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Redox: Oxidation and Reduction

Redox: Oxidation and Reduction. Definitions. Oxidation : loss of e- in an atom increase in oxidation number (ex: -1  0 or +1  +2 ) Reduction : gain of e- by an atom decrease in oxidation number (ex: 0  -1 or +2  +1). Trick to remember this :. OIL RIG

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Redox: Oxidation and Reduction

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  1. Redox: Oxidation and Reduction

  2. Definitions • Oxidation: loss of e- in an atom increase in oxidation number (ex: -1  0 or +1  +2) • Reduction: gain of e- by an atom decrease in oxidation number (ex: 0 -1 or +2  +1)

  3. Trick to remember this: • OIL RIG • Oxidation is loss • Reduction is gain

  4. Another trick…. • LEO (the lion says…) GER • LEO- Lose Electrons Oxidation • GER- Gain Electrons Reduction

  5. Assigning Oxidation Numbers (What are they?) Definition: The apparent charge assigned to an atom of an element. • Numbers may be +, -, or 0 • These are not ionic charges, but maybe the same as the ionic charge. To distinguish oxidation numbers from charges on ions, the sign of the oxidation number precedes the number. Example: Ionic Charge = Mg2+ Oxidation Number = Mg+2

  6. Assigning Oxidation Numbers (Why do we use them?) • During an Oxidation-Reduction reaction electrons are exchanged or transferred. • Its not always possible to determine whether atoms have exchanged electrons by simply reading an equation. • Chemists have devised a system that makes it easier to keep track of the number of electrons lost or gained during a reaction. • Positive, negative or neutral values known as Oxidation Numbers or States can be assigned to atoms. • Oxidation numbers help show which atoms and how many electrons are either gained or lost by an atom. • By using these values we can define what, if any, substance is being oxidized and reduced during the reaction.

  7. Assigning Oxidation Numbers (How do we assign them?) • You must!!! Learn the rules for assigning oxidation states to atoms in an equation. • These numbers will be used to identify what is being oxidized and reduced. Rules for assigning Oxidation #’s • Each uncombined or lone element has an oxidation number of “0”. Ex. Na + Cl2 >>> NaCl – both Na and Cl2 have oxidation numbers of “0” 2. Monatomic ions have an oxidation number equal to its ionic charge. Ex. Na + Cl2 >>> NaCl – the Na ion in NaCl will have an oxidation # of +1 the Cl ion in NaCl will have an oxidation # of -1. 3. The metals in group 1 always have oxidation numbers of +1 in compounds, and the metals in group 2 always have an oxidation number of +2 in compound. 4. Fluorine always has an oxidation number -1 in compound. All the halogens (group 17) will be -1 if they are the most electronegative element in the compound.

  8. Assigning Oxidation Numbers (How do we assign them? continued) 5. Hydrogen is +1 in compounds unless it is combined with a metal, in which it is -1. Example: Hydrogen is +1 in HCl but -1 in LiH 6. Oxygen is usually -2 in compounds. Except when its is combined with fluorine, which is more electronegative, its +2. Example: Oxygen is -2 in H2O, and +2 in OF2. In the peroxide polyatomic ion (O22-) oxygen is -1. These 6 rules can be used along with the following 2 additional rules to calculate oxidation numbers for other elements in compounds or polyatomic ions in an equation. 7. The sum of the oxidation numbers in all compounds must be zero. 8. The sum of the oxidation numbers in polyatomic ions must be equal to the charge on the ion.

  9. Practice Assigning Oxidation Numbers • Example # 1 HNO3 • What are the oxidation numbers of the atoms? • Use the first six rules to assign oxidation numbers first then if you still have some to assign make sure the sum of the oxidation numbers is 0 (rule 7) • Rule 5 states H must be +1 • Rule 6 states each oxygen must be -2, and since there are three oxygen atoms they must total (3 x -2) -6 • So the total of 1 hydrogen and 3 oxygen atoms is (+1 + -6) -5 which means using rule 7 each nitrogen must be +5 to make the total charge of the compound “0”. • Final answer H = +1 • N = +5 • O = -2

  10. More Practice Assigning Oxidation Numbers • What is the oxidation number of chromium in the dichromate ion? • Cr2O72- • Use as many of the first six rules as possible • O has an oxidation number -2 (rule6), producing a total of (7 x -2) -14 for the seven O atoms. • Using rule 8 the sum of the oxidation numbers must be equal to the charge of the polyatomic ions charge • 2 x (the oxidation of Cr) + (-14) = -2 • 2 x (the oxidation of Cr) must = +12 • oxidation of Cr = +12/2 • oxidation of Cr = +6 Fill in chart on the next page!!!!

