cl−, Br−, and I− salts are soluble (NOT Ag+, Pb2+, Hg22+ )
4. Most sulfate salts are soluble (NOT BaSO4, PbSO4, HgSO4, CaSO4)
5. Most OH− salts are only slightly soluble (NaOH, KOH are soluble, Ba(OH)2, Ca(OH)2 are marginally soluble)
6. Most S2−, CO32−, CrO42−, PO43− salts are only slightly soluble.
ex. Solutions of silver nitrate and lithium bromide are mixed. Write a balanced molecular equation for this reaction.
The LiNO3 (aq) dissociates (separates) as soon as it is formed, since all nitrates are soluble in water (see rules below). The
Li+ and NO3- are soluble and have not formed a new visible product, so we call these spectator ions. In a net ionic
equation, we leave these out, so the only action we observe is the formation of the precipitate.
ex. Solutions of silver nitrate and lithium bromide are mixed. Write a balanced net ionic equation for this reaction.
You Try: Write a balanced net ionic equation for each reaction.
ex. A clear, colorless solution of calcium chloride is dropped into a solution of lead (II) nitrate.
ex. Magnesium nitrate and sodium hydroxide solutions are mixed together.
ex. Powdered iron (III) nitrate is stirred into a solution of potassium carbonate.
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
Section 4.6: Solution Stoichiometry and Gravimetric Analysis
You must have a correctly balanced chemical equation!
Remember # moles of solute = (molarity of solution)(volume of solution)
(Zumdahl, pg. 182, #39) What mass of NaCl is required to precipitate all the silver ions from 50.0 mL of a 0.0500 M
solution of AgNO3?
(Zumdahl, pg. 182, #41) What mass of solid aluminum hydroxide is produced when 50.0 mL of 0.200 M Al(NO3)3 is
added to 200.0 mL of 0.100 M KOH? Which is the limiting reactant? The excess reactant? What is the concentration
of the excess reactant left over?
ex. Calculate the mass of Ag2S produced when 125 mL of 0.300 M AgNO3 is added to 125 mL of 0.100 M sodium sulfide.
Which is the limiting reactant? The excess reactant? What is the molarity of the excess reactant left over?
ex. 100. mL of 0.100 M barium nitrate and 100. mL of 0.200 M barium chloride are mixed together.
a) Will a precipitate form? Why or why not?
b) What volume of 1.00 M sodium sulfate do you need to precipitate all the barium ion?
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
(Zumdahl, pg. 184, #75) Douglasite is a mineral with the formula 2KCl∙FeCl2∙2H2O. Calculate the mass percent of
douglasite in a 455.0-mg sample of rock if it took 37.20 mL of a 0.1000 M AgNO3 solution to precipitate all of the Cl- as
AgCl. Assume the douglasite is the only source of chloride ion.
Gravimetric Analysis
One method for determining the amount of a given substance in solution is to form a precipitant that includes the
substance. The precipitant is then filtered, dried, and weighed. This process is called gravimetric analysis.
Technique: Gravimetric Analysis
 Weigh the solid containing the desired ion (ex. Cl-)
 Dissolve the solid
 Add a large excess of an aqueous ion to the solid in solution that will precipitate the desired ion (ie.
the Ag+ in silver nitrate will precipitate the desired Cl- in the solid)
 Filter the precipitate (now AgCl)
 Wash the precipitate, then dry it in a drying oven
 Weigh the precipitate once it is cool.
Since you know the identity of the precipitate, you can calculate the % of the element in the precipitate (ie. find the %
composition of the element in the precipitate) and then use stoichiometry to determine the # moles and mass of the
desired ion in the original solid.
ex. 0.365 g of an unknown contains Ag. If chloride ion is added to a solution of the unknown and 0.206 g of silver
chloride is produced, what is the amount of Ag in the unknown?
