Transcript No Slide Title
Reactions in Aqueous Solution
Chapter 4
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A
solution
is a homogenous mixture of 2 or more substances The
solute
is(are) the substance(s) present in the smaller amount(s) The
solvent
is the substance present in the larger amount Solution Soft drink (
l
) Air (
g
) Soft Solder (
s
) Solvent H 2 O N 2 Pb Solute Sugar, CO 2 O 2 , Ar, CH 4 Sn aqueous solutions of KMnO 4 2
An
electrolyte
is a substance that, when dissolved in water, results in a solution that can conduct electricity.
weak electrolyte is a substance that, when dissolved in water, results in a solution that can conduct small electricity A
nonelectrolyte
is a substance that, when dissolved, results in a solution that does not conduct electricity.
nonelectrolyte weak electrolyte 3 strong electrolyte
Conduct electricity in solution?
Cations (+) and Anions (-) Strong Electrolyte – 100% dissociation NaCl (
s
) H 2 O Na + (
aq
) + Cl (
aq
) Weak Electrolyte – not completely dissociated CH 3 COOH CH 3 COO (
aq
) + H + (
aq
) 4
CH 3
Ionization
of acetic acid COOH CH 3 COO (
aq
) + H + (
aq
) A
reversible
reaction. The reaction can occur in both directions.
Acetic acid is a
weak electrolyte
because its ionization in water is incomplete.
5
Hydration
is the process in which an ion is surrounded by water molecules arranged in a specific manner.
d d + H 2 O 6
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution C 6 H 12 O 6 H 2 O (
s
) C 6 H 12 O 6 (
aq
) 7
Precipitation Reactions
Precipitate – insoluble solid that separates from solution precipitate PbI 2 Pb(NO 3 ) 2 (
aq
) + 2NaI (
aq
) PbI 2 (
s
)
molecular equation
+ 2NaNO 3 (
aq
) Pb 2+ + 2NO 3 + 2Na + + 2I PbI 2 (
s
) + 2Na + + 2NO 3 -
ionic equation
Pb 2+ + 2I PbI 2 (
s
)
net ionic equation
Na + and NO 3 are
spectator
ions 8
Solubility
is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
9
Solubility Rules
1. Group IA and ammonium compounds are soluble.
2. Acetates and nitrates are soluble.
3. Most chlorides, bromides, and iodides are soluble.
Exceptions: AgCl, Hg 2 Cl 2 , PbCl 2 ; AgBr, HgBr 2 , Hg 2 Br 2 , PbBr 2 ; AgI, HgI 2 , Hg 2 I 2 , PbI 2 4. Most sulfates are soluble.
Exceptions: CaSO 4 , SrSO 4 , BaSO 4 , Ag 2 SO 4 , Hg 2 SO 4 , PbSO 4 10
5. Most carbonates are insoluble.
Exceptions: Group IA carbonates and (NH 4 ) 2 SO 4 6. Most phosphates are insoluble.
Exceptions: Group IA phosphates and (NH 4 ) 3 PO 4 7. Most sulfides are insoluble.
Exceptions: Group IA sulfides and (NH 4 ) 2 S 8. Most hydroxides are insoluble.
Exceptions: Group IA hydroxides, Ca(OH) 2 , Sr(OH) 2 , Ba(OH) 2 11
Decide whether the following reaction occurs. If it does, write the molecular, ionic, and net ionic equations.
KBr + MgSO 4 1. Determine the product formulas: K + Mg and SO 4 2− 2+ and Br − make K 2 SO make MgBr 4 2 2. Determine whether the products are soluble: K 2 SO 4 is soluble MgBr 2 is soluble KBr + MgSO 4 no reaction 12
Examples of Insoluble Compounds CdS PbS Ni(OH) 2 Al(OH) 3 13
Writing Net Ionic Equations 1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes completely dissociated into cations and anions.
3. Cancel the spectator ions on both sides of the ionic equation 4. Check that charges and number of atoms are balanced in the net ionic equation Write the net ionic equation for the reaction of silver nitrate with sodium chloride.
