Download Chapter 4 (Hill/Petrucci/McCreary/Perry Chemical Reactions in

Chapter 4 (Hill/Petrucci/McCreary/Perry
Chemical Reactions in Aqueous Solutions
This chapter deals with reactions that occur in aqueous solution …these solutions all use water as
the solvent. We will look at some properties of these solutions and also look briefly at three
different general types of reactions that occur in aqueous solutions.
“water is such a good solvent for so many ionic and molecular substances that it has been called
the universal solvent.” (Hill, p.125)
Electrical Properties of Aqueous Solutions
Nonelectrolytes do not conduct electricity – electrical conductance requires the presence of
charged particles.
Electrolytes do conduct electricity, in proportion to the concentrations of their ions in solution.
Strong electrolyte : almost all molecules or neutral “units” present form ions in aqueous solution
HCl(aq) + H2 O " H3 O+(aq) + Cl1-(aq)
(strong electrolyte)
100% ionized (converted to ions)
Weak electrolyte: relatively few molecules or neutral “units” present form ions in aqueous
solution
HC2 H3 O2 (l) + H2 O " H3 O+(aq) + C2 H3 O2 1- (aq)
(weak electrolyte)
<<100% ionized
Carefully read pp. 125-128 in Hill!
Molarities of Ions in Strong Electrolytes:
to calculate the molarity of an ion in a solution of strong electrolyte, simply multiply the
subscript for that ion in the compound by the given molarity of the electrolyte
See Example 4.1, Hill, p.129
See Exercise 4.1A, Hill, p.129
Reactions of Acids in Aqueous Solution
Recall: an acid is a proton donor in aqueous solution; a base is a proton acceptor in aqueous
solution.
Strong Acids Are Strong Electrolytes - Six strong acids that you should recognize:
HCl hydrochloric acid
HBr hydrobromic acid
HI hydroiodic acid
See also Table 4.1, p. 131, Hill
HNO3 nitric acid
H2 SO4 sulfuric acid
HClO 4 perchloric acid
Most of the other acids that you will encounter are weak acids that are weak electrolytes.
Ionization of Acids in Aqueous Solution
By convention, the chemical formulas for acids have their “ionizable” protons (H+ ions) at the
“front” of the formula.
Recall the strong acids:
HCl
HBr
HI
HNO3
H2 SO4
HClO4
hydrochloric acid
hydrobromic acid
hydroiodic acid
nitric acid
sulfuric acid
perchloric acid
1 ionizable proton
1 ionizable proton
1 ionizable proton
1 ionizable proton
2 ionizable protons
1 ionizable proton
Ionization of Acids in Aqueous Solution
0.10 M
0
0
HCl(aq) + H2 O " H3 O+(aq) + Cl1-(aq)
0.0 M
0.10 M
0.10 M
" means 100%
100% ionization in water = strong acid
0.10 M
0
0
HC2 H3 O2 (aq) + H2 O D H3 O+(aq) + C2 H3 O21-(aq)
0.09 M
0.01 M
0.01 M at equilibrium
D means <<100% ionization in water = weak acid
Reactions of Bases in Aqueous Solution
Recall: a base is a proton acceptor in aqueous solution.
Strong Bases Are Strong Electrolytes
Most strong bases are Group IA and Group IIA hydroxides:
IA:
IIA:
LiOH, NaOH, KOH, RbOH, CsOH
Mg(OH)2 , Ca(OH)2 , Sr(OH)2 , Ba(OH)2 ,
See also Table 4.1, p. 131, Hill
Ionization of Bases in Aqueous Solution
0.10 M
0
0
+
NaOH(aq) (+ H2 O) " Na (aq) + OH1-(aq)
0.0 M
0.10 M
0.10 M
100% ionization in water = strong base
0.10 M
0
0
+
NH3 (aq) + H2 O(l) D NH4 (aq) + OH1-(aq)
0.09 M
0.01 M
0.01 M
at equilibrium
<<100% ionization in water = weak base
Most common weak bases: NH3 and amines
Reaction of Acids with Bases: Neutralization
neutralization reaction: the reaction of ionizable H+ ions on acid molecules with OH1- or other
anions (such as HCO3 1- or CO3 2-) on base “molecules”
Example. We represent an acid-base reaction as a “molecular equation.” (no ions involved)
NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq)
But, underlying reaction: H+ + OH1- " H2 O(l)
Classically, acid-base reactions produce a salt and water.
