Reactions in aqueous solutions Redox reactions

Redox reactions

• In precipitation reactions, cations and anions come together to form an insoluble ionic compound.

• In neutralization reactions, H+ ions and OH- ions come together to form H2O molecules.

• In a third kind of reaction electrons are transferred from one reactant to another. Such reactions are called either oxidation- reduction reactions or redox reactions.

One of the most familiar redox reactions is corrosion of a metal.In some cases corrosion is limited to the surface of the metal, with the green coating that forms on copper roofs and statues. In other instances the corrosion goes deeper, eventually compromising the structural integrity of the metal. Iron rusting is an important example.

Redox reactions

Redox reactions

Corrosion is the conversion of a metal into a metal compound by a reaction between the metal and some substance in its environment. When a metal corrodes, each metal atom loses electrons and so forms a cation, which can combine with an anion to form an ionic compound.

The green coating on some buildings contains Cu2+ combined with carbonate and hydroxide anions, rust contains Fe3+ combined with oxide and hydroxide anions, and oxidized silver contains Ag+ combined with sulfide anions.

SECTION 4.4 Oxidation-Reduction Reactions 131

4.4 | OXIDATION-REDUCTION REACTIONSIn precipitation reactions, cations and anions come together to form an insoluble ioniccompound. In neutralization reactions, ions and ions come together to formH2O molecules. Now let’s consider a third kind of reaction, one in which electrons aretransferred from one reactant to another. Such reactions are called either oxidation-reduction reactions or redox reactions. In this chapter we concentrate on redox reac-tions where one of the reactants is a metal in its elemental form.

Oxidation and ReductionOne of the most familiar redox reactions is corrosion of a metal (! FIGURE 4.11). Insome instances corrosion is limited to the surface of the metal, with the green coatingthat forms on copper roofs and statues being one such case. In other instances the cor-rosion goes deeper, eventually compromising the structural integrity of the metal. Ironrusting is an important example.

Corrosion is the conversion of a metal into a metal compound by a reaction betweenthe metal and some substance in its environment. When a metal corrodes, each metalatom loses electrons and so forms a cation, which can combine with an anion to form anionic compound. The green coating on the Statue of Liberty contains Cu2+ combinedwith carbonate and hydroxide anions, rust contains Fe3+ combined with oxide andhydroxide anions, and silver tarnish contains Ag+ combined with sulfide anions.

When an atom, ion, or molecule becomes more positively charged (that is, when itloses electrons), we say that it has been oxidized. Loss of electrons by a substance is calledoxidation. The term oxidation is used because the first reactions of this sort to be stud-ied were reactions with oxygen. Many metals react directly with O2 in air to form metaloxides. In these reactions the metal loses electrons to oxygen, forming an ionic com-pound of the metal ion and oxide ion. The familiar example of rusting involves the reac-tion between iron metal and oxygen in the presence of water. In this process Fe isoxidized (loses electrons) to form .

The reaction between iron and oxygen tends to be relatively slow, but other metals,such as the alkali and alkaline earth metals, react quickly upon exposure to air." FIGURE 4.12 shows how the bright metallic surface of calcium tarnishes as CaOforms in the reaction

[4.23]

In this reaction Ca is oxidized to Ca2+ and neutral O2 is transformed to ions. Whenan atom, ion, or molecule becomes more negatively charged (gains electrons), we saythat it is reduced. The gain of electrons by a substance is called reduction. When one reac-tant loses electrons (that is, when it is oxidized), another reactant must gain them. Inother words, oxidation of one substance must be accompanied by reduction of someother substance.

O2-2 Ca1s2 + O21g2 ¡ 2 CaO1s2

Fe3+

OH-H+

(a) (b) (c)

# FIGURE 4.11 Familiar corrosion products. (a) A green coating forms when copper is oxidized. (b) Rust forms when iron corrodes.(c) A black tarnish forms as silver corrodes.

Oxidation When an ion become more positively charged (it loses electrons), we say that it has been oxidized. Loss of electrons by a substance is called oxidation.The term oxidation is used because the first reactions of this sort to be studied were reactions with oxygen.

Oxidation Vs Reduction

Many metals react directly with O2 in air to form metal oxides. In these reactions the metal loses electrons to oxygen, forming an ionic compound of the metal ion and oxide ion.

