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CHAPTER 20 “ Oxidation-Reduction Reactions” LEO SAYS GER
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Section 20.1 The Meaning of Oxidation and Reduction (called “redox”) <ul><li>OBJECTIVES </li></ul><ul><ul><li>Define oxidation and reduction in terms of the loss or gain of oxygen, and the loss or gain of electrons. </li></ul></ul>
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Section 20.1 The Meaning of Oxidation and Reduction (Redox) <ul><li>OBJECTIVES </li></ul><ul><ul><li>State the characteristics of a redox reaction and identify the oxidizing agent and reducing agent. </li></ul></ul>
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Section 20.1 The Meaning of Oxidation and Reduction (Redox) <ul><li>OBJECTIVES </li></ul><ul><ul><li>Describe what happens to iron when it corrodes. </li></ul></ul>
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Oxidation and Reduction (Redox) <ul><li>Early chemists saw “ oxidation ” reactions only as the combination of a material with oxygen to produce an oxide. </li></ul><ul><ul><li>For example, when methane burns in air, it oxidizes and forms oxides of carbon and hydrogen, as shown in Fig. 20.1, p. 631 </li></ul></ul>
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Oxidation and Reduction (Redox) <ul><li>But, not all oxidation processes that use oxygen involve burning: </li></ul><ul><ul><li>Elemental iron slowly oxidizes to compounds such as iron (III) oxide, commonly called “rust” </li></ul></ul><ul><ul><li>Bleaching stains in fabrics </li></ul></ul><ul><ul><li>Hydrogen peroxide also releases oxygen when it decomposes </li></ul></ul>
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Oxidation and Reduction (Redox) <ul><li>A process called “ reduction ” is the opposite of oxidation, and originally meant the loss of oxygen from a compound </li></ul><ul><li>Oxidation and reduction always occur simultaneously </li></ul><ul><li>The substance gaining oxygen (or losing electrons) is oxidized , while the substance losing oxygen (or gaining electrons) is reduced . </li></ul>
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Oxidation and Reduction (Redox) <ul><li>Today, many of these reactions may not even involve oxygen </li></ul><ul><li>Redox currently says that electrons are transferred between reactants </li></ul><ul><li>Mg + S -> Mg 2+ + S 2- </li></ul><ul><li>The magnesium atom (which has zero charge) changes to a magnesium ion by losing 2 electrons, and is oxidized to Mg 2+ </li></ul><ul><li>The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S 2- </li></ul>(MgS)
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Oxidation and Reduction (Redox) Each sodium atom loses one electron: Each chlorine atom gains one electron:
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LEO says GER : L ose E lectrons = O xidation Sodium is oxidized G ain E lectrons = R eduction Chlorine is reduced
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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
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Oxidation and Reduction (Redox) <ul><li>Conceptual Problem 20.1, page 634 </li></ul><ul><li>It is easy to see the loss and gain of electrons in ionic compounds, but what about covalent compounds? </li></ul><ul><li>In water, we learned that oxygen is highly electronegative , so: </li></ul><ul><ul><li>the oxygen gains electrons (is reduced and is the oxidizing agent), and the hydrogen loses electrons (is oxidized and is the reducing agent) </li></ul></ul>
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Not All Reactions are Redox Reactions - Reactions in which there has been no change in oxidation number are NOT redox reactions. Examples:
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Corrosion <ul><li>Damage done to metal is costly to prevent and repair </li></ul><ul><li>Iron , a common construction metal often used in forming steel alloys, corrodes by being oxidized to ions of iron by oxygen. </li></ul><ul><ul><li>This corrosion is even faster in the presence of salts and acids , because these materials make electrically conductive solutions that make electron transfer easy </li></ul></ul>
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Corrosion <ul><li>Luckily, not all metals corrode easily </li></ul><ul><ul><li>Gold and platinum are called noble metals because they are resistant to losing their electrons by corrosion </li></ul></ul><ul><ul><li>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 </li></ul></ul><ul><ul><li>Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily </li></ul></ul>
