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Chapter 4
Reactions in Aqueous Solution
Copyright © The McGraw-Hill Companies, Inc. Permission required for reproduction or display.
Describing Chemical Reactions
A chemical reaction is the process by which one or more substances are changed into one or more different substances
(NH4)2Cr2O7(s) N2(g) + Cr2O3(s) + 4H2O(g)
“The reactant ammonium dichromate yields the products nitrogen, chromium (III) oxide and water”
A CHEMICAL EQUATION represents, with symbols and formulas, the identifies and relative amounts of the reactants and products in a chemical equation
reactants
products
2
Indications of a Chemical Reaction1. Evolution of heat and light is strong evidence
that a chemical reaction has taken place! But, the evolution of heat or light by itself is not necessarily a sign of a chemical change since many physical changes
also release either heat or light. 2. Production of gas! (aka bubbles when two substances
are mixed)
3. Formation of precipitate! A solid that is produced as a result of a chemical reaction in solution and that separates from the solution is known as a precipitate
4. Color Change!
3
Characteristics of Chemical Equations1. The equation must represent all reactants and
products.2. The equation must contain the correct formulas for
the reactants and products3. The law of conservation of mass MUST be satisfied!!
Law of conservation of mass – atoms are neither created nor destroyed in ordinary chemical reactions
To equalized numbers of atoms, coefficients are added in front of the formulas where necessary
4
Types of Chemical ReactionsA + B ↔ CC ↔ A + B
A + BC ↔ AC + BAB + CD ↔ AD + CB
Acid + Base ↔ salt + water
Change of oxidation stateHydrocarbon + O2 ↔ CO2
+ H2O
5
Synthesis (Combination)DecompositionSingle ReplacementPrecipitation Reactions
(Double Replacement Reactions)
Neutralization Reactions (Acid/Base)
Redox ReactionsCombustion
PRECIPITATION REACTIONS
Double Replacement ReactionsAB + CD ↔ AD + BC
6
7
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 Solvent Solute
Soft drink (l)
Air (g)
Soft Solder (s)
H2O
N2
Pb
Sugar, CO2
O2, Ar, CH4
Snaqueous solutions
of KMnO4
8
An electrolyte is a substance that, when dissolved in water, results in a solution that can conduct electricity.
A nonelectrolyte is a substance that, when dissolved, results in a solution that does not conduct electricity.
nonelectrolyte weak electrolyte strong electrolyte
9
Strong Electrolyte – 100% dissociation
NaCl (s) Na+ (aq) + Cl- (aq)H2O
Weak Electrolyte – not completely dissociated
CH3COOH CH3COO- (aq) + H+ (aq)
Conduct electricity in solution?
Cations (+) and Anions (-)
10
Ionization of acetic acid
CH3COOH CH3COO- (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.
11
Hydration is the process in which an ion is surrounded by water molecules arranged in a specific manner.
H2O
12
Nonelectrolyte does not conduct electricity?
No cations (+) and anions (-) in solution
C6H12O6 (s) C6H12O6 (aq)H2O
Precipitation Reactions
13
Precipitate – insoluble solid that separates from solution
molecular equation
ionic equation
net ionic equation
Pb2+ + 2NO3- + 2K+ + 2I- PbI2 (s) + 2K+ + 2NO3
-
K+ and NO3- are spectator ions
PbI2
Pb(NO3)2 (aq) + 2KI (aq) PbI2 (s) + 2KNO3 (aq)
precipitate
Pb2+ + 2I- PbI2 (s)
Precipitation of Lead Iodide
14
PbI2Pb2+ + 2I- PbI2 (s)
15
Solubility is the maximum amount of solute that will dissolve in a given quantity of solvent at a specific temperature.
