Oxidation & Reduction I.

Oxidation & Reduction

oxidation – atoms or ions gain a higher oxidation # reduction – atoms or ions drop to a lower oxidation # OIL RIG  Oxidation Is Loss, Reduction Is Gain ♦ oxidation  loss of electrons (increase oxidation #) ♦ reduction  gain of electrons (decrease oxidation #) ♦ ♦

♦

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

♦ oxidation: Mg  Mg2+ + 2e♦ reduction: S + 2e-  S-2 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. 2. 3. 4. 5. 6.

An uncombined element has an oxidation # of 0. A monatomic ion has an oxidation number equal to its charge. Fluorine is –1 in all compounds. Oxygen is –2 in almost all compounds (except in peroxides such as H2O2, where it is –1). Hydrogen is +1 in all compounds except when combined with metals (e.g. – NaH – sodium hydride), where it is –1. 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 the all atoms in the following rxns: (1) F2 + HCl  HF + Cl2 (2) SO2 + H2O  H2SO4 ♦ The Roman numeral in ionic compounds represents the oxidation state of the metal. Fe2O3 = iron (III) oxide  ex: Cu2CrO4 = copper (I) chromate NiSO4 = nickel (II) sulfate ♦ oxidation-reduction reaction (REDOX reaction) – chemical reaction in which elements undergo changes in oxidation numbers ♦ example: ♦

0 0 +1 -1 H2(g) + Cl2(g)  2HCl(g) H is oxidized (0  +1) Cl is reduced (0  -1) ♦

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)

♦ Half-reactions are also known as half-equations ♦ The electrons produced by the oxidation half-reaction go to drive the reduction half-reaction. ♦ oxidising agent – the substance that is reduced  In the previous example, Cl2 is the oxidising agent.  Oxidising agents accept the electrons that the oxidised 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 ♦ Corrosion is a prime example of a redox reaction. ♦ One specific example is the rusting of iron: 0 0 +2 -2 2Fe(s) + O2(g)  2FeO(s) (here, iron is oxidized from 0  +2; oxygen is reduced from 0  -2) 0 0 +3 -2 4Fe(s) + 3O2(g)  2Fe2O3(s) (here, iron is oxidized from 0  +3; oxygen is reduced from 0  -2)

combustion ♦ Combustion is another good example of a common redox reaction. ♦ Consider the combustion of methane as one example. -4 +4 +1

0

+4 -4 -2

+2 -2 +1

CH4(l) + O2(g)  CO2(g) + H2O(g)

II.

Reactivity

♦ It is possible to deduce whether or not a redox reaction will happen based upon a reactivity series (activity series) of elements. ♦ There are two separate reactivity series: one for metals and one for halogens.  Metals can replace other metals, and halogens can replace other halogens only.  How? In a single replacement reaction, if the metal which stands by itself is above the metal it would like to replace in the activity series, then the reaction will occur. If the metal by itself is below the metal in the salt, then it will not occur. Likewise, only if the halogen by itself is above the halogen it is trying to replace, will the reaction occur. ♦ Consider the following activity series:

Activity Series of the Elements Metals Li Rb K Ba Sr Ca Na Mg Al Mn Zn Cr Fe Cd Co Ni Sn Pb H2 Sb Bi Cu Hg Ag Pt Au

Halogens F2 Cl2 Br2 I2

♦ Now consider the following reactions, using the activity series as your guide: 1) Zn(s) + 2HCl(aq)  ZnCl2(aq) + H2(g) In this reaction, zinc is by itself, and is trying to replace hydrogen. Since Zn is above H (H2) on the activity series of metals, this reaction will occur and produce the products listed. 2) Cl2(g) + 2KBr(aq)  2KCl(aq) + Br2(l) •

In this reaction, chlorine is by itself, and is trying to replace bromine. Since Cl2 is above Br2 on the activity series of halogens, this reaction will likewise occur as written. 3) Cu(s) + Fe(NO3)2(aq)  NR •

In this particular reaction, copper is BELOW iron on the activity series. Thus Cu is not able to replace Fe, and no chemical reaction will occur. (NR = No Reaction) ♦ Of course all of these reactions that do occur are redox reactions, which is why we mention them in this section. Prove this in reactions #1 and #2 by: (1) identifying the substances oxidized and reduced; (2) identifying the oxidising and reducing agents; and (3) writing the oxidation and reduction half reactions. •

♦ In this unit, you will also have to do a lab where you experimentally deduce a reactivity series based on the chemical behaviour of a group of oxidising and reducing agents. You will be given test tubes, a series of metals, and concentrated hydrochloric acid only. Consider how you will conduct this laboratory exercise… ♦ Also – very important: You will not be given an activity series in your data booklet on the IB exam! Instead, refer to the Standard Electrode Potential data table (#15) in your data booklet. This data table lists a series of reduction potentials for reduction half reactions.  In other words, the half-reaction at the top of the chart shows that lithium is the least likely element to be reduced, while fluorine, at the bottom is most likely to accept electrons and be reduced.  Put another way, lithium, at the top is most easily oxidized.  At this point a light bulb comes on over your head as you realize that the table of Standard Reduction Potentials lists all metals in exactly the same order as they appear on the activity series, and all halogens in exact reverse order of the activity series…

IIIa.

