AP Chemistry
Unit #4 (Key)
Chapter 4 – Zumdahl & Zumdahl
Types of Chemical Reactions
& Solution Stoichiometry
Students should be able to:
Predict to some extent whether a substance will be a strong electrolyte, weak electrolyte, or nonelectrolyte.
Predict the ions that an electrolyte dissociates into.
Identify substances as acids, bases, and salts.
Using solubility rules, predict if a precipitate forms in a metathesis reaction. Next, predict its products and write a balanced
equation.
Predict the products and write a balanced chemical equation for neutralization reactions.
After constructing molecular reactions for metathesis reactions, be able to identify spectator ions and write the net ionic
equations.
Assign oxidation numbers to atoms.
Determine whether a reaction is Redox (single replacement) or not.
Use the activity series to predict whether a Redox reaction will occur and be able to write the molecular and net ionic equations if
it does.
Calculate moles of solute, volume of solution, or Molarity of the solution from the other two.
Recognize and work dilution problems.
Calculate the volume of a certain molarity solution required to react with another solution of known molarity.
Calculate the mass of a substance that would be required to react with a given volume of a solution of known molarity.
Calculate mass of solute or concentration of an unknown solution from titration data.
Keywords:
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concentration
titration
indicators
solvent
strong electrolyte
precipitate
molecular equation
spectator ions
salts
reduction
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molarity
standard solution
aqueous
electrolyte
weak electrolyte
solubility
(complete) ionic equation
acids
neutralization
redox reaction
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dilution
equivalence point
solute
nonelectrolyte
activity series
metathesis
net ionic equation
bases
oxidation
oxidation number
I.
Aqueous Solutions
A. What is an aqueous solution?
• Solutions in which the solvent is water.
B. What is the solvent?
• It’s the substance that does the dissolving—the dissolving medium
Ex: Water is the universal solvent
C. What is the solute?
• It’s the substance that gets dissolved into liquid form—goes from solid to
aqueous
D. What does concentration mean?
• The amounts of chemicals present in solutions—expressed in terms of
MOLARITY (M) during this unit
E. How do you measure molarity (M)?
M = Molarity =
_____Moles____
liters of solution
F. Sample Exercise 4.1 – Calculate the molarity of a solution made by dissolving 23.4 g of
sodium sulfate, Na2SO4, in enough water to form 125 mL.
G. Sample Exercise 4.2 – How many grams of Na2SO4 are required to make 0.350 L of
0.500 M Na2SO4?
2
H. How do you make a dilution?
1. What is the formula that you can use?
I. Sample Exercise 4.7 p. 145
II. Electrolytes
A. Electrolyte vs. Nonelectrolyte
Electrolyte – solutes that exists as ions (charged particles) in
solution. These solutions are capable of conducting electricity.
Example soluble ionic compounds.
Nonelectrolyte – solute that does not conduct electricity because
it does not ionize in solution.
Example many organic substances (sugars)
B. Strong vs. Weak Electrolytes
Strong Electrolytes ionize completely and are very soluble solutes.
Weak Electrolytes do not have complete ionization in solution.
1. What happens when an ionic substance dissolves?
IONIZATION
H 2O
NaCl (s)
+
–
Na (aq) + Cl (aq)
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C. Sample Exercise 4.4 – What are the molar concentrations of all ions present in a 0.025
M aqueous solution of calcium nitrate?
III.
Acids, Bases, and Salts
A. What is an Acid? Substances that are able to ionize in solution to
form a hydrogen ion (H+ (aq))
1. What is the difference between a Monoprotic Acid and a Diprotic Acid?
Monoprotic Acids = 1 H+ ion added per compound
Diprotic Acids = 2 H+ ions added per compound
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2. What is the difference between a strong acid and a weak acid?
Strong Acids Ionize completely (are Strong Electrolytes), Weak
Acids incompletely ionize.
3. What are the strong acids?
HClO3, HBr, HCl, HI, HNO3, HClO4, H2SO4
B. What is a Base?
Substances that can react with or accept H+ ion.
