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Chapter 4

The Major Classes of Chemical Reactions

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The Major Classes of Chemical Reactions

4.6 Elements in Redox Reactions

4.1 The Role of Water as a Solvent

4.2 Writing Equations for Aqueous Ionic Reactions

4.3 Precipitation Reactions

4.4 Acid-Base Reactions

4.5 Oxidation-Reduction (Redox) Reactions

4.7 Reversible Reactions: An Introduction to Chemical Equilibrium

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Figure 4.1 Electron distribution in molecules of H2 and H2O.

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Figure 4.2 The dissolution of an ionic compound.

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Figure 4.3 The electrical conductivity of ionic solutions.

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Sample Problem 4.1 Determining Moles of Ions in Aqueous Ionic Solutions

PROBLEM: What is the molarity of each ion are in the following solutions?

(a) 0.1 M of sodium chloride dissolved in water

(b) 0.015 M of ferric sulfite dissolved in water

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Figure 4.4 The hydrated proton.

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Sample Problem 4.2 Determining the Molarity of H+/H3O+ Ions in Aqueous Solutions of Acids

PROBLEM: Nitric acid is a major chemical in the fertilizer and explosivesindustries. In aqueous solution, each molecule dissociates and the H becomes a solvated H+ ion. What is the molarity of H+(aq) in 1.4M nitric acid?

PLAN:

SOLUTION:

Use the formula to find the molarity of H+.

One mole of H+(aq) is released per mole of nitric acid (HNO3)

HNO3(l) H+(aq) + NO3-(aq)

1.4M HNO3(aq) should have 1.4M H+(aq).

H2O

HNO3(l) + H2O(l)) H3O+(aq) + NO3-(aq)

1.4M HNO3(aq) should have 1.4M H3O+ (aq).

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Writing Equations for Aqueous Ionic Reactions

The molecular equation

shows all of the reactants and products as intact, undissociated compounds.

The total ionic equation

shows all of the soluble ionic substances dissociated into ions.

The net ionic equation

eliminates the spectator ions and shows the actual chemical change taking place.

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Figure 4.5 A precipitation reaction and its equation.

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Figure 4.7 The reaction of Pb(NO3)2 and NaI.

double displacement reaction (metathesis)

NaI(aq) + Pb(NO3)2 (aq)

PbI2(s) + NaNO3(aq)

2NaI(aq) + Pb(NO3)2 (aq)

PbI2(s) + 2NaNO3(aq)

2Na+(aq) + 2I-(aq) + Pb2+(aq) + 2NO3-(aq)

PbI2(s) + 2Na+(aq) + 2NO3-(aq)

2NaI(aq) + Pb(NO3)2(aq)

PbI2(s) + 2NaNO3(aq)

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Predicting Whether a Precipitate Will Form

1. Note the ions present in the reactants.

2. Consider the possible cation-anion combinations.

3. Decide whether any of the ion combinations is insoluble.

See Table 4.1 (next slide) for solubility rules.

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Table 4.1 Solubility Rules For Ionic Compounds in Water

1. All common compounds of Group 1A(1) ions (Li+, Na+, K+, etc.) and ammonium ion (NH4

+) are soluble.2. All common nitrates (NO3

-), acetates (CH3COO- or C2H3O2-) and most

perchlorates (ClO4-) are soluble.

3. All common chlorides (Cl-), bromides (Br-) and iodides (I-) are soluble, except those of Ag+, Pb2+, Cu+, and Hg2

2+.

1. All common metal hydroxides are insoluble, except those of Group 1A(1) and the larger members of Group 2A(2)(beginning with Ca2+).

2. All common carbonates (CO32-) and phosphates (PO4

3-) are insoluble, except those of Group 1A(1) and NH4

+.

3. All common sulfides are insoluble except those of Group 1A(1), Group 2A(2) and NH4

+.

Soluble Ionic Compounds

Insoluble Ionic Compounds

4. All common sulfates (SO42-) are soluble, except those of Ca2+, Sr2+,

Ba2+, Ag+, and Pb2+.

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Sample Problem 4.3 Predicting Solubility

PROBLEM: Classify the following ionic compounds as soluble or insoluble in water:

(a) Sodium carbonate

(b) Lead (II) sulfate

(c) Cobalt (II) hydroxide

(d) Barium nitrate

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Sample Problem 4.4 Predicting Whether a Precipitation Reaction Occurs; Writing Ionic Equations

PROBLEM: Predict whether a reaction occurs when each of the following pairs of solutions are mixed. If a reaction does occur, write balanced molecular, total ionic, and net ionic equations, and identify the spectator ions.

