Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint

CHAPTER 20 “ Oxidation-Reduction Reactions” LEO SAYS GER

Section 20.1 The Meaning of Oxidation and 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 20.1 The Meaning of Oxidation and Reduction (Redox)
OBJECTIVES
State the characteristics of a redox reaction and identify the oxidizing agent and reducing agent.

Section 20.1 The Meaning of Oxidation and Reduction (Redox)
OBJECTIVES
Describe what happens to iron when it corrodes.

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

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

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 .

Today, many of these reactions may not even involve oxygen
Redox currently says that electrons are transferred between reactants
Mg + S -> Mg 2+ + S 2-
The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg 2+
The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S 2-
(MgS)

Oxidation and Reduction (Redox) Each sodium atom loses one electron: Each chlorine atom gains one electron:

LEO says GER : L ose E lectrons = O xidation Sodium is oxidized G ain E lectrons = R eduction Chlorine is reduced

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 Mg 2+ ion S is reduced : gains e - = S 2- ion Mg is the reducing agent S is the oxidizing agent

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 Reactions - Reactions in which there has been no change in oxidation number are NOT redox reactions. Examples:

Corrosion
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

Corrosion
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

Corrosion
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 20.2 Oxidation Numbers
OBJECTIVES
Determine the oxidation number of an atom of any element in a pure substance.

Section 20.2 Oxidation 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 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 Numbers
The oxidation number of any uncombined element is zero .
The oxidation number of a monatomic ion equals its charge .

The oxidation number of oxygen in compounds is -2 , except in peroxides, such as H 2 O 2 where it is -1.
The oxidation number of hydrogen in compounds is +1 , except in metal hydrides, like NaH, where it is -1.

The sum of the oxidation numbers of the atoms in the compound must equal 0.
2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H

The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge.
X + 3(-2) = -1 N O thus X = +5 thus X = +6 X + 4(-2) = -2 S O

Reducing Agents and Oxidizing Agents
An increase in oxidation number = oxidation
A decrease in oxidation number = reduction
Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent

Trends in Oxidation and Reduction
Active metals :
Lose electrons easily
Are easily oxidized
Are strong reducing agents
Active nonmetals :
Gain electrons easily
Are easily reduced
Are strong oxidizing agents

Section 20.3 Balancing Redox Equations
OBJECTIVES
Describe how oxidation numbers are used to identify redox reactions.

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

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

Identifying Redox Equations
In general, all chemical reactions can be assigned to one of two classes:
oxidation-reduction, in which electrons are transferred:
Single-replacement, combination, decomposition, and combustion
this second class has no electron transfer, and includes all others:
Double-replacement and acid-base reactions

Identifying Redox Equations
In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide:
N 2(g) + O 2(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.
YES!

Balancing 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:
Use oxidation-number changes
Use half-reactions

Using Oxidation-Number Changes
Sort of like chemical bookkeeping, you compare the increases and decreases in oxidation numbers.
start with the skeleton equation
Step 1 : assign oxidation numbers to all atoms; write above their symbols
Step 2 : identify which are oxidized/reduced
Step 3 : use bracket lines to connect them
Step 4 : use coefficients to equalize
Step 5 : make sure they are balanced for both atoms and charge – Problem 20.5, 649

Using 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.)

Choosing a Balancing Method
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.
The half-reaction method works best for reactions taking place in acidic or alkaline solution .

Using 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

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

Anode
Anode – 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 (E o ) 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

Electrodes
Cathode – negative electrode (reduction takes place here)
In electrolytic cells POSITIVE ions are REDUCED at the cathode

Electrodes
Anode – 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**

End of Chapter 20

Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint

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