The familiar example of rusting involves the reaction between iron metal and oxygen in the presence of water. In this process Fe is oxidized (loses electrons) to form Fe3+.

4 Fe(s) + 3O2(g) ⟶ 2Fe2O3(s)

Oxidation Vs Reduction

Many metals, such as the alkali and alkaline earth metals, react quickly upon exposure to air.

E.g. The bright metallic surface of calcium tarnishes as CaO(s) forms in the reaction. 2 Ca(s) + O2(g) ⟶ 2CaO(s)

132 CHAPTER 4 Reactions in Aqueous Solution

Oxidation NumbersBefore we can identify an oxidation-reduction reaction, we must have a bookkeepingsystem—a way of keeping track of electrons gained by the substance being reduced andelectrons lost by the substance being oxidized. The concept of oxidation numbers (alsocalled oxidation states) was devised as a way of doing this. Each atom in a neutral sub-stance or ion is assigned an oxidation number. For monatomic ions the oxidation num-ber is the same as the charge. For neutral molecules and polyatomic ions, the oxidationnumber of a given atom is a hypothetical charge. This charge is assigned by artificiallydividing up the electrons among the atoms in the molecule or ion. We use the followingrules for assigning oxidation numbers:

1. For an atom in its elemental form, the oxidation number is always zero. Thus, eachH atom in the H2 molecule has an oxidation number of 0 and each P atom in theP4 molecule has an oxidation number of 0.

2. For any monatomic ion the oxidation number equals the ionic charge. Thus, hasan oxidation number of , has an oxidation number of , and so forth. Inionic compounds the alkali metal ions (group 1A) always have a charge andtherefore an oxidation number of . The alkaline earth metals (group 2A) are al-ways , and aluminum (group 3A) is always in ionic compounds. (In writingoxidation numbers we will write the sign before the number to distinguish themfrom the actual electronic charges, which we write with the number first.)

3. Nonmetals usually have negative oxidation numbers, although they can sometimesbe positive:

(a) The oxidation number of oxygen is usually in both ionic and molecular com-pounds. The major exception is in compounds called peroxides, which containthe ion, giving each oxygen an oxidation number of .

(b) The oxidation number of hydrogen is usually when bonded to nonmetals andwhen bonded to metals.

(c) The oxidation number of fluorine is in all compounds. The other halogenshave an oxidation number of in most binary compounds. When combinedwith oxygen, as in oxyanions, however, they have positive oxidation states.

4. The sum of the oxidation numbers of all atoms in a neutral compound is zero. Thesum of the oxidation numbers in a polyatomic ion equals the charge of the ion. Forexample, in the hydronium ion the oxidation number of each hydrogen is

and that of oxygen is . Thus, the sum of the oxidation numbers is, which equals the net charge of the ion. This rule is useful in

obtaining the oxidation number of one atom in a compound or ion if you know theoxidation numbers of the other atoms, as illustrated in Sample Exercise 4.8.

3(+1) + 1-2) = +1-2+1

H3O+

-1-1

-1+1

-1O22-

-2

+3+2+1

1+-2S2-+1

K+

O2(g) is reduced(gains electrons)

2 CaO(s)2 Ca(s) ! O2(g)ProductsReactants

Ca(s) is oxidized(loses electrons)

Ca2! and O2" ionscombine to form CaO(s)

e"

e"

e"

! FIGURE 4.12 Oxidation of calcium metal by molecular oxygen. The oxidation involves transfer of electrons from thecalcium metal to the O2, leading to formation of CaO.

In this reaction Ca is oxidized to Ca2+ and neutral O2 is transformed to O2- ions.

Reduction When an atom, ion, or molecule becomes more negatively charged (gains electrons), we say that it is reduced.

The gain of electrons by a substance is called reduction. When one reactant loses electrons (when it is oxidized), another reactant must gain them. In other words, oxidation of one substance must be accompanied by reduction of some other substance.