16
Examples of Insoluble Compounds
CdS PbS Ni(OH)2 Al(OH)3
Problem 4.20
17
Characterize the following compounds as (a)soluble or (b) insoluble in water:1.CaCO3
2.ZnSO4
3.Hg(NO3)2
4.HgSO4
5.NH4ClO4
Writing Net Ionic Equations
18
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
AgNO3 (aq) + NaCl (aq) AgCl (s) + NaNO3 (aq)
Ag+ (aq) + NO3
- (aq) + Na+
(aq) + Cl- (aq) AgCl (s) + Na+ (aq)
+ NO3- (aq)
Ag+ (aq) + Cl- (aq) AgCl (s)
Write the net ionic equation for the reaction of silver nitrate with sodium chloride.
19
Predict what happens when a potassium hydroxide solution is mixed with a solution of sodium chloride. Write a net ionic equation for the reaction.
20
What do we expect to see if we add copper (II) sulfate to sodium hydroxide? Write the molecular equation, ionic equation, and net ionic equation.
21
What do we expect to see if we add copper (II) sulfate to sodium hydroxide? Write the molecular equation, ionic equation, and net ionic equation.
22
Predict what happens when a potassium phosphate solution is mixed with a solution of calcium nitrate. Write a net ionic equation for the reaction.
EXTRA PRACTICE
23
Predict what happens when a silver nitrate solution is mixed with a solution of potassium hydroxide. Write a net ionic equation for the reaction.
Types of Chemical ReactionsA + B ↔ CC ↔ A + B
A + BC ↔ AC + BAB + CD ↔ AD + BC
Acid + Base ↔ salt + water
Change of oxidation stateHydrocarbon + O2 ↔ CO2
+ H2O
24
Synthesis (Combination)DecompositionSingle ReplacementPrecipitation Reactions
(Double Replacement Reactions)
Neutralization Reactions (Acid/Base)
Redox ReactionsCombustion
NEUTRALIZATION REACTIONS
Acid/Base ReactionsAcid + Base ↔ Salt + Water
25
Properties of Acids
26
Have a sour taste. Vinegar owes its taste to acetic acid. Citrusfruits contain citric acid.
React with certain metals to produce hydrogen gas.
React with carbonates and bicarbonatesto produce carbon dioxide gas
Cause color changes in plant dyes.
2HCl (aq) + Mg (s) MgCl2 (aq) + H2 (g)
2HCl (aq) + CaCO3 (s) CaCl2 (aq) + CO2 (g) + H2O (l)
Aqueous acid solutions conduct electricity.
27
Have a bitter taste.
Feel slippery. Many soaps contain bases.
Properties of Bases
Cause color changes in plant dyes.
Aqueous base solutions conduct electricity.
Examples:
28
Arrhenius acid is a substance that produces H+ (H3O+) in water
Arrhenius base is a substance that produces OH- in water
29
Hydronium ion, hydrated proton, H3O+
30
A Brønsted acid is a proton donorA Brønsted base is a proton acceptor
acidbase acid base
A Brønsted acid must contain at least one ionizable proton!