Ex) OH – (aq) + H + (a) H2O (l)
1. What types of compounds make strong bases?
Group 1 Metal Hydroxides (Ex. = NaOH)
Group 2 Heavy Metal Hyrdoxides (Ex. = Sr(OH)2)
C. What are Salts?
Ionic Compounds that are not acids or bases. Examples: NaCl,
AlCl3, Ca(NO3)2
*Salts do NOT produce H+ or OH- when dissolved in H2O
D. What is a neutralization reaction?
Acid + Base Salt + Water
E. Sample Exercise 4.12 (p. 159) – What volume of a 0.100 M HCl solution is
needed to neutralize 25.0 mL of 0.350 M NaOH?
F. Sample Exercise 4.13 (p. 160) – In a certain experiment, 28.0 mL of 0.250 M
HNO3 and 53.0 mL of 0.320 M KOH are mixed. Calculate the amount of water
formed in the resulting reaction. What is the concentration of H+ or OH- ions
in excess after the reaction goes to completion?
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IV.
Ionic Equations
A. Spectator Ions – Ions that don’t change during a reaction. These Ions are
normally left out of reactions.
B. Net Ionic Equations – A Type of equation where the spectors ions are omitted
from the complete ionic equation.
C. Sample Exercise 4.9 (p. 155) – For each of the following reactions, write the
molecular equation, the complete ionic equation, and the net ionic equation.
a. Aqueous potassium chloride is added to aqueous silver nitrate to form a
silver chloride precipitate plus aqueous potassium nitrate.
b. Aqueous potassium hydroxide is mixed with aqueous iron(III) nitrate to
form a precipitate of iron(III) hydroxide and aqueous potassium nitrate.
V.
Metathesis Reactions
A. What is a Metathesis Reaction? A reaction that involves an exchange of ion
partners. AKA - Double Displacement Rxn, & Double Replacement Rxn
X + B
+ BX
B. What are the driving forces for a metathesis reaction?
1. The formation of an insoluble solid (a precipitate) in the mixing solutions.
2. The formation of either a soluble weak electrolyte or a soluble
nonelectrolyte.
3. The formation of a gas that escapes from the solution.
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C. Precipitation Reactions - any double replacement (metathesis) rxn that
produces an insoluble solid.
1. Precipitate – an insoluble solid formed in a solution.
Ex) 2KI (aq) + Pb(NO3)2 (aq) 2 KNO3 (aq) + PbI2 (s)
2. Solubility – the measured amount that a particular substance can
dissolve in a given quantity of solvent.
* Any substance whose solubility is less than 0.01 mol/L will be
considered insoluble.
D. Solubility Rules
SOLUBLE SALTS
Group I compounds and ammonium
compounds
Nitrates, hydrogen carbonates and chlorates
Chlorides, bromides and iodides
2+
+
2+
(EXCEPT those of Pb , Ag and Hg2 )
Sulfates
+
2+
2+
2+
2+
(EXCEPT Ag , Sr , Ba , Pb and Ca )
E.
INSOLUBLE SALTS
Hydroxides
(EXCEPT Group I and ammonium, hydroxides
2+
2+
2+
of Ca , Sr and Ba are slightly soluble)
Carbonates, phosphates, chromates and
sulfides
(EXCEPT group I and ammonium salts,
sulfides of group II are soluble)
Sample Exercise 4.10 (p. 156) – Determining the Mass of Product Formed.
Calculate the mass of solid NaCl that must be added to 1.50 L of a 0.100 M AgNO3
solution to precipitate all the Ag+ ions in the form of AgCl.
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F.
Reactions in which a weak electrolyte or nonelectrolyte forms:
• Strong elecrolytes (100% dissociation) =
1. Strong acids (memorize all 7)
2. Strong bases (memorize)
3. All ionic compounds (metal & nonmetal)
• Weak elecrolytes (less than 100% dissociation) =
4. Weak acids (any that aren’t part of the 7)
5. Weak bases (for now, just NH3)
6. NO ionic compounds
G.
Reactions in which a Gas forms:
•
methesis reaction may lead to the formation of a neutral molecule
that has low boiling point as well as low solubility in water. Thus, a
gas is formed. For example:
2 H+ + CO32- = H2O + CO2(g)
H.