Oxidation Vs Reduction

Each atom is assigned an oxidation number. For monatomic ions the oxidation number is the same as the charge. We use the following rules for assigning oxidation numbers:

Oxidation States rules (a review)

1. For an atom in its elemental form, the oxidation number is always zero

2. For any monatomic ion the oxidation number equals the ionic charge.

3. Nonmetals usually have negative oxidation numbers a. The oxidation number of oxygen is usually -2 in both ionic and molecular compounds.b. The oxidation number of hydrogen is usually +1 when bonded to nonmetals and -1 when

bonded to metals.c. The oxidation number of fluorine is -1 in all compounds. The other halogens have an

oxidation number of - 1 in most binary compounds. When combined with oxygen, as in oxyanions, however, they have positive oxidation states.

4. The sum of the oxidation numbers of all atoms in a neutral compound is zero.

5. The sum of the oxidation numbers in a polyatomic ion equals the charge of the ion.

The reaction between a metal and either an acid or a metal salt conforms to the general pattern

Oxidation of Metals by Acids and Salts

A + BX ⟶ AX + B

Examples: Zn (s) + 2 HBr (aq) ⟶ ZnBr2 (aq) + H2 (g)

Mn (s) + Pb(NO3)2 (aq) ⟶ Mn(NO3)2 (aq) + Pb (s)

These reactions are called displacement reactions because the ion in solution is displaced (replaced) through oxidation of an element.

Many metals undergo displacement reactions with acids, producing salts and hydrogen gas. For example, magnesium metal reacts with hydrochloric acid to form magnesium chloride and hydrogen.

Oxidation of Metals by Acids and Salts 134 CHAPTER 4 Reactions in Aqueous Solution

Many metals undergo displacement reactions with acids, producing salts andhydrogen gas. For example, magnesium metal reacts with hydrochloric acid to formmagnesium chloride and hydrogen gas (! FIGURE 4.13):

The oxidation number of Mg changes from 0 to , an increase that indicates the atomhas lost electrons and has therefore been oxidized. The oxidation number of in theacid decreases from to 0, indicating that this ion has gained electrons and has there-fore been reduced. Chlorine has an oxidation number of both before and after thereaction, indicating that it is neither oxidized nor reduced. In fact the Cl– ions are spec-tator ions, dropping out of the net ionic equation:

[4.26]

Metals can also be oxidized by aqueous solutions of various salts. Iron metal, forexample, is oxidized to by aqueous solutions of such as Ni(NO3)2(aq):

Molecular equation: [4.27]

Net ionic equation: [4.28]

The oxidation of Fe to in this reaction is accompanied by the reduction ofto Ni. Remember: Whenever one substance is oxidized, another substance must be reduced.

Ni2+Fe2+

Fe1s2 + Ni2+1aq2 ¡ Fe2+1aq2 + Ni1s2Fe1s2 + Ni1NO3221aq2 ¡ Fe1NO3221aq2 + Ni1s2Ni2+Fe2+

Mg1s2 + 2 H+1aq2 ¡ Mg2+1aq2 + H21g2-1

+1H+

+2

Mg(s) ! 2 HCl(aq) MgCl2(aq) ! H2(g)

!1 "1 00Oxidation number !2 "1

[4.25]

e!

e!

e!

H!(aq) is reduced(gains electrons)

H!

Mg2!

H2

MgCl"

"

"

Mg(s) is oxidized(loses electrons)

H2(g) ! MgCl2(aq)2 HCl(aq) ! Mg(s)Oxidationnumber

!1 !2 "10"1 0

ProductsReactants

! FIGURE 4.13 Reaction of magnesium metal with hydrochloric acid. The metal is readily oxidized by the acid, producing hydrogengas, H2(g), and MgCl2(aq).

Many metals undergo displacement reactions with acids, producing salts and hydrogen gas. For example, magnesium metal reacts with hydrochloric acid to form magnesium chloride and hydrogen

Oxidation of Metals by Acids and Salts

Mg(s) + 2 HCl (aq) ⟶ MgCl2 (aq) + H2 (g)

0 +2+1 0

The oxidation number of Mg changes from 0 to + 2, an increase that indicates the atom has lost electrons and has therefore been oxidized. The oxidation number of H+ in the acid decreases from +1 to 0, indicating that this ion has gained electrons and has therefore been reduced. Chlorine has an oxidation number of -1 both before and after the reaction, indicating that it is neither oxidized nor reduced. In fact the Cl– ions are spectator ions, dropping out of the net ionic equation:

Mg(s) + 2 H+ (aq) ⟶ Mg2+ (aq) + H2 (g)

Oxidation of Metals by Acids and Salts

Metals can also be oxidized by aqueous solutions of various salts. Iron metal, for example, is oxidized to Fe2+ by aqueous solutions of Ni2+ such as Ni(NO3)2 (aq):

Fe(s) + Ni(NO3)2 (aq) ⟶ Fe(NO3)2 (aq) + Ni (s)

Fe(s) + Ni2+ (aq) ⟶ Fe2+ (aq) + Ni (s)

Molecular equation

Ionic equation

The oxidation of Fe to Fe2+ is accompanied by the reduction of Ni2+ to Ni.