31
Monoprotic acids
HCl H+ + Cl-
HNO3 H+ + NO3-
CH3COOH H+ + CH3COO-
Strong electrolyte, strong acid
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Diprotic acidsH2SO4 H+ + HSO4
-
HSO4- H+ + SO4
2-
Strong electrolyte, strong acid
Weak electrolyte, weak acid
Triprotic acidsH3PO4 H+ + H2PO4
-
H2PO4- H+ + HPO4
2-
HPO42- H+ + PO4
3-
Weak electrolyte, weak acid
Weak electrolyte, weak acid
Weak electrolyte, weak acid
32
33
Identify each of the following species as a Brønsted acid, base, or both. (a) HI, (b) CH3COO-, (c) H2PO4
-
HI (aq) H+ (aq) + I- (aq) Brønsted acid
CH3COO- (aq) + H+ (aq) CH3COOH (aq) Brønsted base
H2PO4- (aq) H+ (aq) + HPO4
2- (aq)
H2PO4- (aq) + H+ (aq) H3PO4 (aq)
Brønsted acid
Brønsted base
Problem 4.32
34
Identify each of the following as either a(a)Brønsted acid, (b) Brønsted base, or (c) both.1.PO4
3-
2.ClO2-
3.NH4+
4.HCO3-
Neutralization Reaction
35
acid + base salt + water
HCl (aq) + NaOH (aq) NaCl (aq) + H2O
H+ + Cl- + Na+ + OH- Na+ + Cl- + H2O
H+ + OH- H2O
Neutralization Reaction Involving a Weak Electrolyte
36
weak acid + base salt + water
HCN (aq) + NaOH (aq) NaCN (aq) + H2O
HCN + Na+ + OH- Na+ + CN- + H2O
HCN + OH- CN- + H2O
Neutralization Reaction Producing a Gas
37
acid + base salt + water + CO2
2HCl (aq) + Na2CO3 (aq) 2NaCl (aq) + H2O +CO2
2H+ + 2Cl- + 2Na+ + CO32- 2Na+ + 2Cl- + H2O + CO2
2H+ + CO32- H2O + CO2
Types of Chemical ReactionsA + B ↔ CC ↔ A + B
A + BC ↔ AC + BAB + CD ↔ AD + BC
Acid + Base ↔ salt + water
Change of oxidation stateHydrocarbon + O2 ↔ CO2
+ H2O
38
Synthesis (Combination)DecompositionSingle ReplacementPrecipitation Reactions
(Double Replacement Reactions)
Neutralization Reactions (Acid/Base)
Redox ReactionsCombustion
OXIDATION REDUCTION REACTIONS
Redox ReactionsSynthesis Reactions: A + B ↔ CDecomposition Reactions: C ↔ A + BSingle Replacement Reactions: A + BC ↔ AC + BCombustion Reactions: hydrocarbon + O2 ↔ CO2 + H2O
39
Oxidation-Reduction Reactions
40
(electron transfer reactions)
2Mg 2Mg2+ + 4e-
O2 + 4e- 2O2-
Oxidation half-reaction (lose e-)
Reduction half-reaction (gain e-)
2Mg + O2 + 4e- 2Mg2+ + 2O2- + 4e-
2Mg + O2 2MgO
41
42
Zn (s) + CuSO4 (aq) ZnSO4 (aq) + Cu (s)
Zn is oxidizedZn Zn2+ + 2e-
Cu2+ is reducedCu2+ + 2e- Cu
Zn is the reducing agent
Cu2+ is the oxidizing agent
Copper wire reacts with silver nitrate to form silver metal.What is the oxidizing agent in the reaction?
Oxidation number
43
The charge the atom would have in a molecule (or anionic compound) if electrons were completely transferred.
1. Free elements (uncombined state) have an oxidation number of zero.
Na, Be, K, Pb, H2, O2, P4 = 0
2. In monatomic ions, the oxidation number is equal to the charge on the ion.
Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2
3. The oxidation number of oxygen is usually –2. In H2O2
and O22- it is –1.
4.4
44
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.
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.
5. Group IA metals are +1, IIA metals are +2 and fluorine is always –1.
HCO3-
O = –2 H = +1
3x(–2) + 1 + ? = –1
C = +4
What are the oxidation numbers of all the elements in HCO3
- ?
7. Oxidation numbers do not have to be integers. Oxidation number of oxygen in the superoxide ion, O2
-, is –½.
