An increase in oxidation state occurs during oxidation Oxidation is the loss of electrons For example:

Zn(s) ĺ Zn2+DTH– Oxidation state of zinc changes from 0 to 2; it increases. 2I–DTĺ I₂DTH– Oxidation state of iodine changes from –1 to 0; it increases.

A decrease in oxidation state occcurs during reduction. Reduction is the gain of electrons. For example:

Ag+DTH– ĺ Ag(s) Oxidation state of silver changes from +1 to 0; it decreases. Cl₂(g) + 2e– ĺ 2Cl–DT 2[LGDWLRQVWDWHRIFKORULQHFKDQJHVIURPWR±LWGHFUHDVHV 4.5.2 ‡ 8QFRPELQHGHOHPHQWVKDYHDQR[LGDWLRQVWDWHRIHJ1D +₂ = 0).

 ‡ ,RQVKDYHDQR[LGDWLRQVWDWHHTXDOWRWKHFKDUJHRQWKHLULRQHJ1D+

= +1, S2– = –2).  ‡ 2[\JHQLQFRPSRXQGVKDVDFKDUJHRI±LQR[LGHVDQG±LQSHUR[LGHV

 ‡ +\GURJHQLQFRPSRXQGVKDVDFKDUJHRIZKHQFRPELQHGZLWKQRQPHWDOVHJLQ+₂S) and –1 when combined with metals (e.g. in NaH).

 ‡ 7KHR[LGDWLRQVWDWHRIDFRPSRXQGRUSRO\DWRPLFLRQLVWKHVXPRIWKHR[LGDWLRQVWDWHVRIDOOLWVDWRPV)RUD compound the sum is 0.

 ‡ 2[LGDWLRQLQYROYHVDQLQFUHDVHLQR[LGDWLRQVWDWH  ‡ 5HGXFWLRQLQYROYHVDGHFUHDVHLQR[LGDWLRQVWDWH 4.5.3 (a) +2 (b) +3 (c) 0 (d) –2 (e) –1 (f) +7 4.6.1 $JDOYDQLFYROWDLFHOHFWURFKHPLFDOFHOOLVDQDUUDQJHPHQWRIFKHPLFDOVDQGHTXLSPHQWWKDWDOORZVDVSRQWDQHRXVUHGR[ UHDFWLRQWRWDNHSODFHLQVXFKDZD\WKDWHOHFWULFLW\LVSURGXFHG,WFRQVLVWVRIWZRKDOIFHOOVR[LGDWLRQRFFXUVLQRQHDQG UHGXFWLRQLQWKHRWKHU(DFKKDOIFHOOFRQWDLQVDQHOHFWURGHLQDQHOHFWURO\WH

4.6.2 (a) reaction that involves the loss of electrons (b) reaction that involves the gain of electrons

(c) electron transfer reaction involving oxidation and reduction (d) a substance that causes oxidation and is itself reduced (e) a substance that causes reduction and is itself oxidised

4.6.3 ,QDJDOYDQLFFHOOWKHUHDUHWZRKDOIFHOOV,QRQHKDOIFHOOWKHUHLVDVSRQWDQHRXVR[LGDWLRQUHDFWLRQDQGLQWKHRWKHUD reduction reaction occurs. Oxidation occurs at the anode, releasing electrons. Reduction occurs at the cathode which gains electrons. A wire connects the two electrodes so that electrons can transfer from the anode to the cathode.

4.7.1 7ZRKDOIFHOOVDUHVHWXSHDFKFRQWDLQLQJDQHOHFWURGHLQDQHOHFWURO\WHVROXWLRQ

 $ZLUHFRQQHFWVWKHWZRHOHFWURGHVWRDOORZWKHHOHFWURQVSURGXFHGWRÀRZIURPRQHWRWKHRWKHU A salt bridge, saturated with potassium nitrate, is set up so that it dips into both electrolytes.

4.7.2 D (OHFWURQVÀRZIURPWKHDQRGHZKHUHR[LGDWLRQRFFXUVWRWKHFDWKRGHZKHUHUHGXFWLRQRFFXUV (b) V salt bridge zinc anode negative ions positive ions copper cathode electrons electrolyte, e.g. CuSO₄ electrolyte, e.g. ZnSO₄ Oxidation Zn Zn2+ _+2e– Zn is a reductant Reduction Cu2+ _+2e– Cu Cu2+ is an oxidant

4.8.1 (a) A device which carries electric current into and out of a cell. (b) An electrolyte is a chemical that conducts electric current. (c) The electrode where oxidation occurs.

