o·

1 . ~

Regents Chemistry Date: ___ _ Name: -----------~

• • •

•

• • • •

• • •

• Test Objectives for Unit 13: Oxidation/Reduction

Define oxidation . Define reduction . Assign oxidation-numbers to each element in a compound or a polyatomic ion .

0 Know the 7 rules for assigning oxidation numbers. Identify whether a chemical equation is a redox reaction or not.

o (Hint: what kind of reaction is always redox?) Determine what is being o~dized and what is being reduced in a redox reaction . Memorize "An Ox Ate a Red Cat" and ''Oil Rig." Identify the oxidizing agent and the reducing agent in a redox reaction . Write correct equations for the oxidation half-reaction and the reduction half....,reaction in a redox reaction. Balance a redox equation, by balancing electrons lost with electrons gained . Use Table J to predict whether a particular reaction will occur or not. In an electrochemical (or galvanic or voltaic) cell, a spontaneous chemical reaction is used to produce electricity. (spontaneous reaction -7 electricity) This is an exothermic process. What type of energy is produced?

• Identify the components of an electrochemical cell and describe their function. (half-cell, salt bridge, wire, electrode, anode where oxidation occurs, cathode where reduction occurs) Use Table J to determine which electrode is the anode and which electrode is the

'\......../ cathode. (The higher metal donates electrons, so it's the anode. The lower electrode is the site of reduction and will be the cathode.) Assign + and - signs to the electrodes. Opposites attract - the polarity is determined by the direction of electron flow.

~

• Describe what happens during the operation of a given half-cell. (Which electrode is dissolving, which electrode is gaining mass, which solution is getting niore concentrated, which solution is getting more dilute, which way and where are the electrons traveling, which way and where are the positive ions traveling) Write a half-reaction for the process occurring at each electrode.

• Define the term electrolysis. • Know that in an electrolytic cell, electricity is used to force a nonspontaneous reaction to

occur. (electricity~ nonspontaneous reaction). This is an endothermic process. \\That type of energy is absorbed?

• Identify the components of a fused salt cell and an electroplating cell and describe their function. (anode, cathode, power supply,+ and- electrodes - polarity is determined by power supply) Write half-reactions for the process occurring at each electrode.

11 Know these differences between voltaic and electrolytic cells: o Voltaic cell: the redox reaction is spontaneous and exothermic o Electrolytic cell:._ the redox reaction is ncmspontaneous and endothermic. o · Voltaic cell: Anode is negative, cathode is positive. o Electrolytic cell: Anode is positive, cathode is negative.

• Know these similarities between voltaic and electrolytic cells: o Both use redox reactions. o The anode is the site of oxidation. o The cathode is the site of reduction. o The electrons flow through the wire from the anode to the cathode.

Regents Chemistry: Vocabulary for Chapters 20 & 21 Name:

-~~~-----~~-~-~~~~~--

1. Oxidation

2. Reduction

3. Half-reaction

4. Redox reaction

5. Oxidizing agent

6. Reducing agent

7. Species

8. Anode

9. Cathode

10. Battery

11. Corrosion

12. Galvanic Cell

13. Electrolytic Cell

14. Half-Cell

15. Salt Bridge

16. Electroplating

17. Fused Salt Cell 6 ~'

Date: ------ ~

,~

~-

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REDOX AND ELECTROCHEMISTRY

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Name

Date Period ~~

Study the rules for assigning oxidation numbers and examine the sample problem below. Then determine the unknown oxidation state in each example.