  11. Oxidation Losing electrons Any chemical change in which the oxidation numberincreases. Example 2Mg(s) + O2(g)→ 2MgO(s) (Magnesium loses, oxygen gains) *In many cases there is the addition of oxygen.

  12. Rusting • Occurs when metals react with oxygen. (accelerated by water and salts). • Certain metals corrode more than others (iron more than aluminum).

  13. Reduction Gaining electrons. Any chemical change in which an oxidation number of an element decreases or becomes negative. Example: 2Fe2O3(s) + 3C (s) 4Fe (s) + 3CO2

  14. Metal Ores • Occur naturally in combination or in oxidized state. • To obtain a purified metal, it must be “reduced” from its ore.

  15. Practice Given a redox reaction… • Which is being oxidized? • Which is being reduced? • Magnesium + Oxygen yields Magnesium oxide • 2Mg0 + O20 2Mg+2O-2 • Zinc + Copper (II) Sulfate Yields Zinc Sulfate + Copper • Zn0 + Cu+2S+6O-24 Zn+2S+6O-24 + Cu0

  16. Redox Reactions • Oxidation cannot occur without reduction. • When electrons from one atom are lost, they must be gained by another. • We call reactions that involve oxidation and reduction REDOX reactions. • Not all reactions are redox reactions. To determine whether or not a reaction is redox, assign oxidation numbers to each atom, both on the reactant and product side. • If there is a change in oxidation number for any atom, the reaction is redox.

  17. Redox Reactions (Continued) • If an element exists alone on one side and in a compound on the other, then it is a Redox Reaction. • Most single replacement, decomposition and synthesis reactions are Redox • Double Replacement Reactions are NOT Redox

  18. Reducing Agents: Metals • The reducing agent is the species being oxidized. • Oxidation number increases.

  19. Oxidizing Agents: Non-metals The oxidizing agent is the species being reduced. Oxidation number decreases.

  20. Half-Reactions • A half-reaction shows either the oxidation or reduction portion of a redox reaction, including the electrons gained or lost. • A reduction half-reaction shows an atom or an ion gaining one or more electrons while its oxidation number decreases • Ex. Na+1 + 1e- Na0 • An oxidation half-reaction shows an atom or an ion losing one or more electrons while its oxidation number increases. • Ex. Na0 Na+1 + 1e-

  21. Half-Reactions (continued) • Half-reactions must show conservation of mass and charge when written properly. • In a half-reaction there will be only one type of atom or ion shown on both reactant and product sides. • In a half-reaction the net charge on both sides of the equation must be equal, but it does not have to be “0”. • Ex. Sn+4 + 2e- Sn+2 Net Charge/side = +2 Mass is conserved with 1 Sn ion/side • Try the practice!!

  22. Writing Half-Reactions • To write the half-reactions for an equation such as MgCl2 + 2Na 2NaCl + Mg • Make sure the equation is balanced • Next assign oxidation numbers to each atom, then write a partial half-reaction to show the change in oxidation number including the coefficients from the balanced equation. Oxidation: 2Na0 2Na+1 Reduction Mg+2 Mg0 • Then place the correct number of electrons on one side of the equation to make the net charge equal on both sides. (remember electrons will always be a product in oxidation and a reactant in reduction) Oxidation: 2Na0 2Na+1 + 2e- Reduction: 2e- + Mg+2 Mg0 • When you have written a correct oxidation and reduction half-reaction the electrons lost (oxidation) should be equal to the electrons gained (reduction) • Try the practice!!!

  23. ELECTROCHEMICAL CELL • The exchange of electrons during a redox reaction can be useful to us. • One practical use of a redox reaction is in an electrochemical cell. An electrochemical cell involves a chemical reaction and the flow of electrons. • There are two common electrochemical cells. • Voltaic cells, which use a spontaneous reaction to produce a flow of electrons or an electric current. • Electrolytic Cells, which require an external power source to force a nonspontaneous reaction to occur.

