chemistry - chp 20 - oxidation reduction reactions - powerpoint

CHAPTER 20CHAPTER 20

““Oxidation-Reduction Reactions”Oxidation-Reduction Reactions”

LEO SAYS GER

Section 20.1Section 20.1The Meaning of Oxidation and The Meaning of Oxidation and

Reduction (called “redox”)Reduction (called “redox”)

OBJECTIVES

Define oxidation and reduction in terms of the loss or gain of oxygen, and the loss or gain of electrons.

Section 20.1Section 20.1The Meaning of Oxidation and The Meaning of Oxidation and

Reduction (Redox)Reduction (Redox)

OBJECTIVES

State the characteristics of a redox reaction and identify the oxidizing agent and reducing agent.

Section 20.1Section 20.1The Meaning of Oxidation and The Meaning of Oxidation and

Reduction (Redox)Reduction (Redox)

OBJECTIVES

Describe what happens to iron when it corrodes.

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox)

Early chemists saw “oxidation” reactions only as the combination of a material with oxygen to produce an oxide.

• For example, when methane burns in air, it oxidizes and forms oxides of carbon and hydrogen, as shown in Fig. 20.1, p. 631

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox)

But, not all oxidation processes that use oxygen involve burning:

•Elemental iron slowly oxidizes to compounds such as iron (III) oxide, commonly called “rust”•Bleaching stains in fabrics•Hydrogen peroxide also releases oxygen when it decomposes

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox)

A process called “reduction” is the opposite of oxidation, and originally meant the loss of oxygen from a compound Oxidation and reduction always occur

simultaneously The substance gaining oxygen (or

losing electrons) is oxidized, while the substance losing oxygen (or gaining electrons) is reduced.

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox) Today, many of these reactions may

not even involve oxygen Redox currently says that electrons

are transferred between reactants

Mg + S → Mg2+ + S2-

•The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg2+

•The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S2-

(MgS)

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox)11

2

00

22

ClNaClNa

Each sodium atom loses one electron:

Each chlorine atom gains one electron:

eNaNa10

10 CleCl

LEO says GER :LEO says GER :

eNaNa10

Lose Electrons = Oxidation

Sodium is oxidized

Gain Electrons = Reduction

10 CleCl Chlorine is reduced

LEO says GER :LEO says GER : - Losing electrons is oxidation, and the substance that loses the electrons is called the reducing agent. - Gaining electrons is reduction, and the substance that gains the electrons is called the oxidizing agent.

Mg(s) + S(s) → MgS(s)

Mg is oxidized: loses e-, becomes a Mg2+ ion

S is reduced: gains e- = S2- ion

Mg is the reducing

agent

S is the oxidizing agent

Oxidation and Reduction (Redox)Oxidation and Reduction (Redox) Conceptual Problem 20.1, page 634

It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds?

In water, we learned that oxygen is highly electronegative, so:

the oxygen gains electrons (is reduced and is the oxidizing agent), and the hydrogen loses electrons (is oxidized and is the reducing agent)

Not All Reactions are Redox ReactionsNot All Reactions are Redox Reactions

- Reactions in which there has been no

change in oxidation number are NOT redox reactions.

Examples:

)()()()( 3

2511111

3

251

aqONNasClAgaqClNaaqONAg

)()()()(22

2

1

4

26

2

1

4

26

2

1121

lOHaqOSNaaqOSHaqHONa

CorrosionCorrosion•Damage done to metal is costly to prevent and repair•Iron, a common construction metal often used in forming steel alloys, corrodes by being oxidized to ions of iron by oxygen.

•This corrosion is even faster in the presence of salts and acids, because these materials make electrically conductive solutions that make electron transfer easy

CorrosionCorrosion•Luckily, not all metals corrode easily

•Gold and platinum are called noble metals because they are resistant to losing their electrons by corrosion•Other metals may lose their electrons easily, but are protected from corrosion by the oxide coating on their surface, such as aluminum – Figure 20.7, page 636•Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily

CorrosionCorrosion•Serious problems can result if bridges, storage tanks, or hulls of ships corrode

•Can be prevented by a coating of oil, paint, plastic, or another metal•If this surface is scratched or worn away, the protection is lost

•Other methods of prevention involve the “sacrifice” of one metal to save the second

•Magnesium, chromium, or even zinc (called galvanized) coatings can be applied

Section 20.2Section 20.2Oxidation NumbersOxidation Numbers

OBJECTIVES

Determine the oxidation number of an atom of any element in a pure substance.

Section 20.2Section 20.2Oxidation NumbersOxidation Numbers

OBJECTIVES

Define oxidation and reduction in terms of a change in oxidation number, and identify atoms being oxidized or reduced in redox reactions.

Assigning Oxidation NumbersAssigning Oxidation Numbers

• An “oxidation number” is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction.

• Generally, a bonded atom’s oxidation number is the charge it would have if the electrons in the bond were assigned to the atom of the more electronegative element

Rules for Assigning Oxidation NumbersRules for Assigning Oxidation Numbers

1)The oxidation number of any uncombined element is zero.

2)The oxidation number of a monatomic ion equals its charge.

11

2

00

22

ClNaClNa

Rules for Assigning Oxidation NumbersRules for Assigning Oxidation Numbers

3)The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1.

4)The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1.