Sample Exercise 4.9 – Write balanced complete ionic and net ionic equations for any
reactions that occur when the following compounds are mixed: (a) Cr(C2H3O2)2 (aq)
and HNO3 (aq) (b) FeCO3 (s) and HCl (aq) (c) PbS (s) and H2SO4 (aq).
VI.
Reactions of Metals
A. Oxidation and Reduction
1. Oxidation – The loss of electrons by a substance during a rxn.
2. Reduction – The gain of electrons by a substance during a rxn.
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“LEO says GER” or “OIL RIG”
Ca (s) + 2 H+ (aq) Ca2+ (aq) + H2 (g)
B. Oxidation of Metals by Acids and Salts
Metal + Acid Salt + Hydrogen Gas
Fe (s) + H2SO4 (aq) FeSO4 (aq) + H2 (g)
C. Sample Exercise 4.10 – Write the balanced molecular and net ionic equations for the
reaction of aluminum with hydrobromic acid.
D. The Activity Series – A list of decreasing ease of Oxidation
Metal
React with Acid?
React with Steam
Li
K
Ca
Na
Mg
Al
Zn
Fe
Sn
Pb
H
Cu
Ag
Pt
Au
YES
YES
YES
YES
YES
YES
YES
YES
YES
YES
NO
NO
NO
NO
YES
YES
YES
YES
YES
YES
YES
YES
NO
NO
NO
NO
NO
NO
NO
React with Cold
Water?
YES
YES
YES
YES
NO
NO
NO
NO
NO
NO
NO
NO
NO
NO
NO
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E. Sample Exercise 4.11 – Will an aqueous solution of iron (II) chloride oxidize
magnesium metal? If so, write the balanced molecular and net ionic equations for the
reaction.
VII.
Solution Stoichiometry
A. How to Solve Solution Stoichiometry Problems:
Laboratory Units
Chemical Units
Laboratory Units
Grams of
Substance
A
Moles of
Substance
A
Volume or
Molarity of
Substance
A
Grams of
Substance
B
Moles of
Substance
B
Volume or
Molarity of
Substance
B
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B. Sample Exercise 4.12 – How many moles of H2O form when 25.0 mL of 0.100 M HNO3
solution is completely neutralized by NaOH?
C. Titrations – A Laboratory method that is used to find the
concentration of an unknown solution by undergoing a specific
chemical reaction of known stoichiometric proportions with a
solution of known concentration.
1. Standard Solutions – The solution of known concentration in a
titration.
2. Equivalence Point – The point at which stoichiometric proportions
are reached.
Example) When a pink color appears during an acid –
base titration
3. Indicators – substances that are used to identify when the
equivalence point has been reached.
Example) Phenolphthalein
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D. Sample Exercise 4.13 – One method used commercially to peel potatoes is to soak
them in a solution of NaOH for a short time, remove them from the NaOH, and spray off
the peel. The concentration of NaOH is normally in the range 3 to 6 M. The NaOH is
analyzed periodically. In one such analysis, 45.7 mL of 0.500 M H2SO4 is required to react
completely with a 20.0 mL sample of NaOH solution:
H2SO4 (aq) + 2 NaOH (aq) 2 H2O (l) + Na2SO4 (aq)
What is the concentration of the NaOH solution?
E. Sample Exercise 4.13 – The quantity of Cl – in a water supply is determined by
titrating the sample with Ag
Ag + (aq) + Cl
–
+
:
(aq) AgCl (s)
What mass of chloride ion is present in 10.0 g sample of the water if 20.2 mL of 0.100
M Ag + is required to react with all the chloride in the sample?
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VIII. Rules for assigning OXIDATION STATES (numbers):
A. UNCOMBINED ELEMENTS (ELEMENTS NOT BONDED TO ANY OTHER
TYPE OF ELEMENT) have an oxidation number of ZERO. This includes any
formula that has only one chemical symbol in it (single elements &
diatomic elements).
