Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

Topic 13 – Oxidation–Reduction Reactions

I. Oxidation & ReductionOIL RIG : Oxidation Is Loss, Reduction Is Gain (of ELECTRONS)

oxidation = loss of electrons (increase in oxidation #) ex: Mg Mg2+ + 2e–

reduction = gain of electrons (decrease in oxidation #) S + 2e– S–2

Electrons are transferred from the substance oxidized to the substance reduced.

Independent Practice: Consider each of the following atom/ion pairs and indicate how many e–

have been gained or lost, and tell if that represents oxidation or reduction:

1) K K+ 3) Fe+3 Fe+2

2) F– F2 4) Mn+2 Mn+7

RULES FOR ASSIGNING OXIDATION NUMBERS:

1. Elements not combined with other elements have (Fe, Ar, H2, N2, P4, etc.) an oxidation # of 0.2. A monatomic ion has an oxidation number equal to its charge.3. Fluorine is –1 in all compounds.4. Oxygen is –2 in almost all compounds.

(except in peroxides, such as H2O2, where it is –1, or when combined with F)5. Hydrogen is +1 in almost all compounds.

(except when combined with metals, such as NaH, sodium hydride, where it is –1)6. The more electronegative element in a binary compound is assigned the oxidation number equal to

the charge it would have as an ion.7. The sum of the oxidation #s of all atoms in a neutral compound is 0.8. The sum of the oxidation #s of all atoms in a polyatomic ion is equal to the charge of the ion.

Assign oxidation #s to all elements in the following compounds:

1) CO2 3) CuCl 5) C6H12O6 7) Na2O2

2) CO32– 4) KMnO4 6) HNO3 8) CaH2

The Roman numeral in ionic compounds represents the oxidation state of the metal.

ex: Fe2O3 = iron (III) oxide (iron’s oxidation # here is +3)Cu2CrO4 = copper (I) chromate (copper’s oxidation # here is +1)NiSO4 = nickel (II) sulfate (nickel’s oxidation # here is +2)

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

II. Redox Equations Redox reactions (oxidation-reduction reactions) are chemical reactions in which elements undergo

changes in oxidation numbers. In redox reactions, there must an element that is oxidized and an element that is reduced. You

can’t have reduction unless you also have oxidation. The electrons have to travel from from the element which is oxidized to the element which is reduced. While it is usually two different elements that are oxidized and reduced, you will sometimes find a reaction where the same element is undergoing both processes.

2Na(s) + Cl2(g) 2NaCl(s)

Na Na+ + e– (= oxidation: sodium has gone from 0 to +1) Cl2 + 2e– 2Cl– (= reduction: chlorine has gone from 0 to –1)

overall:2Na 2Na+ + 2e– (oxidation half-reaction)

Cl2 + 2e – 2Cl – (reduction half-reaction) 2Na + Cl2 2Na+ + 2Cl– (redox reaction)

The electrons produced by the oxidation half-reaction go to drive the reduction half-reaction.

oxidizing agent – the substance that is reduced In the previous example, Cl2 is the oxidizing agent. Oxidizing agents accept the electrons that the oxidized species loses.

reducing agent – the substance that is oxidized In the previous example, Na is the reducing agent. Reducing agents lose electrons that go to drive the reduction half-reaction.

Corrosion and combustion are prime examples of redox reactions:

0 0 +2 -22Fe(s) + O2(g) 2FeO(s)

0 0 +3 -24Fe(s) + 3O2(g) 2Fe2O3(s)

-4 +4 0 +4 -4 +2 -2 +1 -2 +1 CH4(l) + O2(g) CO2(g) + H2O(g)

♦ Here are a few good movies illustrating the processes of oxidation and reduction:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/073_OxiReduPart1.MOV http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/074_OxiReduPart2.MOV

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

III. Practical Applications of Redox Reactions

ELECTROCHEMICAL CELLS There are two types of electrochemical cells:

voltaic cells, in which a spontaneous redox reaction produces electrons (an electrical current), which are harnessed to do work (like lighting up a flashlight)

electrolytic cells, in which an electric current drives a non-spontaneous redox reaction

We’ll look at voltaic cells first, but let’s first consider some terms you need to know:

electrode – a conductor used to establish electric contact with a non-metallic part of a circuit (such as between electrolyte solutions – remember that electrolytes are solutes which conduct electricity when dissolved in aqueous solution)

anode – the electrode where oxidation occurs cathode – the electrode where reduction occurs

To reiterate, then, a voltaic cell is a system of electrodes and electrolytes in which a spontaneous redox reaction produces an electrical current (which is then used as a source of electrical energy) In other words, an electrochemical cell converts chemical energy into electrical energy.

♦ A common example of a voltaic cell is the dry cell (better known as a battery).

The zinc cup (shell) is the anode, and the carbon rod in the middle is the cathode.

at the anode: 0 +2 Zn Zn2+ + 2e– (oxidation)

+4 +3 at the cathode: 2MnO2 + 2NH4 + 2e– Mn2O3 + 2NH3 + H2O (reduction)

half-reactions – occur at one electrode in an electrochemical cell Electrons produced by the half-reaction at the anode drive the other half-reaction at the cathode.

Note that in an electrochemical cell, the anode is the negative electrode and the cathode is the positive electrode.

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

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This illustrates an electrochemical cell formed between copper and silver immersed in solutions of their own salts. Electrons flow from copper (which is higher on the activity series, and thus more easily oxidized) to silver, which is more easily reduced. The salt bridge permits the flow of ions to maintain a balance of charges.

Zinc atoms lose electrons and enter the solution as Zn2+ ions (red arrow). Electrons flow through the external circuit from the Zn to the Cu electrode. At the copper cathode, Cu2+ ions gain electrons and deposit as solid Cu atoms (blue arrow). Ion movement in the solutions is noted by the black arrows. Porous plugs prevent the bulk flow of solution but permit the ions to migrate through the salt bridge.

These are two examples of wet cells, in which the electrolytes are liquids, instead of pastes (as in a dry cell).

Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

ELECTROLYTIC CELLS Again, an electrolytic cell is a system of electrodes and electrolytes where an electric current drives a

non-spontaneous redox reaction. In other words, an electrolytic cell converts electrical energy into chemical energy.

ELECTROPLATING is one example of electrolysis.

Say, for instance, you wanted to silverplate something (as in silverware, or jewelry).

You’ve got a silver (Ag) anode and a soluble silver salt solution (ex: AgNO3).

The cathode is always the object to be plated (since it’s the site of reduction).

0 +1 Silver is oxidized at the anode (Ag Ag+ + e–)

and Ag+ ions are kicked off the anode.

+1 0

Ag+ ions are reduced at the cathode (Ag+ + e– Ag) and silver coats the metal object.

You can also take the apparatus for a voltaic cell like we looked at earlier, and force electrons to flow in the “wrong” (non-spontaneous) direction by means of a battery.

In other words, you can force electrons from the less easily oxidized metal onto the less easily reduced metal. The electrons would naturally tend to flow in the opposite direction.

In this case, the less easily oxidized metal is forced to act as an anode, and the less easily reduced metal is forced to act as a cathode.

In an electrolytic cell, the cathode is negative, and the anode is positive – their polarities are reversed from the spontaneous voltaic cells.

♦ Here is a movie illustrating the electroplating of chromium onto copper:http://cwx.prenhall.com/petrucci/medialib/media_portfolio/text_images/078_Electroplating.MOV

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Name: ______________________Chemistry 1 Notes, 2012–13; P. Holloman

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