Oxidation = a loss of electrons; an element which loses electrons is said to be oxidized. Reduction = a gain of electrons; an element which gains electrons is said to be reduced. Oxidizing Agent = a substance which "takes" electrons from another substance causing that substance to lose electrons or be oxidized; an oxidizing agent is itself reduced. Reducing Agent = a substance which "gives" electrons to another substance causing that substance to be reduced; a reducing agent is itself oxidized. If a substance gives up electrons readily, it is said to be a strong reducing agent. Its oxidized form, however, is normally a poor oxidizing agent. If a substance gains electrons readily, it is said to be a strong oxidizing agent. Its reduced form is a weak reducing agent. Oxidation Number = the charge which an atom has, or appears to have, when electrons are counted according to arbitrary rules (bookkeeping method). 1
Rules for assigning Oxidation Numbers: 1. The oxidation number of a free element is zero. 2. In ionic compounds, the oxidation number of an element is equal to the charge of the ion. 3. All metals of Group 1 (alkali metals) form only +1 ions and their oxidation number is +1 in all their compounds. 4. All metals of Group 2 (alkaline earth metals) form only +2 ions and their oxidation number is +2 in all their compounds. 5. In compounds of the halogens (Group 17) containing only one other metallic element, the halogen is assigned an oxidation number of 1. 6. The oxidation number of oxygen in all its compounds is 2 except in peroxides (such as H 2O 2) and superoxides (which contain an oxygenoxygen bond) where it is 1. 7. The oxidation number of hydrogen is +1 in all its compounds except the metal hydrides (e.g. LiH, CaH 2, etc.) where it is 1. 8. All oxidation numbers must be consistent with the principle of conservation of charge: a) For neutral compounds, the arithmetic sum of the oxidation numbers of all the atoms in the compound must be zero. b) For molecular ions (radicals) the oxidation numbers of all the atoms must add up to the charge of the ion. Another Definition: Oxidation = any chemical change in which there is an increase in oxidation number. The increase results from a loss of electrons (e.g. 3 to 1, 1 to +2). Reduction = any chemical change in which there is a decrease in oxidation number. The decrease results from a gain of electrons (e.g. 3 to 5, 1 to 2). 2
Problems in Assigning Oxidation Numbers example: dichromate ion, Cr 2O 7 2 Cr 2O 7 2 example: ammonium ion, NH 4 + NH 4 + example: sodium sulfate: Na 2SO 4 Na 2SO 4 example: copper(ii)nitrate: Cu(NO 3) 2 Cu(NO 3) 2 3
Homework Assignment 1: Assign oxidation numbers to each of the following: oxidation number(s) name 1. Cl 2 2. Cl 3. Na 4. Na + 5. H 2S 6. H 2SO 4 7. NO 3 8. CrO 4 2 9. NH 4Cl 10. NH 3 11. CaH 2 12. Na 2O 2 4
For each of the following unbalanced reactions identify: 1. substance oxidized (SO) 2. substance reduced (SR) 3. oxidizing agent (OA) 4. reducing agent (RA) 1. H 2 (g) + N 2 (g) > NH 3 (g) 2. C(s) + H2O(l) --> CO(g) + H2(g) 3. H2O2(aq) + PbS(s) --> PbSO4(s) + H2O(l) 4. HNO3(aq) + I2(s) --> HIO3(aq) + NO2(g) + H2O(l) 5
Balancing Redox Reactions You are familiar with balancing chemical quations by using the inspection method balancing by adjusting coefficients. Redox reactions are more complicated in that the numbers of atoms of each element may be balanced by inspection but the charge may not be. We will examine two methods for balancing more xcomplex redox equations. 1. Oxidation Number Method example: Cu(s) + HNO 3(aq) NO(g) + Cu(NO 3 ) 2 (aq) + H 2 O(l) 1. Assign oxidation numbers: Cu + HNO 3 NO + Cu(NO 3 ) 2 + H 2 O 1. Write the half reactions: 2. oxidation: 3. reduction: 4. Adjust coefficients to balance electrons: 5. oxidation: 6. reduction: 7. Transfer coefficients back to skeleton (unbalanced) equation: 8. Cu(s) + HNO 3(aq) NO(g) + Cu(NO 3 ) 2 (aq) + H 2 O(l) 9. Balance the rest by inspection (note: original coefficients may need to be changed): 6