45
The Oxidation Numbers of Elements in their Compounds
46
NaIO3
Na = +1 O = -2
3x(-2) + 1 + ? = 0
I = +5
IF7
F = -1
7x(-1) + ? = 0
I = +7
K2Cr2O7
O = -2 K = +1
7x(-2) + 2x(+1) + 2x(?) = 0
Cr = +6
What are the oxidation numbers of all the elements in each of these compounds? NaIO3 IF7 K2Cr2O7
Problem 4.50
47
Give the oxidation number for the underlined atoms or in each of the following species:a)Mg3N2 b) CsO2 c) CaC2
d)CO32- e) C2O4
2- f) ZnO22-
g) NaBH4 h) WO42-
Types of Oxidation-Reduction Reactions
48
Synthesis/Combination Reaction
A + B C
2Al + 3Br2 2AlBr3
Decomposition Reaction
2KClO3 2KCl + 3O2
C A + B
0 0 +3 -1
+1 +5 -2 +1 -1 0
Types of Oxidation-Reduction Reactions
49
Combustion Reaction
A + O2 B
S + O2 SO2
0 0 +4 -2
2Mg + O2 2MgO0 0 +2 -2
Types of Oxidation-Reduction Reactions
50
Displacement Reaction
A + BC AC + B
Sr + 2H2O Sr(OH)2 + H2
TiCl4 + 2Mg Ti + 2MgCl2
Cl2 + 2KBr 2KCl + Br2
Hydrogen Displacement
Metal Displacement
Halogen Displacement
0 +1 +2 0
0+4 0 +2
0 -1 -1 0
51
The Activity Series for Metals
M + BC MC + B
Hydrogen Displacement Reaction
M is metalBC is acid or H2O
B is H2
Ca + 2H2O Ca(OH)2 + H2
Pb + 2H2O Pb(OH)2 + H2
Problem 4.52
52
Which of the following metals can react with water to produce H2 (g)?
a)Aub)Lic)Hgd)Cae)Pt
53
The Activity Series for Halogens
Halogen Displacement Reaction
Cl2 + 2KBr 2KCl + Br2
0 -1 -1 0
F2 > Cl2 > Br2 > I2
I2 + 2KBr 2KI + Br2
Types of Oxidation-Reduction Reactions
54
The same element is simultaneously oxidized and reduced.
Example:
Disproportionation Reaction
Cl2 + 2OH- ClO- + Cl- + H2O0 +1 -1
oxidized
reduced
55
Ca2+ + CO32- CaCO3
NH3 + H+ NH4+
Zn + 2HCl ZnCl2 + H2
Ca + F2 CaF2
Classify each of the following reactions.
Problem 4.54
56
Predict the outcome of the reactions represented by the following equations by using the activity series, and balance the equations.
Cu (s) + HCl (aq) I2 (g) + NaBr (aq)
Mg (s) + CuSO4 (aq)
Cl2 (g) + KBr (aq)
Problem 4.56
57
Classify the following redox reactions by type:
P4 + 10Cl2 4PCl52NO N2 + O2
Cl2 + 2KI I2 + 2KCl
Solution Stoichiometry
MolarityDilutionsGravimetric AnalysisTitrations
58
59
Solution Stoichiometry
The concentration of a solution is the amount of solute present in a given quantity of solvent or solution.
M = molarity =moles of solute
liters of solution
What mass of KI is required to make 5.00 x 102 mL of a 2.80 M KI solution?
volume of KI solution moles KI grams KIM KI M KI
5.00x102 mL = 232 g KI166 g KI
1 mol KIx
2.80 mol KI
1 L solnx
1 L
1000 mLx
Problem 4.60
60
Calculate the mass in grams of sodium nitrate required to prepare 2.50 x 102 mL of a 0.707 M solution.
61
Preparing a Solution of Known Concentration
62
Dilution is the procedure for preparing a less concentrated solution from a more concentrated solution.
DilutionAdd Solvent
Moles of solutebefore dilution (i)
Moles of soluteafter dilution (f)=
MiVi MfVf=
63
How would you prepare 60.0 mL of 0.200 M HNO3 from a stock solution of 4.00 M HNO3?
