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1. Electron Transfer Reactions • Electron transfer reactions are oxidation-reduction or redox reactions. • Results in the generation of an electric current (electricity) or be caused by imposing an electric current. • Therefore, this field of chemistry is often called ELECTROCHEMISTRY.

2. Terminology for Redox Reactions • OXIDATION—loss of electron(s) by a species; increase in oxidation number; increase in oxygen. • REDUCTION—gain of electron(s); decrease in oxidation number; decrease in oxygen; increase in hydrogen. • OXIDIZING AGENT—electron acceptor; species is reduced. (an agent facilitates something; ex. Travel agents don’t travel, they facilitate travel) • REDUCING AGENT—electron donor; species is oxidized.

3. You can’t have one… without the other! • Reduction (gaining electrons) can’t happen without an oxidation to provide the electrons. • You can’t have 2 oxidations or 2 reductions in the same equation. Reduction has to occur at the cost of oxidation LEO the lion says GER! ose lectrons xidation ain lectrons eduction GER!

4. Another way to remember • OIL RIG s s xidation ose eduction ain

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

6. 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.

7. 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

8. Rules for Assigning Oxidation Numbers • The oxidation number of any uncombined element is zero. • The oxidation number of a monatomic ion equals its charge.

9. Rules for Assigning Oxidation Numbers • The oxidation number of oxygen in compounds is -2, except in peroxides, such as H2O2 where it is -1. • The oxidation number of hydrogen in compounds is +1, except in metal hydrides, like NaH, where it is -1.

10. Rules for Assigning Oxidation Numbers • 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

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

12. Review of Oxidation numbers The charge the atom would have in a molecule (or an ionic compound) if electrons were completely transferred. • Free elements (uncombined state) have an oxidation number of zero. Na, Be, K, Pb, H2, O2, P4 = 0 • In monatomic ions, the oxidation number is equal to the charge on the ion. Li+, Li = +1; Fe3+, Fe = +3; O2-, O = -2 • The oxidation number of oxygen is usually–2. In H2O2 and O22- it is –1. 4.4

13. Oxidation numbers of all the atoms in HCO3- ? • The oxidation number of hydrogen is +1except when it is bonded to metals in binary compounds. In these cases, its oxidation number is –1. • Group IA metals are +1, IIA metals are +2 and fluorine is always –1. 6. The sum of the oxidation numbers of all the atoms in a molecule or ion is equal to the charge on the molecule or ion. HCO3- O = -2 H = +1 3x(-2) + 1 + ? = -1 C = +4 4.4

14. Oxidation and Reduction (Redox) • 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.

15. 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 a magnesium ion by losing 2 electrons, and is oxidized to Mg2+ • The sulfur atom (which has no charge) is changed to a sulfide ion by gaining 2 electrons, and is reduced to S2-

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

17. LEO says GER : Lose Electrons = Oxidation Sodium is oxidized Gain Electrons = Reduction Chlorine is reduced

18. 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 Mg2+ ion Mg is the reducing agent S is the oxidizing agent S is reduced: gains e- = S2- ion

19. Oxidation and Reduction (Redox) • 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)

20. Not All Reactions are Redox Reactions - Reactions in which there has been no change in oxidation numberare NOT redox reactions. Examples:

21. 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 electricallyconductive solutions that make electron transfer easy

22. 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 • Iron has an oxide coating, but it is not tightly packed, so water and air can penetrate it easily

23. 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

24. 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

25. 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

26. Balancing Redox Equations • OBJECTIVES • Describe how oxidation numbers are used to identify redox reactions.

27. Balancing Redox Equations • OBJECTIVES • Balance a redox equation using the oxidation-number-change method.

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

29. 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

30. Identifying Redox Equations • In an electrical storm, nitrogen and oxygen react to form nitrogen monoxide: • N2(g) + O2(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!

31. 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

32. 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

33. 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

34. 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

35. Choosinga 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.

36. Fe2+ + Cr2O72- Fe3+ + Cr3+ +2 +3 Fe2+ Fe3+ +6 +3 Cr2O72- Cr3+ Cr2O72- 2Cr3+ Balancing Redox Equations The oxidation of Fe2+ to Fe3+ by Cr2O72- in acid solution? • Write the unbalanced equation for the reaction ion ionic form. • Separate the equation into two half-reactions. Oxidation: Reduction: • Balance the atoms other than O and H in each half-reaction. 19.1

37. Fe2+ Fe3+ + 1e- 6Fe2+ 6Fe3+ + 6e- 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O 14H+ + Cr2O72- 2Cr3+ + 7H2O Cr2O72- 2Cr3+ + 7H2O Balancing Redox Equations • For reactions in acid, add H2O to balance O atoms and H+ to balance H atoms. • Add electrons to one side of each half-reaction to balance the charges on the half-reaction. • If necessary, equalize the number of electrons in the two half-reactions by multiplying the half-reactions by appropriate coefficients. 19.1

38. 14H+ + Cr2O72- + 6Fe2+ 6Fe3+ + 2Cr3+ + 7H2O 6Fe2+ 6Fe3+ + 6e- 6e- + 14H+ + Cr2O72- 2Cr3+ + 7H2O Balancing Redox Equations • Add the two half-reactions together and balance the final equation by inspection. The number of electrons on both sides must cancel. You should also cancel like species. Oxidation: Reduction: • Verify that the number of atoms and the charges are balanced. 14x1 – 2 + 6x2 = 24 = 6x3 + 2x3 • For reactions in basic solutions, add OH- to both sides of the equation for every H+ that appears in the final equation. You should combine H+ and OH- to make H2O. 19.1

39. CHEMICAL CHANGE --->ELECTRIC CURRENT • To obtain a useful current, we separate the oxidizing and reducing agents so that electron transfer occurs thru an external wire. This is accomplished in a GALVANIC or VOLTAIC cell. A group of such cells is called a battery. http://www.mhhe.com/physsci/chemistry/essentialchemistry/flash/galvan5.swf

40. Galvanic Cells anode oxidation cathode reduction - + spontaneous redox reaction 19.2