Molecular Equations to Ionic Equations
ionic equation: all ionizable species written as ions, i.e. in their ionized or dissociated forms
Example. If the molecular equation is
NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq)
we must “break” up the aqueous ionizable species into their respective ions:
NaOH(aq) ionizes to Na+(aq) + OH1-(aq)
HCl(aq) ionizes in water to H3 O+(aq) + Cl1-(aq)
Molecular Equation: NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq)
NaOH(aq): Na+(aq) + OH1-(aq)
HCl(aq): H3 O+(aq) + Cl1-(aq)
NaCl(aq): ionizes to Na+(aq) + Cl1-(aq)
Corresponding Ionic Equation:
Na+(aq) + OH1-(aq) + H3 O+(aq) + Cl1-(aq) " H2 O(l) + Na+(aq) + Cl1-(aq)
] Species in bold that appear on both sides are called spectator ions and “cancel” out.
These species do not participate in the chemical reaction.
Molecular Equations to Ionic Equations
Molecular Equation: NaOH(aq) + HCl(aq) " H2 O(l) + NaCl(aq)
Ionic Equation: Na+(aq) + OH1-(aq) + H3 O+(aq) + Cl1-(aq) " H2 O(l) + Na+(aq) + Cl1-(aq)
Net Ionic Equation: H3 O+(aq) + OH1-(aq) " H2 O(l)
See Example 4.2 and following Exercises 4.2A and 4.2B on Hill, p. 133
Acid-Base Reactions That Form Gases
1. Carbonates (compounds that contain CO3 2-)
CaCO3 (s) + HCl(aq) " H2 O(l) + CaCl2 (aq) + CO2 (g)#
2. Sulfites (compounds that contain SO3 2-)
K2 SO3 (s) + H2 SO4 (aq) " H2 O(l) + K2 SO4 (aq) + SO2 (g)#
3. Sulfides (compounds that contain S2-)
Na2 S(aq) + 2 HCl(aq) " 2 NaCl(aq) + H2 S(g)#
Acid-Base Reactions: Another Example
When balancing acid-base equations that have hydroxyl bases, use the lowest common
denominator for the number of ionizable protons and the number of OH1- ions per base unit. Use
this number for the number of H2 O molecules formed.
Example. (unbalanced)
H3 PO4 (aq) + Ca(OH)2 (aq) " ???
2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + ??
2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + Ca2+(aq) + PO4 3-(aq)
2 H3 PO4 (aq) + 3 Ca(OH)2 (aq) " 6 H2 O(l) + Ca3 (PO4 )2 (aq)
Reactions That Form Precipitates
precipitate: a solid product formed from the reaction of two soluble ions (a cation and an anion);
a precipitate is, by definition, insoluble (not soluble) in the solvent used
Example of a Precipitation Reaction: Ba2+(aq) + SO4 2-(aq) " BaSO4 (s)
] The chemical equation above is the net ionic equation for the reaction between barium
chloride and sodium sulfate: BaCl2 (aq) + Na2 SO4 (aq) " 2NaCl(aq) + BaSO4 (s)
Can you get from this equation to the net ionic equation?
Reaction of Ag+ with I1The reaction: AgNO3 (aq) + KI(aq) " " AgI(s) + KNO3 (aq)
Net ionic reaction: Ag+(aq) + I1-(aq) " AgI(s) ]Silver iodide, AgI precipitates!
Solubility Rules for Common Ionic Compounds
1.Group IA ions and NH4 + are almost always SOLUBLE when paired with NO3 1-, C2 H3 O21- and
ClO 4 12. Most salts of Cl1-, Br1-, and I1- are SOLUBLE; exceptions are combinations of these anions
with Pb2+, Ag+ or Hg2 2+.
3. Compounds containing SO4 2- are SOLUBLE except those with Sr2+, Ba2+, Pb2+, and Hg2 2+;
CaSO4 is slightly soluble.
4. Compounds containing CO3 2-, OH1-, PO4 3- and S2- are INSOLUBLE except Group IA cations,
NH4 +; combinations of OH1- and S2- with Ca2+, Sr2+, Ba2+ are slightly to moderately soluble.
Memorize these solubility rules! See also Table 4.3 on p. 136 (Hill)
See Example 4.4 p. 137
Exercise 4.4A, p. 137
(a) MgSO4 (aq) + KOH(aq) " ?
(b) FeCl3 (aq) + Na2 S(aq) " ?
(c) Sr(NO3 )2 (aq) + Na2 SO4 (aq) " ?