AP Chemistry Unit 2: Types of Chemical Reactions, Reaction Prediction and Solution Stoichiometry
(Zumdahl, Pg. 184 #77) Chlorisondiamine, C14H18Cl6N2, is a drug used to treat hypertension. A 1.28-g sample of
medication containing this drug was treated to release all the chlorine as chloride ion. When the filtered solution
containing chloride ion was treated with an excess of silver nitrate, 0.104 g silver chloride was recovered. Calculate the
mass percent of chlorisondiamine in the medication, assuming the drug is the only source of chloride.
ex: Phosphorite, also called phosphate salt, is a mineral containing PO4 -3 and OH- anions and Ca+2 cations. It is treated
with sulfuric acid in the manufacture of phosphate fertilizers. A chemist analyzes for the calcium content in an impure
phosphate salt by weighing out a 0.4367 g sample, dissolving it in water, and precipitating the Ca+2 ions as the insoluble
hydrated salt CaC2O4•H2O. After being filtered and dried the weight of the hydrated salt’s mass was determined to be
0.2920 g. Calculate the weight percent of calcium in the sample of phosphate rock.
ex. An ore sample is to be analyzed for sulfur. As part of the procedure, the ore is dissolved and the sulfur is converted
to sulfate ions, SO4-2. Barium nitrate is added which causes the sulfate to precipitate as BaSO4. The original sample had a
mass of 3.187 g. The dried BaSO4 has a mass of 2.005g. What is the percent sulfur in the ore sample?
8
Section 4.8: Acid-Base Reactions
a)
Predicting Reactions Involving Acids and/or Bases
Formation of a gas: Gases may form directly in a double replacement reaction or can form from the decomposition of a
product such as H2CO3 or H2SO3. (Don't leave carbonic or sulfurous acids, or ammonium hydroxide as products of double
replacement reactions, decompose them!)
H2CO3 (aq) 
H2SO3 (aq) 
NH4OH (aq) 
Ex. Excess hydrochloric acid solution is added to a solution of potassium sulfite.
Ex. A solution of sodium hydroxide is added to a solution of ammonium chloride.
Neutralization: When a molecular substance such as water or acetic acid is formed, ions are removed from solution and
the reaction "works". Note: Depending on the volumes and concentrations of the reactants, you may not actually
produce a pH of 7 upon complete reaction. The excess concentration of ions will determine the pH. Only if there is NO
excess of either reactant will you get a pH of 7. This will change your net ionic equation.
Ex. Dilute solutions of lithium hydroxide and hydrobromic acid are mixed.
Ex. Equal volumes of 0.4 M calcium hydroxide and 0.2 M sulfuric acid are mixed.
Will the pH of the resulting solution be above 7, below 7, or neutral? Explain your answer.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
9
b) Volumetric analysis – or titration - is a widely used quantitative method of analysis in which the volume of one
reagent (the titrant), that is required to react completely with an exact amount of another reagent, is measured
precisely. The method presumes that the titration reaction goes to completion.
In most titrations the concentration of titrant is known, in which case the exact amount of analyte, the substance being
analyzed, can be found.
In other cases the exact amount of substance being titrated is known, from which the concentration of titrant can be
determined. This latter case is the basis for solution standardization. A standard solution is one that has a concentration
known precisely, often to four significant figures.
Generally, a solution is standardized by placing a portion of it in a buret and measuring the volume of it required to
react with an accurately and precisely weighed sample of primary standard. Primary standards are ultra-pure solid
compounds with high molecular weights and reliable
stability.
Once a solution is standardized, it may be used as a
secondary standard for determining the concentration
of other solutions with which it reacts. For example, KHP
(potassium hydrogen phthalate) is a solid primary
standard used to standardize sodium hydroxide (NaOH)
solutions. A standardized NaOH solution can then serve
as a secondary standard in determining the acid content
of other samples.