Whenever one substance is oxidized, another substance must be reduced.

Predicting redox reactions - the Activity SeriesWe can predict whether a metal will be oxidized either by an acid or a salt as different metals vary in the ease with which they are oxidized. Zn is oxidized by aqueous solutions of Cu2+ but Ag is not. Zn, therefore, loses electrons more readily than Ag; that is, Zn is easier to oxidize than Ag.

Experimentally we can create a list of metals arranged in order of decreasing ease of oxidation is called an activity series.

Activity Series of Metals in Aqueous Solution Metal Oxidation reactioLithium Li (s) ⟶ Li2+ (aq) + e-

Potassium K (s) ⟶ K+ (aq) + 2e-

Barium Ba (s) ⟶ Ba2+ (aq) + 2e-

Calcium Ca (s) ⟶ Ca2+ (aq) + 2e-

Sodium Na (s) ⟶ Na+ (aq) + e-

Magnesium Mg(s) ⟶ Mg2+ (aq) + 2e-

Aluminum Al(s) ⟶ Al3+ (aq) + 3e-

Manganese Mn(s) ⟶ Mn2+ (aq) + 2e-

Zinc Zn(s) ⟶ Zn2+ (aq) + 2e-

Chromium Cr(s) ⟶ Cr3+ (aq) + 3e-

Iron Fe(s) ⟶ Fe2+ (aq) + 2e-

Cobalt Co(s) ⟶ Co2+ (aq) + 2e-

Nickel Ni(s) ⟶ Ni2+ (aq) + 2e-

Tin Sn(s) ⟶ Sn2+ (aq) + 2e-

Lead Pb(s) ⟶ Pb2+ (aq) + 2e-

Hydrogen H2(g) ⟶ 2H+ (aq) + 2e-

Copper Cu(s) ⟶ Cu2+ (aq) + 2e-

Silver Ag(s) ⟶ Ag+ (aq) + e-

Mercury Hg(l) ⟶ Hg2+ (aq) + 2e-

Platinum Pt(s) ⟶ Pt2+ (aq) + 2e-

Gold Au(s) ⟶ Au3+ (aq) + 3e-

Ease

of o

xida

tion

incr

ease

s

Any metal on the list can be oxidized by the elements below it.E.g., Cu is above Ag in the series. Thus, copper metal is oxidized by silver ions:

SECTION 4.4 Oxidation-Reduction Reactions 137

e!

e!

Cu(s) is oxidized(loses electrons)

Ag!(aq) is reduced(gains electrons)

Cu(NO3)2(aq) ! 2 Ag(s)2 AgNO3(aq) ! Cu(s)ProductsReactants

NO3"Ag! Ag

NO3" Cu2!

Cu

++

! FIGURE 4.14 Reaction of coppermetal with silver ion. When copper metal isplaced in a solution of silver nitrate, a redoxreaction forms silver metal and a bluesolution of copper(II) nitrate.

SAMPLE EXERCISE 4.10 Determining When an Oxidation-ReductionReaction Can Occur

Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write thebalanced molecular and net ionic equations for the reaction.

SOLUTIONAnalyze We are given two substances—an aqueous salt, FeCl2, and a metal, Mg—and askedif they react with each other.

Plan A reaction occurs if the reactant that is a metal in its elemental form (Mg) is locatedabove the reactant that is a metal in its oxidized form (Fe2+) in Table 4.5. If the reaction occurs,the ion in FeCl2 is reduced to Fe, and the Mg is oxidized to .