Oxidation-Reduction Reactions (Redox Reactions)
In addition to acid-base and precipitation reactions, there is a third type of reaction: oxidationreduction or redox reactions.
Oxidation-reduction reactions are electron exchange reactions (electron = e1-)
Oxidation - loss of 1 or more electrons by an ion or molecule
Reduction - gain of 1 or more electrons by an ion/molecule
Example of an Oxidation Reaction.
Example of a Reduction Reaction.
Fe0 " Fe3+ + 3 e1Cl2 + 2 e1- " 2 Cl1-
Oxidation Numbers and the Oxidation Number Concept
oxidation number: the charge on a monatomic ion or the nominal charge on an atom in a unit of
a compound (oxidation number is also referred to as the oxidation state of an atom)
Short List of Oxidation Number Rules.
1. The oxidation number of a neutral, uncharged atom is 0
2. Ions: IA metals = +1; IIA metals = +2
3. Hydrogen : H is usually +1; sometimes -1 in hydrides
4. Oxygen: O is usually -2; sometimes -1 in peroxides
5. The sum of all the oxidation numbers in a molecule or an ion is equal to the charge on the
molecule (0) or ion.
Examples.
See Example 4.7, p. 141 (Hill)
Exercise 4.7 A, p. 142
Assign known oxidation numbers and then set sum of the oxidation numbers equal to the charge
and solve algebraically.
Al2 O3
P4
HAsO 4 3-
NaMnO 4
Oxidation Numbers in Nitrogen, Sulfur and Chlorine Species (see text Figure 4.12)
Identifying Oxidation and Reduction Reactions
To classify a reaction as an oxidation process or as a reduction process, first assign oxidation
numbers to all atoms on both sides of the equation.
1. Oxidation. If the oxidation number for an element increases (becomes more positive) from
reactant to product, the process is an oxidation process
2. Reduction. If the oxidation number for an element decreases (becomes more negative) from
reactant to product, the process is a reduction process
Identifying Oxidation and Reduction Reactions
Example. The “thermite” reaction: 2 Al(s) + Fe2 O3 (s) " 2 Fe(l) + Al2 O3 (s)
Here, Al0 " Al3+ (oxidation) and Fe3+ " Fe0 (reduction)
We say that Al was oxidized to Al3+ and that Fe3+ was reduced to Fe0
The species oxidized (Al) is the reducing agent, and the species reduced (Fe3+) is the oxidizing
agent.
Two Other Examples.
4 HCl(aq) + O2 (g) " 2 Cl2 (aq) + 2 H2 O(l)
Cl1- in HCl is oxidized to Cl0 in Cl2 , and O2 (O 0 ) is reduced to O2- in H2 O
Ag(s) + H+(aq) + NO3 1- " Ag+(aq) + H2 O(l) + NO(g)
Ag0 is oxidized to Ag+, and N+5 in NO3 1- is reduced to N+2 in NO
Oxidation-Reduction Reactions (Redox Reactions)
See Hill, Figure 4.14, pp. 146: HNO3 oxidizes Cu to Cu 2+, but HCl doesn’t … why?
Activity Series of the Metals
A metal will “displace from solution the ions of any metal that lie below it in the activity series.
Example.
metal = Mg and ion = Ni2+
Mg0 (s) + Ni2+(aq) " Mg2+(aq) + Ni0 (s)
Read Hill, Section 4.5, pp. 148-150
Figure 14.15. Maryland uses the Breathalyzer to determine blood alcohol levels of drivers.
Cr2 O7 2- + ethanol " Cr3+
Titrations
We are interested in quantitatively determining the concentration of a chemical species (called
the analyte) in a sample.
The sample is placed in a flask or beaker, and a solution containing a known concentration of a
chemical reagent (the titrant ) that will react with the analyte is added until no more analyte
remains (the titration endpoint).
The chemical reaction between the analyte and the titrant is known. We also know the
concentration of the titrant solution and the volume of the solution required to just react with all
of the analyte.
A titration is carried out using a tube ( a buret) calibrated along its length, typically in 0.1 mL
increments. The volume before titrant is measured and recorded; the volume after reaching the
endpoint is then measured and recorded.
The volume of titrant used is then:
Vrequired = Vfinal - Vinitial
Acid-Base Titrations
See Examples and Exercises on pp. 153-154 (Hill)
Precipitation Titrations
See Examples and Exercises on pp. 155-156 (Hill)
Redox Titrations
See Examples and Exercises on pp. 156-157 (Hill)