The point during the reaction where enough titrant has
been added to completely react all of the solution being
analyzed is called the equivalence point. The solution
should have an indicator added to it in order to signal
the endpoint. The endpoint is the volume of titrant
required to make the analyte turn color. If you select
the correct indicator, the endpoint will be at or near the
equivalence point.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
10
Technique: Preparing the Buret for Titration
1. Rinse a clean buret several times with 5 mL portions of distilled water and then with portions of
titrant.
2. Allow the titrant to drain through the buret stopcock so that the tip gets rinsed with titrant as well.
3. Discard the rinse solution in a waste beaker. Clamp the buret into place, and fill it with the titrant.
Remove air bubbles from the tip of the buret and the stopcock by draining several milliliters of titrant.
Dispose of the drained titrant in a waste beaker.
ex. How many mL of a 0.800 M NaOH solution is needed to neutralize 40.00 mL of a 0.600 M HCl solution?
ex. You want to determine the molar mass of an acid that contains one acidic hydrogen. You weigh out a
2.879 g sample of the pure acid and dissolve it in distilled water adding a few drops of phenolphthalein indicator in
the acid solution. You titrate the acid with 0.1704 M NaOH. Endpoint is reached after 42.55 mL of the base is added.
Calculate the molar mass of the acid.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
11
ex. A student carries out an experiment to standardize a sodium hydroxide solution. To do this the student weighs out a
1.3009-g sample of potassium hydrogen phthalate (KHC8H4O4). KHP (molar mass 204.22 g/mol) has one acidic
hydrogen. The student dissolves the KHP in distilled water, adds phenolphthalein as an indicator, and titrates the
resulting solution with the sodium hydroxide solution to the phenolphthalein endpoint. The difference between the
final and initial buret readings indicates that 41.20 mL of the sodium hydroxide solution is required to react exactly
with the 1.3009 g KHP. Calculate the concentration of the sodium hydroxide solution.
ex.
An environmental chemist analyzed the effluent (the released waste material) from an industrial process known to
produce the compounds carbon tetrachloride (CC14) and benzoic acid (HC7H5O2), a weak acid that has one acidic
hydrogen atom per molecule. A sample of this effluent weighing 0.3518 g was shaken with water, and the
resulting aqueous solution required 10.59 mL of 0.1546 M NaOH for neutralization. Calculate the mass percent of
HC7H5O2 in the original sample.
AP Problem
Concentrated sulfuric acid (18.4-molar H2SO4) has a density of 1.84 grams per milliliter. After dilution with water to 5.20molar, the solution has a density of 1.38 grams per milliliter and can be used as an electrolyte in lead storage batteries
for automobiles.
(a) Calculate the volume of concentrated acid required to prepare 1.00 liter of 5.20-molar H2SO4.
(b) Calculate the volume of 5.20-molar H2SO4 that can be completely neutralized with 10.5 grams of sodium
bicarbonate, NaHCO3.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
12
Section 4.9: Oxidation-Reduction Reactions
Oxidation reduction reactions involve a transfer of electrons.
Oxidation is a loss of electrons. Reduction is a gain of electrons.
Oxidizing agent: is the reactant that causes oxidation of another substance by taking electrons from the other substance.
The oxidizing agent gains electrons and therefore is reduced in the process.
Reducing agent: is the reactant that causes another substance to be reduced by giving the other substance some of its
own electrons and is therefore oxidized in the process.
LEO the lion goes GER
(loss of electrons is oxidation; gain of electrons is reduction)
OIL RIG
(oxidation is loss; reduction is gain)
The substance undergoing… oxidation
LOSES electron(s)
Gets its oxidation number INCREASED
Gets OXIDIZED
Acts as the REDUCING AGENT
a)
The substance undergoing… reduction
GAINS electron(s)
Gets its oxidation number DECREASED
Gets REDUCED
Acts as the OXIDIZING AGENT
Predicting Redox Reactions
Single Replacement: In a single replacement reaction, one element displaces another in a compound. One element is
oxidized and another is reduced.