RULES FOR AsSIGNING OXIDATION NUMBERS 1. Oxidation numbers for atoms that are free elements are always zero 2. The oxidation numbers of ions are !tie same as the charge on the ion 3. Some elements have only one oxidation state

a. group 1 metals always form 1+ ions and always have a +1 ·oxida!ion state

b. group 2 metals always form 2+ ions and always have a +2 oxidation state

4. Some elements usuaHy have a particular oxidation state a. oxygen has a -2 oxidation stale except in peroxides where it is -1

and in compounds with fluorine (OFJ where it is +2 b. hydrogen has a + 1 oxidation state except in hydrides with group 1

and group 2 metals 5. the sum of the oxidation numbers

a. in a compound it is always zero b. in a polyatomic ion it is equal to the charge on the ion

Sample Problem Find the oridation state of the elements in K2Cr20 7•

Element K Cr 0

Subscript 2 2 7 TOTAL

O:ddation state +I ? -2

Sum of oxidation +2 ?? -14 0 states

[a} potassitnn is a group one metal; i1s oxidation sta1:e is ahvays+l

[bJ m .. -ygen usually has an oxidation state of-2 [cJ the sum of oxidation states of each element is the product

of the subscript and the oxidation state {dJ find the sum of the oxidation states of chromium(??) by

setting the sum of ail the oxidation states to zero (+2i + ?? +_ (-14) = 0

?? =+!2 . ff] . find the oxidation state of chromium m by dlviding the .

sum (+12) by the subscript (21 +U _,.. 2 == +6

1. Chlorine in KCIO 4 I. ---

2. Nitrogen in Ba(NO.J1 2. ---

3. Phosphorus in CaiPOJ2 3. __ _

4. Manganese in LiMn04 4. __ _

5. Sulfur in l\..:.:.303 5. __ _

6. Chromium in CaCrO 4 6. ---

7. Sulfur in MgS:P3 7. --

8. Nitrogen in Zn(NO;J2 8. __ _

9. Chlorine in HC103 9. __ _

10 ... Carbon in CaC20 4 10. --

11. Sulfur in KHSO 4 11. __

·~

©Evan P. Silberstein, 2003

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I

I ASSIGNING OXIDATION NUMBERS Name -------\SSign oxidation numbers to all of-the elements in each of the compounds or ions below.~'"

1. HCI -

11. H2S03

.2. KN03 12. H2S04

'

3. OH- 13. Ba02

4. Mg3N2 14. KMn04

.'~

5. KC/03 15. LiH

6. Al(N03) 3 16. MnC?2

7. SB 17. OF2

8. H202 1s. so3

9. Pb02 19. NH3

10. NaHS04 20. Na '

f"':'hi.°'."'!.,..,..,l-'--, tr-t"J..,L r

"-....,./\.

\..._,.i

\-......_/

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REDOX AND ELECTROCHEMISTRY Date

0-rtalcjzifi~ · ~:xiS,ati®rt-~·®'uctic©n ~t.·acti©rtg

""When chemical bonds form. electrons are either 1 ost, gained or shared. Metals lose electrons. This is what happens when iron rusts. \¥hen the iron, a metal, combines 'vith o:\.·ygen, a non metal, to form rus~ it loses electrons. This process is called oxidation even \Vnen the nonmetal is not o::\.·ygen. Nonmetals

·gain electrons causing their oxidation states to go dovm.. This is called reduction. It is possible to tell what ·was oxidized and what \Vas reduced in a chemical reaction by checking the oxidation states of the elements before and after the reaction. The element that has an increase 1n oxidation state was oxidized while the one that has a decrease in oxidation state was reduced.

Example

2FeCl2 + CI2 --+ 2FeCI3 Fe+~-+ Fe+3 Iron was oxidized Cl~' -+ c1-1 Chlorine was reduced

-.. ~:~$3;~~$;~;;~*':~::::-:~:::-;;;;o:;;:;:;:::::;;~:=~~it:~~~::i:Z";.';::;;:~~~~~'i=~:i?~~~~lli:::~:::::-

Period ~=

For each of the examples below, determine the oxidation states of the elements on both sides o"f fie equation. Then determine which element was oxidized and which was reduced. Write your answer in the space provided.

Reaction Element:

·-Oxidized Reduced

Example: Cu + 2AgN03 -+ Cu{N03)i + 2Ag 0 +I -+5 -~ +:? +5 -:! 0 Cu Ag

· Cu + 2AgN03 --+ Cu(N03)~ + 2Ag

1. 2Mg+ O::! ~ 2Mg0

2. Zn+ 2HCl-+ ZnCl, + H,

\ ..