Electrochemistry

electrode – a conductor used to establish electric contact with a non-metallic part of a circuit (such as between electrolyte solutions) Recall that electrolytes are solutes which conduct electricity in aqueous solution. ♦ anode – the electrode where oxidation occurs ♦ cathode – the electrode where reduction occurs ♦

Electrochemical Cells ♦

electrochemical cell (Voltaic Cell )– a system of electrodes and electrolytes where 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.

The Dry Cell (Battery) ♦ The zinc cup (shell) is the anode, and the carbon rod in the middle is the cathode. 0 +2 at the anode: Zn  Zn2+ + 2e(oxidation) +4 +3 at the cathode: 2MnO2 + 2NH4 + 2e-  Mn2O3 + 2NH3 + H2O

(reduction)

Electrolytic Cells ♦

electrolytic cell – 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.

♦

electrolysis – when an electric current drives an oxidation-reduction reaction

Electroplating Electroplating is one example of electrolysis. Consider silverplating of silverware as an example of electroplating. • You’ve got a silver (Ag) anode and a soluble silver salt solution (ex: AgNO3) • The cathode is the object to be plated. ♦ This same process can be done using copper (instead of silver) and is called copper plating. This is the particular process IB expects you to know. So just replace silver with copper in the previous example and you have copper plating! ♦

• •

0 +1 Silver is oxidized at the anode (Ag  Ag+ + e-) and Ag+ ions are kicked off the anode (the silver bar). +1 0 Silver is reduced at the cathode (Ag+ + e-  Ag) where silver coats the metal object (spoon in this case).

IIIb. Electrolysis of a Molten Salt ♦ Consider the following electrolytic cell, involving electrolysis of a molten salt ♦ The example we will use is the most common of the salts, sodium chloride. An important thing to note is that solid sodium chloride does not conduct electricity.  However, molten sodium chloride does. In this case we are dealing with molten sodium chloride. That simply means that we've taken sodium chloride and heated it up enough to where it is melted. When it melts, the sodium ions and the chloride ions can separate from one another somewhat, and they are free to move throughout the liquid. ♦ An idealized cell for the electrolysis of sodium chloride is shown in the figure below. A source of direct current is connected to a pair of inert electrodes immersed in molten sodium chloride. Because the salt has been heated until it melts, the Na+ ions flow toward the negative electrode and the Cl- ions flow toward the positive electrode.

 The voltage source (a battery in this case) is forcing electrons to flow from left to right. Electrons are forced onto the electrode on the right, where they are picked up by sodium ions (Na+ ions). When the sodium ions pick up these electrons, they produce sodium metal. Since this is reduction, that electrode is called the cathode.

 When Na+ ions collide with the negative electrode, the battery carries a large enough potential to force these ions to pick up electrons to form sodium metal. Na+ + e-

Negative electrode (cathode):

Na

 On the left side, electrons are being pulled off of the chloride ions to form chlorine gas, which would bubble away unless it was somehow captured. Since this is an oxidation process, the electrode on the left side of this particular diagram would be called the anode because that is where the oxidation occurs. Positive electrode (anode):

2 Cl-

Cl2 + 2 e-

 The net effect of passing an electric current through the molten salt in this cell is to decompose sodium chloride into its elements, sodium metal and chlorine gas. Electrolysis of NaCl: Cathode (-): Anode (+):

Na+ + e2

Cl-

Na Cl2 + 2 e-

♦ Remember that the Cathode is the negative electrode and the Anode is the positive electrode! ♦ The dotted vertical line in the center of the above figure represents a diaphragm that keeps the Cl2 gas produced at the anode from coming into contact with the sodium metal generated at the cathode. The function of this diaphragm can be understood by turning to a more realistic drawing of the commercial Downs cell used to electrolyze sodium chloride shown in the figure below.

♦ Chlorine gas that forms on the graphite anode inserted into the bottom of this cell bubbles through the molten sodium chloride into a funnel at the top of the cell. Sodium metal that forms at the cathode floats up through the molten sodium chloride into a sodium-collecting ring, from which it is periodically drained. The diaphragm that separates the two electrodes is a screen of iron gauze, which prevents the explosive reaction that would occur if the products of the electrolysis reaction came in contact. ♦ Sodium chloride is not the only salt that can be used in this process. We could use any ionic compound that can be melted to free up the ions so that they can move. ♦ Independent practice: Given the following salts, sketch a diagram showing the electrolysis of the molten salt. Indicate the source of the electric current and conductors, positive and negative electrodes (write the words “positive electrode”, “negative electrode”, “cathode”, and “anode”), movement of ions, and products, just like the diagram above. (1) Al2O3

(2) CaCl2

(3) ZnF2

Also, for the first two molten salt systems, explain in detail exactly what is happening from start to finish in the cell.