AgNO 3 (
aq
) + NaCl (
aq
) AgCl (
s
) + NaNO 3 (
aq
) Ag + + NO 3 + Na + + Cl Ag + + Cl AgCl (
s
) + Na + + NO 3 AgCl (
s
) 14
Properties of Acids Have a sour taste. Vinegar owes its taste to acetic acid. Citrus fruits contain citric acid.
Cause color changes in plant dyes.
React with certain metals to produce hydrogen gas.
2HCl (
aq
) + Mg (
s
) MgCl 2 (
aq
) + H 2 (
g
) React with carbonates and bicarbonates to produce carbon dioxide gas 2HCl (
aq
) + CaCO 3 (
s
) CaCl 2 (
aq
) + CO 2 (
g
) + H 2 O (
l
) Aqueous acid solutions conduct electricity.
15
Properties of Bases Have a bitter taste.
Feel slippery. Many soaps contain bases.
Cause color changes in plant dyes.
Aqueous base solutions conduct electricity.
Examples: 16
Arrhenius acid is a substance that produces H + (H 3 O + ) in water Arrhenius base is a substance that produces OH in water 17
Hydronium ion
, hydrated proton, H 3 O + 18
In Figure A, a solution of HCl (a strong acid) illustrated on a molecular/ionic level, shows the acid as all ions.
In Figure B, a solution of HF (a weak acid) also illustrated on a molecular/ionic level, shows mostly molecules with very few ions.
19
A
Br ønsted acid
is a proton donor A
Br ønsted base
is a proton acceptor base acid acid base A Br ønsted acid must contain at least one ionizable proton!
20
Household Acids and Bases 4 | 21 21
Monoprotic
acids HCl H + + Cl HNO 3 H + + NO 3 CH 3 COOH H + + CH 3 COO Strong electrolyte, strong acid Strong electrolyte, strong acid Weak electrolyte, weak acid
Polyprotic Acid
An acid that results in two or more acidic hydrogens per molecule.
Diprotic
acids H 2 SO 4 H + + HSO 4 Strong electrolyte, strong acid HSO 4 H + + SO 4 2-
Triprotic
acids H 3 PO 4 H 2 PO 4 HPO 4 2 H + + H 2 PO 4 H + H + + HPO 4 2 + PO 4 3 Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid Weak electrolyte, weak acid 22
23
Identify each of the following species as a Br ønsted acid, base, or both. (a) HI, (b) CH 3 COO , (c) H 2 PO 4 HI (
aq
) H + (
aq
) + I (
aq
) Br ønsted acid CH 3 COO (
aq
) + H + (
aq
) CH 3 COOH (
aq
) Br ønsted base H 2 PO 4 (
aq
) H + (
aq
) + HPO 4 2 (
aq
) H 2 PO 4 (
aq
) + H + (
aq
) H 3 PO 4 (
aq
) Br ønsted acid Br ønsted base 24
Neutralization Reaction
acid + base salt + water HCl (
aq
) + NaOH (
aq
) NaCl (
aq
) + H 2 O H + + Cl + Na + + OH Na + + Cl + H 2 O H + + OH H 2 O 25
Neutralization Reaction Involving a Weak Electrolyte weak acid + base salt + water HCN (
aq
) + NaOH (
aq
) NaCN (
aq
) + H 2 O HCN + Na + + OH Na + + CN + H 2 O HCN + OH CN + H 2 O 26
Neutralization Reaction Producing a Gas acid + base salt + water + CO 2 2HCl (
aq
) + Na 2 CO 3 (
aq
) 2NaCl (
aq
) + H 2 O +CO 2 2H + + 2Cl + 2Na + + CO 3 2 2H + + CO 3 2 2Na + + 2Cl + H 2 O + CO 2 H 2 O + CO 2 27
Acid−Base Reaction with Gas Formation
Some salts, when treated with an acid, produce a gas. Typically sulfides, sulfites, and carbonates behave in this way producing hydrogen sulfide, sulfur trioxide, and carbon dioxide, respectively.
The photo to the right shows baking soda (sodium hydrogen carbonate) reacting with acetic acid in vinegar to give bubbles of carbon dioxide.
