Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint
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Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint

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    Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint Chemistry - Chp 20 - Oxidation Reduction Reactions - PowerPoint Presentation Transcript

    • CHAPTER 20“Oxidation-Reduction Reactions” LEO SAYS GER
    • Section 20.1The Meaning of Oxidation and Reduction (called “redox”)OBJECTIVES Define oxidation andreduction in terms of the lossor gain of oxygen, and theloss or gain of electrons.
    • Section 20.1The Meaning of Oxidation and Reduction (Redox)OBJECTIVES State the characteristics ofa redox reaction and identifythe oxidizing agent andreducing agent.
    • Section 20.1The Meaning of Oxidation and Reduction (Redox)OBJECTIVES Describe what happens toiron when it corrodes.
    • Oxidation and Reduction (Redox) Early chemists saw “oxidation”reactions only as the combination ofa material with oxygen to producean 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) But, not all oxidation processesthat 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) A process called “reduction” is theopposite of oxidation, and originallymeant the loss of oxygen from acompound Oxidation and reduction always occursimultaneously The substance gaining oxygen (orlosing electrons) is oxidized, while thesubstance losing oxygen (or gainingelectrons) is reduced.
    • 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- (MgS)•The magnesium atom (which has zero charge) changes to amagnesium ion by losing 2 electrons, and is oxidized to Mg2+•The sulfur atom (which has no charge) is changed to asulfide ion by gaining 2 electrons, and is reduced to S2-
    • Oxidation and Reduction (Redox) 0 0 +1 −1 2 Na + Cl 2 → 2 Na ClEach sodium atom loses one electron: 0 +1 − Na → Na + eEach chlorine atom gains one electron: 0 −1 − Cl + e → Cl
    • LEO says GER : Lose Electrons = Oxidation 0 +1 −Na → Na + e Sodium is oxidized Gain Electrons = Reduction0 −1 −Cl + e → Cl 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 is the Mg is oxidized: loses e-, becomes a Mg2+ ionreducing agent Mg(s) + S(s) → MgS(s)S is the oxidizing agent S is reduced: gains e- = S2- ion
    • Oxidation and Reduction (Redox) Conceptual Problem 20.1, page 634 It is easy to see the loss and gain ofelectrons in ionic compounds, but whatabout covalent compounds? In water, we learned that oxygen ishighly 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: +1 +5 −2 +1 −1 +1 −1 +1 +5 −2Ag N O 3 (aq ) + Na Cl (aq ) → Ag Cl ( s ) + Na N O 3 (aq) +1 −2 +1 +1 +6 −2 +1 +6 −2 +1 −22 Na O H (aq ) + H 2 S O 4 (aq ) → + Na 2 S O 4 (aq ) + H 2 O(l )
    • Corrosion•Damage done to metal is costly toprevent and repair•Iron, a common construction metal oftenused in forming steel alloys, corrodes bybeing 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 NumbersOBJECTIVES Determine the oxidationnumber of an atom of anyelement in a pure substance.
    • Section 20.2 Oxidation NumbersOBJECTIVES Define oxidation andreduction in terms of achange in oxidation number,and identify atoms beingoxidized or reduced in redoxreactions.
    • 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 Numbers1) The oxidation number of any uncombined element is zero.2) The oxidation number of a monatomic ion equals its charge. 0 0 +1 −1 2 Na + Cl 2 → 2 Na Cl
    • Rules for Assigning Oxidation Numbers3) 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. +1 −2 H2O
    • Rules for Assigning Oxidation Numbers5) The sum of the oxidation numbers of the atoms in the compound must equal 0. +1 −2 +2 −2 +1 H2O Ca (O H ) 2 2(+1) + (-2) = 0 (+2) + 2(-2) + 2(+1) = 0 H O Ca O H
    • Rules for Assigning Oxidation Numbers6) The sum of the oxidation numbers in the formula of a polyatomic ion is equal to its ionic charge. ? −2 ? −2 − 2− N O3 S O4 X + 3(-2) = -1 X + 4(-2) = -2 N O S O thus X = +5 thus X = +6
    • Reducing Agents and Oxidizing Agents• Conceptual Problem 20.2, page 641• An increase in oxidation number = oxidation• A decrease in oxidation number = reduction 0 +1 − Na → Na + eSodium is oxidized – it is the reducing agent 0 −1 − Cl + e → ClChlorine is reduced – it is the oxidizing agent
    • Trends in Oxidation and ReductionActive metals: Lose electrons easily Are easily oxidized Are strong reducing agentsActive nonmetals: Gain electrons easily Are easily reduced Are strong oxidizing agentsConceptual Problem 20.3, page 643
    • Section 20.3 Balancing Redox EquationsOBJECTIVES Describe how oxidationnumbers are used to identifyredox reactions.
    • Section 20.3 Balancing Redox EquationsOBJECTIVES Balance a redox equationusing the oxidation-number-change method.
    • Section 20.3 Balancing Redox EquationsOBJECTIVES Balance a redox equationby breaking the equation intooxidation and reduction half-reactions, and then using thehalf-reaction method.
    • Identifying Redox Equations In general, all chemical reactions canbe assigned to one of two classes:1) oxidation-reduction, in which electrons are transferred: • Single-replacement, combination, decomposition, and combustion2) 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 andoxygen react to form nitrogen monoxide: N2(g) + O2(g) → 2NO(g)•Is this a redox reaction? YES! •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
    • Balancing Redox Equations It is essential to write a correctlybalanced equation that representswhat 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 changes2) Use half-reactions
    • Using Oxidation-Number Changes Sort of like chemical bookkeeping, youcompare the increases and decreases inoxidation 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 showingjust the oxidation or just the reduction thattakes place they are then balanced separately, andfinally 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 Method1) 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-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 ofElectrochemical Cells
    • CathodeCathode - 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 ofElectrolytic 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**