(d) The electrode where reduction takes place. 4.8.2 (a) To produce an electric current.

(b) Release of electrons occurs during oxidation at the anode. These have to travel to the cathode. By having the anode DQGFDWKRGHLQVHSDUDWHFHOOVZHFDQXVHWKHÀRZRIHOHFWURQVDVDVRXUFHRIHOHFWULFFXUUHQWYROWDJH

4.9.1 (a) Standard reduction potential is a measure of the relative tendency of a substance to gain one or more electrons FRPSDUHGWRWKHVWDQGDUGK\GURJHQKDOIFHOO

These are measured under standard conditions of 25°C, 1 atmosphere pressure and using a 1 mol/L electrolyte solution. The larger the E value, the greater the oxidising power of a substance.

 E $KDOIFHOOFRQWDLQLQJDVWDQGDUGK\GURJHQHOHFWURGHDVVKRZQLQWKHGLDJUDPEHORZ inert metal e.g. Pt H2 gas 1 mol L–1_[H+_] (c) Voltmeter

4.9.2 (a) The redox table is a list of standard reduction potentials (E ). A copy of this table can be found at the back of this book.  E 5HGXFWLRQUHGXFHGUHYHUVLQJOHIWGRZQÀXRULQHHOHFWURQVWRSJLYHXSUHGXFHGORZHUGRZQK\GURJHQR[LGLVLQJ

4.9.3 (a) –0.44 volts

(b) (i) Al(s) U Al3+ + 3e– +1.68 V (ii) Cu(s) U Cu2+ + 2e– –0.34 V (c) There is no reaction with iron in a solution of aluminium ions.

For the iron in copper sulfate solution, the reaction is as follows: Cu2+ + 2e– U Cu(s) +0.34 V

Fe(s) U Fe2+ + 2e– +0.44 V

Cu2+ + Fe(s) U Fe2+ + Cu(s) +0.78 V E

Zn(s) U Zn2+ + 2e– +0.76 V

Cu2+ + Zn(s) U Zn2+ + Cu(s) +1.10 V E

Diagram of lead acid cell:

H SO and water 2 4 e– e– Pb anodes PbO cathodes 2

Chemistry of lead acid cell: Anode: lead plates Pb ĺ Pb2+ + 2e–

Cathode: lead(IV) oxide

PbO₂ + 4H+ + 2e– ĺ Pb2+ + 2H₂O

Pb ions combine with sulfate ions and form lead sulfate. Electrolyte: 6 mol L–1 sulfuric acid

(b) Various, e.g.

Diagram of lithium cell:

Electrolyte Lithium Anode contact

+

-

Positive electrode (mixture used varies with type)

Separator Cathode contact/case Seal

Chemistry of lithium cell:

Anode: lithium

Li ĺ Li+

+ e–

Cathode: carbon

Reaction involves silver chromate or iodine, e.g. I₂ + 2e– ĺ 2I–

4.10.2

Lead acid cell Lithium cell

Cost and Practicality Expensive, but long lasting.

Used as car batteries. Practical for this purpose as they do not need to be portable.

They can be recharged. The reactions above are reversed by the car generator forcing current back into the battery. They work in a wide range of temperatures.

Expensive compared to other batteries. Used in cardiac pacemakers, cellular phones, watches , computers and cameras.

Practical for these uses as they are long-lasting, rechargeable and high voltage.

Lithium batteries are very light and deliver more power (about 3 V) than dry cell and alkaline batteries (1.5 V).

Impact on society Their development meant that cars could be started much

more easily and reliably – they did not have to be cranked. They improved the capability of people to move around and travel long distances.

The development of these long-lasting, rechargeable, very light batteries that produce a constant, relatively high voltage has led to the development of medical applications such as cardiac pacemakers which have saved lives. Their small size and portability have allowed the development of smaller electronic devices such as cameras, watches and phones. These have improved our ability to communicate over distance.

Environmental impact

Contains concentrated sulfuric acid (about 6 mol L–1

) which must be disposed of safely as it is highly corrosive. Also lead is a toxic heavy metal so must be disposed of carefully.

Lithium must be transported and disposed of safely to avoid environmental damage.