A + BC  B + AC
Active nonmetals replace less active nonmetals from their compounds in aqueous solution. Each halogen will displace
heavier halogens from their binary salts. If a less active nonmetal tries to displace an active nonmetal, there will be no
reaction.
ACTIVITY SERIES OF NONMETALS :
Fluorine
Highest
Chlorine
Bromine
Iodine
Lowest
Ex. Chlorine gas is bubbled into a solution of potassium iodide.
Ex. Iodine crystals are dropped into potassium chloride solution.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
13
Active metals replace less active metals or hydrogen from their compounds in aqueous solution. Use an activity series to
determine the activity of metal, or a periodic table. Group I metals are the most active metals on the periodic table.
Activity of metals also decreases as you move down a group. (Just remember: gold and platinum are virtually
nonreactive.)
Ex. Magnesium turnings are added to a solution of iron (III) chloride.
Ex. Sodium is added to water. (exception – good one to know)
Simple Synthesis and Decomposition: Elements react together, usually under very high heat, to form ionic compounds.
Most elements exist as compounds naturally in the earth’s crust. Compounds decompose to form their elements, either
by extreme heat under a Bunsen burner flame or by a process called electrolysis. Electrolyte uses high voltage to
overcome the bond energy between atoms of the different elements. NOTE: You have to know the formulas of the
diatomic elements: H2, N2, O2, F2, Cl2, Br2, and I2.
Ex. Aluminum oxide ore is electrolyzed.
Ex. Calcium strips are reacted with nitrogen gas.
Combustion Reactions: In the presence of excess oxygen, most hydrocarbons and alcohols can undergo complete
combustion reactions to form carbon dioxide and water. The only tricky part is balancing the reaction. Balance the C
first, H second, and O third.
Ex. Propane burns in air.
Ex. Ethanol burns in excess oxygen.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
14
b) Assigning Oxidation Numbers
1. The oxidation number of an atom in its element form is 0.
2. The sum of the oxidation numbers for the atoms in a compound must be 0.
3. The sum of the oxidation numbers for the atoms in a polyatomic ion must equal the charge of the polyatomic ion.
4. The oxidation number for a monatomic ion is the same as the ion’s charge.
5. The oxidation number for metals and nonmetals in binary compounds are the same as assigned for writing formulas.
6. Metals with a single oxidation number will always have the same oxidation number in any compound. group 1 = +1,
group 2 = +2, Zn = +2, Al = +3 etc...
7. Oxygen will have the oxidation number of -2. The exception to this rule occurs in peroxides (O2-2), where each oxygen
is assigned a -1 oxidation number.
Fe(s) + Cu+2(aq) --> Fe+2 (aq) + Cu (s)
ex.
(a) write the oxidation numbers of each reactant and product.
(b) name the reactant that becomes oxidized.
(c) name the reactant that becomes reduced.
(d) name the reducing agent.
(e) name the oxidizing agent.
(f) write the reduction half-equation.
(g) write the oxidation half-equation.
(h) list all the lab observations you would expect to see as this reaction went to completion.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
15
c) Balancing Redox Equations by the Half-Reaction Method
Identify what is oxidized and what is reduced.
Break the reaction into two half reactions; one for the oxidation half-reaction and one for the reduction half-reaction.
In acidic solution, and for EACH half-reaction (ie. you do steps #1-4 twice):
1.
2.
3.
4.
For each half-reaction, balance all atoms, other than H and O.
For each half-reaction, balance O by using H2O (since all redox reactions occur in solution)
For each half-reaction, balance H by using H+
For each half-reaction, balance the charges by adding electrons, e-. Both sides have to have an equal charge to each
other.