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Corrosion <ul><li>Serious problems can result if bridges, storage tanks, or hulls of ships corrode </li></ul><ul><ul><li>Can be prevented by a coating of oil, paint, plastic, or another metal </li></ul></ul><ul><ul><li>If this surface is scratched or worn away, the protection is lost </li></ul></ul><ul><li>Other methods of prevention involve the “ sacrifice ” of one metal to save the second </li></ul><ul><ul><li>Magnesium, chromium, or even zinc (called galvanized) coatings can be applied </li></ul></ul>
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Section 20.2 Oxidation Numbers <ul><li>OBJECTIVES </li></ul><ul><ul><li>Determine the oxidation number of an atom of any element in a pure substance. </li></ul></ul>
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Section 20.2 Oxidation Numbers <ul><li>OBJECTIVES </li></ul><ul><ul><li>Define oxidation and reduction in terms of a change in oxidation number, and identify atoms being oxidized or reduced in redox reactions. </li></ul></ul>
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Assigning Oxidation Numbers <ul><li>An “ oxidation number ” is a positive or negative number assigned to an atom to indicate its degree of oxidation or reduction. </li></ul><ul><li>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 </li></ul>
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Rules for Assigning Oxidation Numbers <ul><li>The oxidation number of any uncombined element is zero . </li></ul><ul><li>The oxidation number of a monatomic ion equals its charge . </li></ul>
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Rules for Assigning Oxidation Numbers <ul><li>The oxidation number of oxygen in compounds is -2 , except in peroxides, such as H 2 O 2 where it is -1. </li></ul><ul><li>The oxidation number of hydrogen in compounds is +1 , except in metal hydrides, like NaH, where it is -1. </li></ul>
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Rules for Assigning Oxidation Numbers <ul><li>The sum of the oxidation numbers of the atoms in the compound must equal 0. </li></ul>2(+1) + (-2) = 0 H O (+2) + 2(-2) + 2(+1) = 0 Ca O H
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Rules for Assigning Oxidation Numbers <ul><li>The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge. </li></ul>X + 3(-2) = -1 N O thus X = +5 thus X = +6 X + 4(-2) = -2 S O
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Reducing Agents and Oxidizing Agents <ul><li>Conceptual Problem 20.2, page 641 </li></ul><ul><li>An increase in oxidation number = oxidation </li></ul><ul><li>A decrease in oxidation number = reduction </li></ul>Sodium is oxidized – it is the reducing agent Chlorine is reduced – it is the oxidizing agent
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Section 20.3 Balancing Redox Equations <ul><li>OBJECTIVES </li></ul><ul><ul><li>Describe how oxidation numbers are used to identify redox reactions. </li></ul></ul>
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Section 20.3 Balancing Redox Equations <ul><li>OBJECTIVES </li></ul><ul><ul><li>Balance a redox equation using the oxidation-number-change method. </li></ul></ul>
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Section 20.3 Balancing Redox Equations <ul><li>OBJECTIVES </li></ul><ul><ul><li>Balance a redox equation by breaking the equation into oxidation and reduction half-reactions, and then using the half-reaction method. </li></ul></ul>
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Identifying Redox Equations <ul><li>In general, all chemical reactions can be assigned to one of two classes: </li></ul><ul><ul><li>oxidation-reduction, in which electrons are transferred: </li></ul></ul><ul><ul><ul><li>Single-replacement, combination, decomposition, and combustion </li></ul></ul></ul><ul><ul><li>this second class has no electron transfer, and includes all others: </li></ul></ul><ul><ul><ul><li>Double-replacement and acid-base reactions </li></ul></ul></ul>
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Identifying Redox Equations <ul><li>In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide: </li></ul><ul><li>N 2(g) + O 2(g) -> 2NO (g) </li></ul><ul><li>Is this a redox reaction? </li></ul><ul><ul><li>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. </li></ul></ul><ul><ul><li>Conceptual Problem 20.4, page 647 </li></ul></ul>YES!