Homework Assignment 3 Balance the following redox reactions and indicate SO, SR, OA, and RA 1. CrCl 3 + MnO 2 + H 2 O MnCl 2 + H 2 CrO 4 2. NH 3 + CuO Cu + H 2 O + N 2 7
3. HNO 3 + H 2 S H 2 SO 4 + NO 2 + H 2 O 4. Zn + H 2 SO 4 ZnSO 4 + SO 2 + H 2 O 8
Electrochemistry (Part II) Electrolytic Cell use electrical energy from a battery to make a nonspontaneous redox reaction occur Voltaic Cell (also called galvanic cell or battery) use a spontaneous redox reaction to produce electric energy (i.e. an electric current) Chemists have set up a Reference Table comparing oxidizing strengths of different substances based of the standard hydrogen half cell. The table is reproduced on the next page. Review this table and note the following: Electrolytic Cells 1. the reactions are all written in the reduction direction 2. substances are arranged in order of decreasing oxidizing strength (i.e. F 2 is strongest oxidizing agent; Li + weakest oxidizing agent) 3. you can couple any two reactions but one of them must be reversed 4. voltage is an intensive property it does not depend on the number of electrons transferred 5. a net volatege (E o ) = + represents a spontaneous redox reaction; a net E o = represents a nonsponatneous redox reaction An electrolytic cell involves electrolysis or electroplating. Let s first consider electrolysis reactions. Again, these reactions involve using an electric current (supplied by a battery) to force a nonspontaneous redox reaction to occur. The electrode at which oxidation occurs is designated the anode while the electrode at which reduction occurs is designated as the cathode. (Students can remember this by An Ox and a Red Cat!). 9
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Problem 1: Consider the electrolysis of molten sodium chloride. Why must the sodium chloride be in the molten state? Diagram an electrolytic cell illustrating this process. In the diagram label the following: anode, cathode, battery, direction of electron flow thru the wires. Label the electrodes with + and signs. Write the half reactions that occur at each electrode, the net reaction and calculate the net E o. Cathode: Anode: Net reaxn: 11
Problem 2: Consider the electrolysis of water. Why must a salt like sodium sulfate be added? 1. What will be the color of litmus in a solution of sodium sulfate? 1. Write the half reaction that occurs at the anode and what will be the color of the indicator at the anode as the reaction proceeds? 2. Write the half reaction that occurs at the cathode and what will be the color of the indicator at the cathode as the reaction proceeds? 3. What is the net reaction and the net voltage? 12
Problem 3: Consider the electrolysis of aqueous potassium iodide. Now we have competing species! H 2O, K +, and I In the anode half reaction H 2O vs I which is more readily oxidized? In the cathode half reaction: H 2O vs K + which is more readily resduced? Anode: Cathode: Net: 13
Electroplating Let s consider electroplating a fork with silver metal. To electroplate an object keep the following in mind: 1. The object to be plated is made the cathode (i.e. hooked up to the negative terminal of the battery) 2. The object is placed in a solution containing ions of the metal to be platted 3. The cathode is made of the platting metal (i.e. silver) cathode (reduction): anode (oxidation): 14