-
5. If necessary, multiply one or both half-reactions by an integer to equalize the number of electrons transferred (e lost
-
= e gained)
6. Add the two half-reactions and cancel species that appear on both sides of the equation.
7. Check and make sure atoms and charge are balanced.
In basic solution
1. Do steps #1-7 from above.
+
-
2. For each H that’s left, add an OH to both sides of the equation.
+
-
3. On the side with H , add the H+ and the OH together to get H2O
4. These should algebraically cancel with the H2O that was already on the other side
5. Check and make sure everything’s balanced (atoms and charges)
ex. Balance each redox equation in acidic solution.
-
2-
2+
MnO4 (aq) + C2O4 (aq)  Mn (aq) + CO2(g) (acidic)
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
16
Cl2 (g)  Cl- (aq) + OCl- (aq)
** this is a disproportionation reaction – the reactant both oxidizes and reduces itself**
Sn2+ + Cr2O7 2-  Sn4+ + Cr3+ (basic solution)
MnO4 2-+ I-  MnO2 + I2 (basic solution)
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
17
d)
Oxidation-Reduction Reaction Stoichiometry (Titration)
The most common oxidizing agents are KMnO4 (purple) & K2Cr2O7 (orange).
When you titrate with permanganate, the solution is colorless until you use up all of the reducing agent (substance being
oxidized). It acts as its own indicator, meaning the purple will fade at the equivalence point to a colorless or very pale
pink solution.
In calculations, work redox titrations like acid-base titrations. You must have a balanced reaction to know the mole
ratio.
Ex. Use the half reaction method to write the balanced net ionic equation for the reaction between MnO4and Fe2+ ion in acid solution. The unbalanced equation is Fe2+(aq) + MnO4-(aq) ➔ Fe3+(aq) + Mn2+ (aq).
How many moles of Fe2+ ion can be oxidized by 0.012 moles MnO4- ion? Show work.
Ex . A solid sample containing some Fe2+ ion weighs 1.923 g. It requires 36.44 mL of 0.0244 M KMnO4 to titrate the
Fe2+ in the dissolved sample to a pink end point. Calculate the grams of iron in the sample and the percentage of
iron in the sample. Show work with units.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
18
Ex. Find the mass percent of copper in an alloy. The alloy is dissolved and excess I– reacts with the copper in this
(unbalanced) reaction: 2 Cu2+ + 5 I – ➔ 2 CuI + I3– . The I3– is then titrated (the I3– will have a deep blue violet if some
starch is present) with sodium thiosulfate, Na2S2O3, in this unbalanced reaction:
I3– + 2 S2O32– ➔ S4O62- + 3 I–.
If 26.23 mL of 0.10 M Na2S2O3 is required for the titration to the equivalence point, then what is the weight percent
of Cu in the 0.251 g alloy?
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000
19
AP test 1981 B A 1.2516 gram sample of a mixture of CaCO3 and Na2SO4 was analyzed by dissolving the sample and
completely precipitating the Ca2+ as CaC2O4. The CaC2O4 was dissolved in sulfuric acid and the resulting H2C2O4 was
titrated with a standard KMnO4 solution.
(a) Write the balanced equation for the titration reaction, shown unbalanced below.
MnO4- + H2C2O4 + H+ → Mn2+ + CO2 + H2O
Indicate which substance is the oxidizing agent and which substance is the reducing agent.
(b) The titration of the H2C2O4 obtained required 35.62 milliliters of 0.1092 M MnO4- solution. Calculate the number
of moles of H2C2O4that reacted with the MnO4-.
(c) Calculate the number of moles of CaCO3 in the original sample.
(d) Calculate the percentage by weight of CaCO3 in the original sample.
Rhonda Alexander, Robert E. Lee High School, Tyler, TX;
Chris Bires, Spring Valley High School, Spring Valley, OR; Educational Testing Service, Princeton, NJ;
George R. Hague and Jane D. Smith, The Ultimate Chemical Equations Handbook, © 2001 Flinn Scientific, Inc.,
Rene McCormick, AP Strategies, Dallas, TX; Mary Payton, Carrollton HS, Carrollton, TX;
Zumdahl, Chemistry, 5th ed. Houghton Mifflin 2000