3. Fe,Q, + 3CO-+ 2Fe + 3CO, ..;. ;> • -

4. 2K:!Cr:P7 + 2H:P + 3S -+ 4KOH + 2Cr20 3 + 3S02

( ~ Go on to the next page.)

c;

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REDOX AND ELECTROCHEMISTRY Pag~

Reaction Element:

Oxidized Reduced

5. 2H:i0 + O:i-+ 2H:i02

6. 2KCI03 -+ 2KCI + 302

7. 4NaOH + Ca(OH),·+ C + 4CIO~-+ 4NaCIO, +.CaCO- + 3HD • - - - .:> .:;..

8. 3P + 5HN03 + 2H20-+ SNO + 3H3P04

9. 3Cu + 8HN03 -+ 2NO + 3Cu(NOJ1 + 4H:i0

10. 2PbS04 + 2H:P-+ Pb02 +Pb+ 2H2SO..- .·~

-

11. 4HC1 + Mn02 -+ MnCI2 + 2H:P + Cl2

12. 4NH3 + 502 -+ 4NO +.6H20

13. 16HCI + 2K.Mn04 -+ 8H20 +2KC1+2MnCI:+ 5Cl2 ..

14. Cu+ 2H2S04 -+ CuS04 +S02 + H20

15. 8HN03 + 6KI-+ 6KN03 + 312 + 2NO + 4H20

16. I:i + SHCIO + H20-+ 2HI03 + 5HC1

17. K2Cr20 7 + 3SnCI2 + 14HCI -+ 2CrC13 + 3SnC14 + 2KC1 + 7H20 ..

18. SnCl 2 + 2HgC12 -+ SnC14 + Hg2Cl2

;g Evan P. Silberstein, 2003

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·~

_;

·Regents Chemistry:· Half-Reactions Name: · Date:· ___ _

Half-reactions must conserve atoms and charge. That means the total number of atoms of . .

each element p:i.ust be the same on qqtb. sid~s of the half-reaction and the total charge must be the same on both sides.

For example, we cannot write 0-2 -+ 02. Neither atoms nor charge is conserved. First balance the atoms: 2 0-2 --+ 02. -

Second balance the charge. Th.ere is a total of.2(-2) or-4 on the left so there must be

....-..~'¢.= .. -

·nega~ve 4 on the right 2 0·2 -+ 02 + 4 e·. tl OTE. ~ Oic t.'!t'>~6~ f..) e>, or" ~ ~re

. e letP"t ~+ \S .Z.E.·F-t> ( c:;)

\......__,'

; ;-.

1) A4d electrons to complete the following half-reactions. Id.en~fy each half-reaction as oxidation or reduction .

. a) Ch + -+ 2 cr1

b) Sn+4 +· --+. sn+2

c) Pb+2 -+ Pb+4 +

d) As+5 + --+ As+2

e) P4 -+ 4 p+S +

For the equation: I

~\ . - .

~ Mg + S -+ MgS

a) Assign oxidation numbers for every element in the reactants and products.

b) What is oxidized? -------'---

c) Whatisreduced? --------­

_d) Write a half-reaction to represent oxidatio~:

e) Write a half-reaction to represent reduction:

· !) Zn(s) + 2 W1(aq) --+ Zn+2(aq) + H2(g)_

a) Assign oxidation numbe~ for_ every element in the reactantS and products.

· ._b) "Whatisoxidized? -------. . . .

\~. c) .~tis.reduc·ed? -------=---. -d) ·Write a half-~~a~tion to r_epi-eS~nt oxidati6~:·

j -: -~;,i{7 ~ ~::-~:L~ ~

! '!

e) '9lrliii e. M1£ rse.stiea ~ re:e:reee?!:t ?eS..e!i~s;

5) Cb(g) + 2 Br"1(aq) ...+ 2 Cr1(aq) .+ Br4(1) . .

,. a) Assign oxidation numbers for every element in the reactants and products . ._.. 1 ~ \ • t Ir• • "' 1"' • J • if. I• :- ,. fl!t ' : ... ·,, VO .. : • •

• ,. .. i:>}:·:What.is ozjdiz~? ....... ·------..

c)· Whatis·~educed? _. ------­

d) Write a'half-r~action to repre~ent oxidation:

· e) Write a half-reaction to represent reduction:.