4 | 28 28
Oxidation-Reduction Reactions
(electron transfer reactions) 2Mg 2Mg 2+ + 4e -
Oxidation
half-reaction (lose e ) O 2 + 4e 2O 2 2Mg + O 2 + 4e -
Reduction
2Mg + O 2 2Mg 2+ half-reaction (gain e ) + 2O 2MgO 2 + 4e 29
Zn (
s
) + CuSO 4 (
aq
) ZnSO 4 (
aq
) + Cu (
s
) Zn Zn 2+ + 2e Zn is oxidized Zn is the
reducing agent
Cu 2+ + 2e Cu Cu 2+ is reduced Cu 2+ is the
oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.
What is the oxidizing agent in the reaction?
Cu (
s
) + 2AgNO 3 (
aq
) Cu(NO 3 ) 2 (
aq
) + 2Ag (
s
) Cu Cu 2+ + 2e Ag + + 1e Ag Ag + is reduced Ag + is the oxidizing agent 30
Oxidation number
The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H 2 , O 2 , P 4 = 0 2. In monatomic ions, the oxidation number is equal to the charge on the ion.
Li + , Li = +1 ; Fe 3+ , Fe = +3 ; O 2 , O = -2 3. The oxidation number of oxygen is and O 2 2 it is –1 .
usually
–2 . In H 2 O 2 31 4.4
4. The oxidation number of hydrogen is +1
except
when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1 .
5. Group IA metals are +1 , IIA metals are +2 always –1 .
and fluorine is 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion.
7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O 2 , is –½ .
HCO 3 What are the oxidation numbers of all the elements in HCO 3 ?
O = – 2 3x( – 2) H = + 1 + ?
+1 = – 1 C = +4 32
The Oxidation Numbers of Elements in their Compounds 33
What are the oxidation numbers of all the elements in each of these compounds?
NaIO 3 IF 7 K 2 Cr 2 O 7 KIO 3 K = +1 O = -2 3x( -2 ) + 1 + ?
= 0 I = +5 Na 2 Cr 2 O 7 O = -2 Na = +1 7x( -2 ) + 2x( +1 ) + 2x( ?) = 0 Cr = +6 IF 7 F = -1 7x( -1 ) + ?
= 0 I = +7 34
Potassium permanganate, KMnO 4 , is a purple−colored compound; potassium manganate, K 2 MnO 4 , is a green−colored compound. Obtain the oxidation numbers of the manganese in these compounds.
K Mn
1(+1) + 1(oxidation number of Mn) + 4(−2) = 0 1 + 1(oxidation number of Mn) + (−8) = 0 (−7) + (oxidation number of Mn) = 0
O
Oxidation number of Mn = +7 35
K Mn O
2(+1) + 1(oxidation number of Mn) + 4(−2) = 0 2 + 1(oxidation number of Mn) + (−8) = 0 (−6) + (oxidation number of Mn) = 0 Oxidation number of Mn = +6 In KMnO 4 , the oxidation number of Mn is +7.
In K 2 MnO 4 , the oxidation number of Mn is +6.
36
What is the oxidation number of Cr in dichromate, Cr 2 O 7 2− ?
Cr O
2(oxidation number of Cr) + 7(−2) = −2 2(oxidation number of Cr) + (−14) = −2 2(oxidation number of Cr) = +12 Oxidation number of Cr = +6 37
Oxidation
The half−reaction in which there is a loss of electrons by a species (or an increase in oxidation number).
Reduction
The half−reaction in which there is a gain of electrons by a species (or a decrease in oxidation number).
38
Oxidizing Agent
A species that oxidizes another species; it is itself reduced.
Reducing Agent
A species that reduces another species; it is itself oxidized.
39
Types of Oxidation-Reduction Reactions
Combination Reaction
A reaction in which two substances combine to form a third substance. A + B C 0 0 2Al + 3Br 2 +3 -1 2AlBr 3 40
For example: 2Na(
s
) + Cl 2 (
g
) 2NaCl(
s
) 41
Decomposition Reaction
C A + B +1 +5 -2 2KClO 3 +1 -1 0 2KCl + 3O 2 A reaction in which a single compound reacts to give two or more substances.