2

2

1

OH

Rules for Assigning Oxidation NumbersRules for Assigning Oxidation Numbers

5) The sum of the oxidation numbers of the atoms in the compound must equal 0.

2

2

1

OH2(+1) + (-2) = 0 H O

2

122

)(

HOCa(+2) + 2(-2) + 2(+1) = 0 Ca O H

Rules for Assigning Oxidation NumbersRules for Assigning Oxidation Numbers

6) The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge.

3

2?

ONX + 3(-2) = -1N O

24

2?

OS

thus X = +5 thus X = +6

X + 4(-2) = -2S O

Reducing Agents and Oxidizing AgentsReducing Agents and Oxidizing Agents

• Conceptual Problem 20.2, page 641

• An increase in oxidation number = oxidation

• A decrease in oxidation number = reduction

eNaNa10

10 CleCl

Sodium is oxidized – it is the reducing agent

Chlorine is reduced – it is the oxidizing agent

Trends in Oxidation and ReductionTrends in Oxidation and ReductionActive metals:

Lose electrons easily Are easily oxidized Are strong reducing agents

Active nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agents

Conceptual Problem 20.3, page 643

Section 20.3Section 20.3Balancing Redox EquationsBalancing Redox Equations

OBJECTIVES

Describe how oxidation numbers are used to identify redox reactions.

Section 20.3Section 20.3Balancing Redox EquationsBalancing Redox Equations

OBJECTIVES

Balance a redox equation using the oxidation-number-change method.

Section 20.3Section 20.3Balancing Redox EquationsBalancing Redox Equations

OBJECTIVES

Balance a redox equation by breaking the equation into oxidation and reduction half-reactions, and then using the half-reaction method.

Identifying Redox EquationsIdentifying Redox Equations In general, all chemical reactions can be assigned to one of two classes:

1)oxidation-reduction, in which electrons are transferred:• Single-replacement, combination,

decomposition, and combustion

2) this second class has no electron transfer, and includes all others:

• Double-replacement and acid-base reactions

Identifying Redox EquationsIdentifying Redox Equations In an electrical storm, nitrogen and

oxygen react to form nitrogen monoxide:

N2(g) + O2(g) → 2NO(g)

•Is this a redox reaction?

•If the oxidation number of an element in a reacting species changes, then that element has undergone either oxidation or reduction; therefore, the reaction as a whole must be a redox.

•Conceptual Problem 20.4, page 647

YES!

Balancing Redox EquationsBalancing Redox Equations It is essential to write a correctly balanced equation that represents what happens in a chemical reaction

• Fortunately, two systematic methods are available, and are based on the fact that the total electrons gained in reduction equals the total lost in oxidation. The two methods:

1)Use oxidation-number changes

2)Use half-reactions

Using half-reactionsUsing half-reactions A half-reaction is an equation showing

just the oxidation or just the reduction that takes place

they are then balanced separately, and finally combined

Step 1: write unbalanced equation in ionic form

Step 2: write separate half-reaction equations for oxidation and reduction

Step 3: balance the atoms in the half-reactions (More steps on the next screen.)

ChoosingChoosing a Balancing Method a Balancing Method

1) The oxidation number change method works well if the oxidized and reduced species appear only once on each side of the equation, and there are no acids or bases.

2) The half-reaction method works best for reactions taking place in acidic or alkaline solution.

Using half-reactionsUsing half-reactions continued

•Step 4: add enough electrons to one side of each half-reaction to balance the charges

•Step 5: multiply each half-reaction by a number to make the electrons equal in both

•Step 6: add the balanced half-reactions to show an overall equation

•Step 7: add the spectator ions and balance the equation

•Rules shown on page 651 – bottom

•Conceptual Problem 20.6, page 652

Electrochemical (Voltaic) Cells

•An apparatus that allows a redox reaction to occur by transferring electrons through an external connector

•Redox reactions that occur spontaneously may be employed to provide a source of electrical energy

•When the two half cells of a redox reaction are connected by an external conductor and a salt bridge that allows the migration of ions, a flow of electrons (electric current) is produced

• In a voltaic cell, a chemical reaction is used to produce a spontaneous electric current by converting chemical energy to electrical energy

Basic Concepts of Electrochemical Cells

Cathode

Cathode - The electrode where reduction occurs

•In an electrochemical (voltaic) cell, the cathode is the POSITIVE electrode

AnodeAnode – the electrode where

oxidation occurs• In an electrochemical

(voltaic) cell, the anode is the NEGATIVE electrode

•The anode is the more active metal (according to table J)

Electrolytic Cells• Sometimes, in combining half reaction,

the potential (Eo) for the overall reaction is negative. In this case, the reaction will not take place spontaneously.

• Redox reactions that do not occur spontaneously can be forced to take place by supplying energy with an externally applied electric current

• The use of an electric current to bring about a chemical reaction is called electrolysis

• In an electrolytic cell, an electric current is used to produce a chemical reaction

Basic Concepts of Electrolytic Cells

ElectrodesCathode – negative electrode

(reduction takes place here)• In electrolytic cells POSITIVE ions

are REDUCED at the cathode

ElectrodesAnode – Positive anode (oxidation

takes place here)• In electrolytic cells NEGATIVE

ions are OXIDIZED at the anode.

**NOTE: The charge of the electrode is the opposite in electrochemical cells**