(a) sodium chlorate(aq) + ammonium nitrate(aq)(b) silver (I) nitrate(aq) + calcium chloride(aq)

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Table 4.2 Selected Acids and Bases

AcidsStrong

hydrochloric acid, HCl

hydrobromic acid, HBr

hydroiodic acid, HI

nitric acid, HNO3

sulfuric acid, H2SO4

perchloric acid, HClO4

Weakhydrofluoric acid, HF

phosphoric acid, H3PO4

acetic acid, CH3COOH (or HC2H3O2)

BasesStrong

Weak

sodium hydroxide, NaOH

calcium hydroxide, Ca(OH)2

potassium hydroxide, KOH

strontium hydroxide, Sr(OH)2

barium hydroxide, Ba(OH)2

ammonia, NH3

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Classification of Acids Strong Acids—Complete Dissociation, e.g.

HCl(aq) → H+(aq) + Cl–(aq)

H2SO4(aq) → H+(aq) + HSO4–(aq)

Weak Acids—Partial Dissociation, e.g.

HF(aq) H+(aq) + F–(aq)

HSO4–(aq) H+(aq) + SO4

2–(aq)

Strong Bases—Complete Dissociation, e.g.

NaOH(aq) → Na+(aq) + OH–(aq)

Weak Bases—most weak bases actually react with H2O(l) to produce hydroxide ions, e.g.

NH3(aq) + H2O(l) NH4+(aq) + OH–(aq)

Na2CO3(aq) + H2O(l) NaHCO3(aq) + NaOH(aq)

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Acid-Base Reactions(Neutralization Reactions)

HCl(aq) + NaOH(aq) → H2O(l) + NaCl(aq)

HCl, NaOH, and NaCl are strong electrolytes, so total ionic equation is:

H+(aq) + Cl–(aq) + Na+(aq) + OH–(aq) → H2O(l) + Na+(aq) + Cl–(aq)

Net ionic reaction is:

H+(aq) + OH–(aq) → H2O(l)

or

H3O+(aq) + OH–(aq) → 2H2O(l)

This is the net ionic reaction for ALL neutralization reactions involving strong acids and strong bases!

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Definitions

Acid-base reactions are proton transfer reactions.

Bronsted-Lowry acid is a proton donor (HCl in reaction above)

Bronsted-Lowry base is proton acceptor (NaOH in reaction above)

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Figure 4.9

An aqueous strong acid-strong base reaction on the atomic scale.

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Other Acid-Base Reactions (Neutralization Reactions)Weak acid / strong base:

CH3COOH(aq) + NaOH(aq) → NaOOCCH3 (aq) + H2O(l)

CH3COOH(aq) + Na+(aq) + OH–(aq) → Na+(aq) + CH3COO–(aq) + H2O(l)

CH3COOH(aq) + OH–(aq) → CH3COO–(aq) + H2O(l)

Strong acid / weak base:

HCl(aq) + NH3(aq) → NH4Cl(aq)

H+(aq) + Cl–(aq) + NH3(aq) → NH4+(aq) + Cl–(aq)

H+(aq) + NH3(aq) → NH4+(aq)

Weak acid / weak base:

CH3COOH(aq) + NH3(aq) → NH4OOCCH3(aq)

CH3COOH(aq) + NH3(aq) → NH4+(aq) + CH3COO–(aq)

No spectator ions! – total ionic equation ≡ net ionic equation!

Net ionic

Total ionic

Net ionic

Total ionic

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Sample Problem 4.5 Writing Ionic Equations for Acid-Base Reactions

PROBLEM: Write balanced molecular, total ionic, and net ionic equations for each of the following acid-base reactions and identify the spectator ions.

(a) sodium hydroxide(aq) + sulfuric acid(aq)

(b) magnesium hydroxide(s) + perchloric acid(aq)

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Figure 4.8 An acid-base titration.

Start of titrationExcess of acid

Point of neutralization

Slight excess of base

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Sample Problem 4.6 Finding the Concentration of Acid from an Acid-Base Titration

PROBLEM: Concentrations of solutions of analytical reagents (the ones measured from the burette) must themselves be determined from standardized solutions. This can be done in 2 ways:

•Accurately weigh a sample of pure solid acid or base (primary standard) and titrate this sample with a solution of base or acid to be standardized.

•Titrate the solution with one already standardized.

Hydrochloric acid is purchased from a chemical supply house as a0.100 M solution. In a titration 29.67 mL of this solution is required to neutralize a 25.00 mL solution of Ba(OH)2 to be standardized. What is the molarity of the Ba(OH)2 solution?

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Sample Problem 4.7 Finding the Purity of an Acid from an Acid-Base Titration

PROBLEM: An ‘impure’ sample of 1.034 g oxalic acid (H2C2O4 – a diprotic acid) is dissolved in water and the solution titrated against 34.47 mL of 0.485 M sodium hydroxide solution. What is the purity of the oxalic acid sample?

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Figure 4.10 An acid-base reaction that forms a gaseous product.