6) Ca(s) +· 2 W1(aq) -:-* Ca+2(aq) + H2(g)

a) Assign oxidation numbers for every element in the reactants and prcd.ucts.

b) What is oridized? --------

c) What is reduced? _. --------

d) Write a half-reaction to represent oxidation:

e) Write a half-reaction to represent reduction:

7) 2 Al(s) + 3 Cu+2(aq) -7 · 2 At3 + 3 Cu(s)

a) Assign oxidation.numbers for every element in the reactants and produpts.

b) What is oxidized? --------

c) What is reduced? --------

d) Write a half-reaction to represent oxidation: -

e) Write a half-re~on to represent reduction: . ..

. ..

..

.. .. 'I

·,

. . . . ·;··. . . :.::.' - . . .. : .... 0 .... : .. - ,,. f' ". :' . ~ . -

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REDOX AND ELECTROCHEMISTRY Pag#

4. Sn + HN03 + H20 ... '.".'.t B~Sx103 +NO

5. K1v1n04 + HCl -+ KCI + MnC12 + H10 + Cl2

6. Fe(OH)2 + B:P2 -+ Fe(OH)J

7. Na+ H:P -+ NaOH + H::!

-~

8. Zn + HN03 -+ Zn(N03)::! + N02 + H::P

9. H:!01 --+ H::O + 0::

10. K,Cr,07 + H,O + S-+ SO,+ .KOH+ Cr,0, .;.. - ... - - ,;,

•t:; EYan P. Silberstein, 2003

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........ ., . .,,..,,,""r'

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/l%~J:!!1~1ll11jt)tW!W1l~·Mf.'i·?s,¥•1rl.''."'t·1·?~?.\·tll<~;i*/ .. ,1!l'i?~i«tt:INi!;,1'!.i,o;,§11 .. ,f.'!i:N10:•:'"·'·""¥.-m~1'1•iN• ·"''•:.!'•:!., r:":·';··'r"-·"'!?''J ,, ... -:• ... '.:'"'"'i.'.'f";>;1.-·%"''""'·'•'""°'·'"''·"k: ..... ,,r .. ,,,,, '""··:'"·.!.· .. , .,• .. ,,__ ... _.p.l)!t.91)N,<N0i .. !l_l:','Jiii:!!.-L\wJ/X\f.·!Y"-"MWP.1~.:

)

Review of seeps in balancing a skelet~n redo~ reactio11. <

1. Fl.nd. the elements that undergo a change In oxidation state. I) Assign oxidation numbers.to everything. · II) See If and how the oxidation numbers change for .er;ich element.

2. . Write the oxidation and reduction half-reactions. i) Balance atoms. ii) Balance charge.

3. Multiply each half~reactlon by~;-;;:: lowest coeffjcient needed to make the# of electrons lost in the oxidation half-reaction = the # of electrons gained In the reduction half-reaction.

4. Add the two half-reactions to produce a baidi iced redox reaction.

5. Transfer the coefficients obtained in your balanced redox reaction to the appropriate terms In the original skeleton equation. Balance any remaining terms by Inspection.

)

I~~.~~:.:1;-i;>•.::•.,111 l·•.151~• • .! I,· .~fi.\!L.~.I. 1. ·~

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~DOX AND ELECTROCHEMISTRY Date

0.y_rlcjini tb.t· 0.ctivitej $t·ri:e·3

. During a single replacement reaction, one element :. >ar.e m.r <'<IE~:-:i~;: . · N:> 'f1:<i~l-:;;t ·. takes the place of another in a compound.. Manv

• • !'.<~:!1$e. .. l m f'u~, .. · J

- ::... · <.·.:. compounds, such as the copper II sulfate, consist of ·.;·.,,.·

.J;i;;i> ~i'cu. :;so.i :: two parts, arn~tal (cop~er)_and anonn:etaI (sulfa~e). ~-:KL;.:=~~,;;:;:,..,_.. When a metal such as zmc 1s dropped mto a solllti on