Examples:
Al(s)0
Na(s)0
Cl20
H20
B. In COMPOUNDS (remember, they are neutral and have 2+ different
elements bonded together), the sum of the CHARGES must ADD UP TO
ZERO so the ions within a compound have oxidation numbers equal to
their OXIDATION # FOUND ON PERIODIC TABLE/INDIVIDUAL
CHARGES.
Ex:
NaCl
Ex:
Na: 1(+1) = +1
Cl: 1(-1) = -1
0!
Mg3N2
Mg: 3(+2) = +6
N: 2(-3) = -6
0!
Ex: HNO3
H: 1(+1) = +1
N: 1(+5) = +5
O: 3(-2) = -6
0!
* The OXIDATION NUMBER is the number INSIDE the PARENTHESES.
It is the charge on JUST ONE atom of that element!
** Remember that we almost always write the + ION FIRST and the - ION
LAST in a compound formula.
EXAMPLE:
HCl
EXCEPTION to this rule:
NH3
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C. GROUP 1 METALS always have an oxidation number of +1 when in a
compound (bonded to another species). Likewise, combined GROUP 2
METALS always therefore have a +2 oxidation number when located
within a compound.
Ex:
LiCl
MgCl2
D. FLUORINE is always a -1 in compounds. The other HALOGENS (ex: Cl,
Br) are also -1 as long as they are the most electronegative element in the
compound.
Ex:
HF
CaCl2
NaBr
E. HYDROGEN is a +1 in compounds unless it is combined with a metal (and is
at the back of the formula), then it is -1.
Ex:
HCl
LiH
F. OXYGEN is USUALLY -2 in compounds.
Ex:
H2O
When combined with fluorine (F), which is more electronegative,
oxygen is +2.
Ex:
OF2
When in a PEROXIDE oxygen is -1. A peroxide is a compound that has a
formula of X2O2.
Ex:
Na2O2
H2O2
G. The sum of the oxidation numbers in polyatomic ions must equal the
CHARGE ON THE ION (SEE TABLE E).
Ex: Cr2O72-
Cr: 2(+6) = +12
O: 7(-2) = -14
-2
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IX. Balancing Oxidation-Reduction Reactions
A. Half – Reactions
Sn2+ (aq) + 2Fe3+ (aq) Sn4+ (aq) + 2Fe2+ (aq)
red. ½ rxn:
ox. ½ rxn:
B. Balancing Redox Equations by the Half Reaction Method (in acidic
solution)
1. Divide the equation into 2 incomplete half-reactions, one for oxidation, one for reduction.
2. Balance each half-reaction.
(a) First, balance the elements other than H and O.
(b) Next, balance the O atoms by adding H2O.
(c) Then, balance the H atoms by adding H+
(d) Finally, balance the charge by adding e – to the side with the greater positive charge.
3. Multiply each half-reaction by an integer so that the number of electrons lost in the
oxidation half-reaction is equal to the number of electrons gained by the reduction halfreaction.
4. Add the two half-reactions and simplify where possible by canceling species appearing on
both sides of the equation.
5. Check the equation to make sure that there are the same number of atoms of each kind
and the same total charge on both sides.
Sample Exercise 4.19 – Potassium dichromate (K2Cr2O7) is a bright orange compound
that can be reduced to a blue-violet solution of Cr+3 ions. Under certain conditions,
K2Cr2O7 reacts with ethyl alcohol (C2H5OH) as follows:
H+(aq) + Cr2O72– (aq) + C2H5OH(l)
Acidic
Cr3+ (aq) + CO2(g) + H2O(l)
Problem #63 p. 184
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C. Balancing Equations for reactions occuring in Basic Solutions
Same as acidic solution except no excess H+ can remain, so:
6. Consume any excess H+ by adding OH- to that side, making water.
Sample Exercise 4.20 – Silver is sometimes found in nature as large nuggets;
more often it is found mixed with other metals and their ores. An aqueous
solution containing cyanide ion is often used to extract the silver using the
following reaction that occurs in basic solution:
Ag(s) + CN –(aq) + O2(g)
Basic
Ag(CN)2-(aq)
Problem #65 p. 184
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