Molecular equation

NaHCO3(aq) + CH3COOH(aq)

CH3COONa(aq) + CO2(g) + H2O(l)

Total ionic equation

Na+(aq)+ HCO3-(aq) + CH3COOH(aq)

CH3COO-(aq) + Na+(aq) + CO2(g) + H2O(l)

Net ionic equation

HCO3-(aq) + CH3COOH(aq)

CH3COO-(aq) + CO2(g) + H2O(l)

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Figure 4.11 The redox process in compound formation.

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Table 4.3 Rules for Assigning an Oxidation Number (O.N.)

1. For an atom in its elemental form (Na, O2, Cl2, etc.): O.N. = 02. For a monoatomic ion: O.N. = ion charge3. The sum of O.N. values for the atoms in a compound equals zero. The sum of O.N. values for the atoms in a polyatomic ion equals the ion’s charge.

General rules

Rules for specific atoms or periodic table groups

1. For Group 1A(1): O.N. = +1 in all compounds

2. For Group 2A(2): O.N. = +2 in all compounds

3. For hydrogen: O.N. = +1 in combination with nonmetals

4. For fluorine: O.N. = -1 in combination with metals and boron

6. For Group 7A(17): O.N. = -1 in combination with metals, nonmetals (except O), and other halogens lower in the group

5. For oxygen: O.N. = -1 in peroxidesO.N. = -2 in all other compounds(except with F)

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Sample Problem 4.8 Determining the Oxidation Number of an Element

PROBLEM: Determine the oxidation number (O.N.) of each element in these compounds:

(a) P4 (b) Al2O3 (c) KMnO4

(d) NaH (e) Na2Cr2O7 (f) Fe3O4

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Figure 4.12

Highest and lowest oxidation numbers of reactive main-group elements.

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Figure 4.13 A summary of terminology for oxidation-reduction (redox) reactions.

X Y

e-

transfer or shift of electrons

X loses electron(s) Y gains electron(s)

X is oxidized Y is reduced

X is the reducing agent Y is the oxidizing agent

X increases its oxidation number

Y decreases its oxidation number

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Sample Problem 4.9 Recognizing Oxidizing and Reducing Agents and Balancing Redox Equations by the Oxidation Number Method

PROBLEM: Identify the oxidizing agent and reducing agent in each of the following and then balance the equation using the oxidation numbers:

(a) Ag+(aq) + Cu(s) Ag(s) + Cu2+(aq)

(b) Fe2O3(s) + CO(g) Fe(s) + CO2(g)

For more complicated redox equations, the oxidation number method is a bit of a nightmare – use the half-reaction method: Ch. 21, p904-908.

Also for this method – we will work with net ionic equations rather than full equations – much more student friendly!

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Sample Problem 4.10 Balancing Redox Equations by the Half-Reaction Method

PROBLEM: Use the half-reaction method to balance the following equations:

(a) Ag+(aq) + Cu(s) Ag(s) + Cu2+(aq)

(b) SO32–(aq) + MnO4

–(aq) SO42–(aq) + Mn2+(aq) in acid solution

(c) Cr(OH)3(aq) + ClO–(aq) CrO42–(aq) + Cl2(g) in base solution

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Figure 4.14 A redox titration.

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Figure 4.15 Combining elements to form an ionic compound.

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Figure 4.16 Decomposing a compound to its elements.

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Figure 4.17 An active metal displacing hydrogen from water.

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Figure 4.18 The displacement of H from acid by nickel.

Ni(s) + 2H+(aq) Ni2+(aq) + H2(g)

0 +1 +2 0

O.N. increasing

oxidation occurring

reducing agent

O.N. decreasing

reduction occurring

oxidizing agent

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Figure 4.19 Displacing one metal with another.

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Figure 4.20

The activity series

of the metals.

can displace Hfrom water

LiKBaCaNa

stre

ngth

as

redu

cing

age

nts

can displace Hfrom steam

MgAlMnAnCrFeCd

H2

cannot displace H from any source

CuHg AgAu

can displace Hfrom acid

CoNiSnPb

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Sample Problem 4.11 Identifying the Type of Redox Reaction

PROBLEM: Classify each of the following redox reactions as a combination,decomposition, or displacement reaction:

(a) 2NaCl(s) → 2Na(s) + Cl2(g)

(b) 2AgNO3(aq) + Cu(s) → Cu(NO3)2(aq) + 2Ag(s)

(c) 2NaIO3(s) → 2NaI(s) + 3O2(g)

(d) P4(s) + 6Cl2(g) → 4PCl3(l)

(e) 2SO2(g) + O2(g) → 2SO3(g)

(f) Ca(s) + 2H2O(l) → Ca(OH)2(aq) + H2(g)

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Equilibrium

For the first time we have seen reactions that do not go to completion.

HF(aq) H+(aq) + F–(aq)

In solution, all three species are present at constant concentrations.

Dynamic Equilibrium

Macroscopic scale: no changes in concentrations of all species when equilibrium state is reached.

Microscopic scale: Forward ( ) and reverse ( ) reactions are occurring at the same rates.

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Figure 4.21

The equilibrium state.