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Balancing Redox Equations <ul><li>It is essential to write a correctly balanced equation that represents what happens in a chemical reaction </li></ul><ul><ul><li>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: </li></ul></ul><ul><ul><li>Use oxidation-number changes </li></ul></ul><ul><ul><li>Use half-reactions </li></ul></ul>
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Using Oxidation-Number Changes <ul><li>Sort of like chemical bookkeeping, you compare the increases and decreases in oxidation numbers. </li></ul><ul><ul><li>start with the skeleton equation </li></ul></ul><ul><ul><li>Step 1 : assign oxidation numbers to all atoms; write above their symbols </li></ul></ul><ul><ul><li>Step 2 : identify which are oxidized/reduced </li></ul></ul><ul><ul><li>Step 3 : use bracket lines to connect them </li></ul></ul><ul><ul><li>Step 4 : use coefficients to equalize </li></ul></ul><ul><ul><li>Step 5 : make sure they are balanced for both atoms and charge – Problem 20.5, 649 </li></ul></ul>
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Using half-reactions <ul><li>A half-reaction is an equation showing just the oxidation or just the reduction that takes place </li></ul><ul><li>they are then balanced separately , and finally combined </li></ul><ul><ul><li>Step 1 : write unbalanced equation in ionic form </li></ul></ul><ul><ul><li>Step 2 : write separate half-reaction equations for oxidation and reduction </li></ul></ul><ul><ul><li>Step 3 : balance the atoms in the half-reactions </li></ul></ul>(More steps on the next screen.)
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Choosing a Balancing Method <ul><li>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. </li></ul><ul><li>The half-reaction method works best for reactions taking place in acidic or alkaline solution . </li></ul>
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Using half-reactions <ul><li>continued </li></ul><ul><ul><li>Step 4 : add enough electrons to one side of each half-reaction to balance the charges </li></ul></ul><ul><ul><li>Step 5 : multiply each half-reaction by a number to make the electrons equal in both </li></ul></ul><ul><ul><li>Step 6 : add the balanced half-reactions to show an overall equation </li></ul></ul><ul><ul><li>Step 7 : add the spectator ions and balance the equation </li></ul></ul><ul><li>Rules shown on page 651 – bottom </li></ul><ul><li>Conceptual Problem 20.6, page 652 </li></ul>
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Electrochemical (Voltaic) Cells <ul><li>An apparatus that allows a redox reaction to occur by transferring electrons through an external connector </li></ul><ul><li>Redox reactions that occur spontaneously may be employed to provide a source of electrical energy </li></ul>
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<ul><li>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 </li></ul><ul><li>In a voltaic cell, a chemical reaction is used to produce a spontaneous electric current by converting chemical energy to electrical energy </li></ul>
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Cathode <ul><li>Cathode - The electrode where reduction occurs </li></ul><ul><li>In an electrochemical (voltaic) cell, the cathode is the POSITIVE electrode </li></ul>
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Anode <ul><li>Anode – the electrode where oxidation occurs </li></ul><ul><li>In an electrochemical (voltaic) cell, the anode is the NEGATIVE electrode </li></ul><ul><li>The anode is the more active metal (according to table J) </li></ul>
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Electrolytic Cells <ul><li>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. </li></ul><ul><li>Redox reactions that do not occur spontaneously can be forced to take place by supplying energy with an externally applied electric current </li></ul><ul><li>The use of an electric current to bring about a chemical reaction is called electrolysis </li></ul><ul><li>In an electrolytic cell, an electric current is used to produce a chemical reaction </li></ul>
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Electrodes <ul><li>Cathode – negative electrode (reduction takes place here) </li></ul><ul><li>In electrolytic cells POSITIVE ions are REDUCED at the cathode </li></ul>
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Electrodes <ul><li>Anode – Positive anode (oxidation takes place here) </li></ul><ul><li>In electrolytic cells NEGATIVE ions are OXIDIZED at the anode. </li></ul><ul><li>**NOTE: The charge of the electrode is the opposite in electrochemical cells** </li></ul>