  24. ELECTROCHEMICAL CELL (continued) • Electrochemical cells have two surfaces called electrodes that can conduct electricity. • An electrode is were oxidation or reduction will occur. • The electrode at which oxidation occurs is called the anode. • The electrode at which reduction occurs is called the cathode. • Use the phrase “AN OX - RED CAT” (anode oxidation – reduction cathode) to help you remember.

  25. Voltaic Cells • Voltaic cells take advantage of electron transfers during spontaneous reactions usually single replacement. • When you have two metals that differ in reactivity one will lose electrons easily (oxidation) and the other will gain the electrons (reduction) the other has lost. • If you separate the metals and connect them with a wire the flow of electrons from one metal to the other can be used to create an electric current.

  26. Voltaic Cells • The metals must be placed in separate ionic solutions that are connected by a salt bridge. • The metals act as electrodes the one that is more reactive, loses electrons (oxidation), will be the anode, the less reactive metal, gains electrons (reduction), is the cathode. • A voltmeter can be attached to the wire connecting the metals to detect electric current. • See the diagram of the Voltaic cell

  27. Voltaic Cell Diagram

  28. Voltmeter Drawing an Voltaic Cell Wire Electrodes Salt Bridge Solution A Solution B

  29. Voltaic Cells • Each vessel is called a half cell because a half reaction takes place here. • The salt bridge connects the two half cell solutions allowing ions to flow from one solution to the other to balance out the charge. • The electrodes (metals) are connected by a wire which acts as an electrical conductor. The flow of electrons through the wire creates an electric current. • This is a battery.

  30. Voltaic Cells & Table J • Oxidation and reduction occur at the electrodes. • To identify which electrode is your anode or the site of oxidation and which is your cathode the sight of reduction you must compare the reactivity of the metals on Table J in your reference table. • The more reactive metal will always lose electrons (oxidation) which makes it your Anode. • The less reactive metal will always gain electrons (reduction) it is always your cathode. • Remember “Red Cat An Ox”

  31. Table J • So in summary the metal higher up on Table J is your Anode (oxidation) the metal lower on Table J is your cathode (reduction). • Once you identify your anode and cathode you can identify the flow of electrons they flow from the Anode to the Cathode. • Try the practice questions on the next page !!!

  32. Electrolytic Cells • In a voltaic cell electrons flow spontaneously from anode to cathode. • The opposite occurs in an electrolytic cell, since the reaction is non-spontaneous there must be an electrical power source placed in the circuit to force the electrons from anode to cathode. • When electricity is used to force a reaction to occur it is called electrolysis.

  33. Electroplating • The most common use of electrolytic cells is an electroplating cell. • Electroplating involves plating a small layer of usually a precious metal on another metal. • The material to be plated for example a spoon or a piece of jewelry, would be the cathode were reduction occurs. • The metal being used to plate for example gold or silver would be the anode were oxidation occurs.

  34. Electrolytic Cell For Electroplating Power Source Anode Cathode • Anode-Oxidation • Ag0 Ag+1 +1e- • Cathode-Reduction • Ag+1 + 1e- Ag0

  35. Electrolysis • Electrolysis can be used to break up compounds to reform the elements it is made up of. • For instance if you wanted to break up water molecules into hydrogen and oxygen gas you could use an Electrolytic Cell. • Since the reaction will be non-spontaneous a power source must be used to force the reaction to occur. • The power source is usually a battery. • Reaction: 2H2O + electrical energy >>>> 2H2 + O2

  36. Electrolytic Cell for Electrolysis Power Source • Write the half-reactions: • Reduction: 4H+1 + 4e- 2H2 (Cathode) • Oxidation: 2O-2 O2 + 4e- (Anode)

  37. Comparing Voltaic and Electrolytic Cells Voltaic and Electrolytic cells have several similarities and differences. • Similarities • Both use redox reactions • The anode is the site of oxidation • The cathode is the site of reduction • The electron flow through the wire is from anode to cathode • Differences • Voltaic cells use spontaneous reactions to produce energy (voltmeter) • Electrolytic use non-spontaneous reactions that requires energy (power source) • Voltaic cells the anode is negative and the cathode is positive • Electrolytic cells the anode is positive and the cathode is negative.

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