For example: 2HgO(
s
) 2Hg(
l
) + O 2 (
g
) 42
Types of Oxidation-Reduction Reactions
Combustion Reaction
A reaction in which a substance reacts with oxygen, usually with the rapid release of heat to produce a flame.
A + O 2 0 0 S + O 2 B +4 -2 SO 2 0 0 2Mg + O 2 +2 -2 2MgO 43
Combustion Reaction
For example: 4Fe(
s
) + 3O 2 (
g
) 2Fe 2 O 3 (
s
) 4 | 44 44
Types of Oxidation-Reduction Reactions
Displacement Reaction
A reaction in which an element reacts with a compound, displacing another element from it.
A + BC AC + B 0 +1 +2 Sr + 2H 2 O Sr(OH) 2 0 + H 2 Hydrogen Displacement +4 TiCl 4 0 0 +2 + 2Mg Ti + 2MgCl 2 Metal Displacement 0 Cl 2 -1 -1 0 + 2KBr 2KCl + Br 2 Halogen Displacement 45
Displacement Reaction
For example: Zn(
s
) + 2HCl(
aq
) H 2 (
g
) + ZnCl 2 (
aq
) 4 | 46 46
Types of Oxidation-Reduction Reactions
Disproportionation Reaction
The same element is simultaneously oxidized and reduced.
Example: reduced 0 Cl 2 + 2OH oxidized +1 ClO -1 + Cl + H 2 O 47
Classify each of the following reactions.
Ca 2+ + CO 3 2 NH 3 + H + CaCO 3 NH 4 + Precipitation Acid-Base Zn + 2HCl ZnCl 2 + H 2 Redox (H 2 Displacement) Ca + F 2 CaF 2 Mg(
s
) Mg 2+ (
aq
) + 2e − Fe 3+ (
aq
) + 3e − Fe(
s
) Redox (Combination) (oxidation) (reduction) 48
Solution Stoichiometry
The
concentration
of a solution is the amount of solute present in a given quantity of solvent or solution.
moles of solute
M
=
molarity
= liters of solution
Molality= amount of substance in mol of solute/mass in Kg of the solvent
What mass of KI is required to make 500. mL of a 2.80
M
KI solution?
volume of KI solution
M
KI moles KI
M
KI grams KI 1 L 500. mL x 1000 mL x 2.80 mol KI x 1 L soln 166 g KI 1 mol KI = 232 g KI 49
To prepare a solution, add the measured amount of solute to a volumetric flask, then add water to bring the solution to the mark on the flask.
50
Preparing a Solution of Known Concentration 51
Dilution
is the procedure for preparing a less concentrated solution from a more concentrated solution.
Dilution Add Solvent Moles of solute before dilution (i)
M
i V i
=
= Moles of solute after dilution (f)
M
f V f 52
Diluting a solution quantitatively requires specific glassware.
The photo at the right shows a volumetric flask used in dilution. 53
You place a 1.52−g of potassium dichromate, K 2 Cr 2 O 7 , into a 50.0−mL volumetric flask. You then add water to bring the solution up to the mark on the neck of the flask. What is the molarity of K 2 Cr 2 O 7 in the solution?
Molar mass of K 2 Cr 2 O 7 is 294 g.
1.52 g 1 mol 294 g 3 50.0 10 L 0.103
M
54
A saturated stock solution of NaCl is 6.00
M
. How much of this stock solution is needed to prepare 1.00−L of physiological saline soluiton (0.154
M
)?
M
i
V
i
M
f
V
f
V
i
M
f
V
f
M
i
V
i
V
i (0.154
M
)(1.00
6.00
M
0.0257
L or L) 25.7
mL 55
How would you prepare 30.0 mL of 0.100
M
HNO 3 from a stock solution of 2.00
M
HNO 3 ?
M
i V i =
M
f V f
M
i = 2.00
M M
f = 0.100
M
V f = 0.0300 L V i = ? L V i =
M
f V f
M
i = 0.100
M
x 0.0300 L 2.00
M
= 0.00150 L = 1.50 mL Dilute 1.50 mL of acid with water to a total volume of 30.0 mL.
56