Voltaic Cells As mentioned before, a voltaic cell is one in which we use a redox reaction to generate an electric current. In this case we physically separate the two half reactions into two half cells. Each half cell consists of a container holding a metal electrode immersed in an aqueous metal salt solution (usually a metal nitrate salt). The two electrodes are connected by a wire and the two half cells are also connected by a salt bridge. Some terms (electrochem language) yoyu need to understand: Current the flow of electrons through a conductor (i.e. metal); current is measured in terms of coulombs (unit of electric charge) that pass per second. This is also known as the amperage (amps). Anode the electrode at which oxidation occurs; it is designated negative in an electrochemical cell. Cathode the electrode at which reduction occurs; it is designated positive in an electrochemical cell. Anions negatively charged ions (e.g. NO 3 ). Cations positively charged ions (e.g. Na + ). Internal Circuit the flow of electrons through a salt bridge or porous cup. External Circuit the flow of electrons through the wire and metal electrodes. Salt Bridge/Porous Cup function to maintain electrical neutrality in each half cell by allowing ions to migrate. 15
Let s consider how to build an electrochemical cell. One of the most common examples involves the reaction between zinc metal and aqueous copper(ii) ions: Zn o (s) + Cu 2+ (aq) Zn 2+ (aq) + Cu o (s) (blue) (colorless) Which substance is oxidized? Which substance is reduced? A schematic representation of the above system is: Zn Zn 2+ anode Cu 2+ Cu cathode What will the oxidation half cell consists of? What will the reduction half cell consist of? Anode Reaction: (note build up of positive charge in the oxidatio half cell due to [Zn 2+ ] increasaing) Cathode Reaction: (note the build up of negative charge in the reductiion half cell because Cu 2+ is reduced) net reaction: What happens to the mass of the zinc electrode? What happens to the mass of the copper electrode? 16
Describe the flow of ions through the salt bridge: 17
Problem: Sketch an electrochemical cell consisting of Mg and Sn metals. Write the half reactions that occur at each electrode, the net reaction, and calculate the net E o. On the diagram label each electrode, include the sign, label the salt bridge, show direction of electron flow (external circuit) and direction of ion flow (internal circuit). First, which is more easily oxidized, Mg or Sn (or which is more readily reduced, Mg 2+ or Sn 2+ ) Oxidation of Mg: reduction of Mg 2+ or Oxidation of Sn: reduction of Sn 2+ Anode: Cathode: Net: 18
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Lead Storage Battery Anode: Pb(s) + H 2SO 4(aq) PbSO 4(s) + 2H+(aq) + 2e Cathode: Overall: PbO 2(s) + H 2SO 4(aq) + 2e PbSO 4(s) + 2H 2O(l) Pb(s) + PbO 2(s) + 2H 2SO 4(l) 2PbSO 4(s) + 2 H 2O(l) 20
First common Dry Cell: Anode (ox): Cathode (red): 21
Corrosion Corrosion is a change of a metal to a compound, the most common example being an oxidation. For iron: oxidation: Fe o Fe 2+ + 2e reduction: ½ O 2 + H 2O + 2e 2OH net: Fe o + ½ O 2 + H 2O Fe(OH) 2 The Fe(OH) 2 outer layer easily flakes off exposing more Fe o. For aluminum: Al o Al 2O 3 The Al 2O 3 does not easily flake off and thus protects the inner Al o. For copper: Cu0 CuO or Cu 2O The greenish copper oxide does not easily flake off and protects the inner copper Methods of Combating Corrosion 1. Coating paint, enamel, grease, etc 1. Alloying for example Fe + Cr or Ni stainless steel 3. Cathodic Protection make the metal you want to protect the cathode; use a sacrificial metal as the anode. A sacrificial metal is one that is much more easily oxidized and relatively cheap, such as Zn or Mg. 22
The standard hydrogen half cell. When coupled with a Zn//Zn2+ half cell, one of two things can happen: a) The Zn atoms will be oxidized to aqueous Zn 2+ ions by the H + ions; the aqueous H + ions are then reduced to H 2(g) OR b) The Zn 2+ ions oxidize H 2 to H + (aq); the Zn 2+ ions arethen reduced to Zn(s) 23
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