Solve Because Mg is above Fe in the table, the reaction occurs. To write the formula for thesalt produced in the reaction, we must remember the charges on common ions. Magnesium isalways present in compounds as ; the chloride ion is . The magnesium salt formed inthe reaction is MgCl2, meaning the balanced molecular equation is

Both FeCl2 and MgCl2 are soluble strong electrolytes and can be written in ionic form, whichshows us that is a spectator ion in the reaction. The net ionic equation is

The net ionic equation shows that Mg is oxidized and is reduced in this reaction.

Check Note that the net ionic equation is balanced with respect to both charge and mass.

PRACTICE EXERCISEWhich of the following metals will be oxidized by Pb(NO3)2: Zn, Cu, Fe?

Answer: Zn and Fe

Fe2+

Mg1s2 + Fe2+1aq2 ¡ Mg2+1aq2 + Fe1s2Cl-

Mg1s2 + FeCl21aq2 ¡ MgCl21aq2 + Fe1s2Cl-Mg2+

Mg2+Fe2+

Cu(s) + 2Ag+ (aq) ⟶ Cu2+ (aq) + 2Ag (s) The oxidation of copper to copper ions is accompanied by the reduction of silver ions to silver metal. The silver metal is evident on the surface of the copper wire.

Predicting redox reactions - the Activity Series

Only metals above hydrogen in the activity series are able to react with acids to form H2. For example, Ni reacts with HCl(aq) to form H2:

Ni(s) + 2HCl (aq) ⟶ NiCl (aq) + H2 (s)

Because elements below hydrogen in the activity series are not oxidized by H+, Cu does not react with HCl(aq).

PRACTICE EXERCISE

• Will an aqueous solution of iron(II) chloride oxidize magnesium metal? If so, write the balanced molecular and net ionic equations for the reaction.

SolutionTwo substances (an aqueous salt, FeCl2, and a metal, Mg)A reaction occurs if the reactant that is a metal in its elemental form (Mg) is located above the reactant that is a metal in its oxidized form (Fe2+) in the “Activity Series of Metals in Aqueous Solution” Table. If the reaction occurs, the Fe2+ ion in FeCl2 is reduced to Fe, and the Mg is oxidized to Mg2+.

Because Mg is above Fe in the table, the reaction occurs. To write the formula for the salt produced in the reaction, we must remember the charges on common ions. Magnesium is always present in compounds as Mg2+; the chloride ion is Cl-. The magnesium salt formed in the reaction is MgCl2, meaning the balanced molecular equation is:

Mg (s) + FeCl2 (aq) ⟶ MgCl2 (aq) + Fe (s) Both FeCl2 and MgCl2 are soluble strong electrolytes and can be written in ionic form, which shows us that Cl- is a spectator ion in the reaction. The net ionic equation is:

Mg (s) + Fe2+ (aq) ⟶ Mg2+ (aq) + Fe (s) The net ionic equation shows that Mg is oxidized and Fe2+ is reduced in this reaction.

• Which of the following metals will be oxidized by Pb(NO3)2: Zn, Cu, Fe? Answers:

Zn and Fe

Determining when an Redox can occur

PRACTICE EXERCISE • Determine the oxidation number for the indicated element in each of the following substances:

a) S in SO2

b) C in COCl2c) Mn in KMnO4

d) Br in HBrOe) As in As4

f) O in K2O2

• Which element is oxidized and which is reduced in the following reactions?a) N2 (g) + 3 H2 (g) ⟶ 2 NH3 (g)b) 3Fe(NO3)2 (aq) + 2 Al (s) ⟶ 3 Fe (s) + 2Al(NO3)3 (aq)c) Cl2 (aq) + 2 NaI (aq) ⟶ I2 (aq) + 2 NaCl (aq)d) PbS (s) +4 H2O2 (aq) ⟶ PbSO4 (s) + 4 H2O (l)

• Write balanced molecular and net ionic equations for the reactions ofa) manganese with sulfuric acid in aqueous solutionb) chromium with hydrobromic acidc) tin with hydrochloric acidd) aluminum with formic acid (HCOOH)

• Using the activity series table, write balanced chemical equations for the following reactions. If no reaction occurs, simply write NR.

a) Iron metal is added to a solution of copper(II) nitrateb) zinc metal is added to a solution of magnesium sulfatec) hydrobromic acid is added to tin metald) hydrogen gas is bubbled through an aqueous solution of nickel(II) chloridee) aluminum metal is added to a solution of cobalt(II) sulfate.

Reduction-oxidation reactions