Chapter 4

Reactions in Aqueous Solution

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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.