MiVi = MfVf
Mi = 4.00 M Mf = 0.200 M Vf = 0.0600 L Vi = ? L
Vi =MfVf
Mi
= 0.200 M x 0.0600 L4.00 M
= 0.00300 L = 3.00 mL
Dilute 3.00 mL of acid with water to a total volume of 60.0 mL.
Problem 4.70
64
Water is added to 25.0 mL of a 0.866 M KNO3 solution until the volume of the solution is exactly 500 mL. What is the concentration of the final solution?
Problem 4.72
65
You have 505 mL of a 0.125 M HCl solution and you want to dilute it to exactly 0.100 M. How much water should you add? (assume that the volumes are additive.)
66
Gravimetric Analysis1. Dissolve unknown substance in water
2. React unknown with known substance to form a precipitate
3. Filter and dry precipitate
4. Weigh precipitate
5. Use chemical formula and mass of precipitate to determine amount of unknown ion
67
A 0.5662 g sample of an ionic compound containing chloride ions and an unknown metal is dissolved in water, and treated with excess AgNO3. If 1.0882 g of AgCl precipitate forms, what is the percent by mass of Cl in the original compound?
%72.24%100AgCl g 143.4
Cl g 35.45 Cl%
g 0.2690 g 1.08820.2472 Cl of mass
%51.47%1005662.0
g 0.2690 %Cl
so sample, original in the Cl ofamount theis This
g
Problem 4.78
68
A sample of 0.6760 g of an unknown compound containing barium ions (Ba2+) is dissolved in water and treated with an excess of Na2SO4. If the mass of the BaSO4 precipitate formed is 0.4105 g, what is the percent by mass of Ba in the original unknown compound?
69
TitrationsIn a titration a solution of accurately known concentration is added gradually added to another solution of unknown concentration until the chemical reaction between the two solutions is complete.
Equivalence point – the point at which the reaction is complete
Indicator – substance that changes color at (or near) the equivalence point
Slowly add baseto unknown acid
UNTIL
the indicatorchanges color
70
Titrations can be used in the analysis of
Acid-base reactions
Redox reactions
H2SO4 + 2NaOH 2H2O + Na2SO4
5Fe2+ + MnO4- + 8H+ Mn2+ + 5Fe3+ + 4H2O
71
What volume of a 1.420 M NaOH solution is required to titrate 25.00 mL of a 4.50 M H2SO4 solution?
WRITE THE CHEMICAL EQUATION!
volume acid moles red moles base volume base
H2SO4 + 2NaOH 2H2O + Na2SO4
4.50 mol H2SO4
1000 mL solnx
2 mol NaOH
1 mol H2SO4
x1000 ml soln
1.420 mol NaOHx25.00 mL = 158 mL
M
acid
rxn
coef.
M
base
Problem 4.86
72
Calculate the concentration (in molarity) of a NaOH solution if 25.0 mL of the solution are needed to neutralize 17.4 mL of a 0.312 M HCl solution.
73
WRITE THE CHEMICAL EQUATION!
volume red moles red moles oxid M oxid
0.1327 mol KMnO4
1 Lx
5 mol Fe2+
1 mol KMnO4
x1
0.02500 L Fe2+x0.01642 L = 0.4358 M
M
red
rxn
coef.
V
oxid
5Fe2+ + MnO4- + 8H+ Mn2+ + 5Fe3+ + 4H2O
16.42 mL of 0.1327 M KMnO4 solution is needed to oxidize 25.00 mL of an acidic FeSO4 solution. What is the molarity of the iron solution?
16.42 mL = 0.01642 L 25.00 mL = 0.02500 L
Problem 4.92
74
The SO2 present in air is mainly responsible for the acid rain phenomenon. Its concentration can be determined by titrating against a standard permanganate solution as follows:
5SO2 + 2MnO4- + 2H2O 5SO4
2- + 2Mn2+ + 4H+
Calculate the number of grams of SO2 in a sample of air if 7.37 mL of 0.00800 M KMnO4 solution are required for the titration.