.:-.·;-:- .. ···,,. containing copper II sulfate, its naiural tendency is to .u;;c!:l.-:· combine 'vith the sulfate by giving electrons to it

The sulfate's. outer shell is already full, however, because it has already gained electrons from the copper. As a result, however, the copper has room for zinc's electrons. If zinc can force copper to take its electrons, zinc can become a cation and take copper's place in the compound.. Whether or not the zinc can take the copper's pla.Ce depends upon which metal has the greater tendency to lose electrons. Scientists have determined by experimentation which metals can replace each other in aqueous solution. This resulted in the development of the Activity Series as sho\11-n in 'Chart J to the right. The most active metals and nonmetals are shovm tov,;ard the top of the chart. Elements at the top of the activity series can replace tho_se below them ·

For each example below, if a reaction will occur based on the eleQ:lents' positions in the Actfrity Series, CQmplete the equation and balance it. If there is no reaction, write no reaction. [NOTE: for metals, the format for single replacement reactions is AB+ C-+ CB+ A; for nonmetals the format is · AB + D -+ AD+ B]

1. Mg(s) + HCI(aq) -+

2. Ag(s) + Cu(NOJ:i(aq)-+

3. Zn(s) + Mn(CH3COO)laq) __.

4. AJ(s) + HCI(aq)-+

5. Cu(s) + HBr(aq) __.

6. Cu(s) + AgCH3COO(aq)-+

7. Sn(s) + H:S04(aq)-+

8. Mg(s) + Pb(NOJ:(aq)-+

9. Pb(s) + AuCI(aq)-+

10. Au(s) + LiCl(aq) _.

©Evan P. Silberstein, 2003

aA

Period

TaM~J Act.wny s~:rt11s~·~·

~'.i:·

f.fost t "llkt~i* ~-unitwt~lsfMt~

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(".<,

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Al

Ti

?,fo

7..i> c·,. f;·

C<,

~'i

.Si1

Vh >"H.~

c-~i

Ag_,

:~.;.:!$ .. ~ L:;;:ol)..."1; i....----1.----.... L~:ll.st :'-"!:\<:~(;;·~ry s~~·t-:;~ h:s~:d {:fc i~~·fis;:\~~!:l: ~~::V~d::sttl

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L TH·E ELECTROCHEMICALC CELL Name ~-

v

Alo ___ _ salt bridge t----Pb

\..__..,/.

•. !a•,.

1.0M

Al(N03)3

1.0 M.

Pb{N03) 2

Answer the questions below referring to the above diagram and a Table ·Of staAdard · T Electrode Potentials.- ·

l. Which Is more easily oxidized, metal, aluminum or lead? ______________ ___

2. What Is the balanced equation showing the spontaneous reaction t~at occurs?

)( What Is the maximum voltage that the above cell can pro.duce?------

4. What Is the direction of electron flow in the wire? ---"'"""· ·----------

5. What Is the direction. of positive Ion flow In the salt bridge? ---------

-· -----6~·-Whlch·electrode is decreeslng-in-size?--- -· --- . . . -· - ·-- ____ .

7. Which electrode Is Increasing in size?------------------

8. What is happening to the concentration of aluml0um Ions? -------­

\. 9. What is happening to the concentration of lead Ions? __ ---------

. ~ ~ What Is the voltG:ige in this ceH when the reaction reaches equilibrium? ___ _

11. ~Nch~theanode?_~~-~----~-~--~-----~

12. Which is the cathode?_·--~-----------------13. What is the positive electrode? ________________ _

~1~1ti.i.s-t~~: Form WSl0.3.lA Name

REDOX P...ND ELECTROCHEMISTRY Date Period ~'tf ---

0 $alt an~ Jba:fjW.·rg

Portable electronic devices run on batteries. 'foe electricity generated by a battery comes from a chemiyal reaction known as an oxidation-reduction reaction. During an a single replacement,. a type of oxidation-reduction reaction, more active metals transfer electrons to less active metals. As a result, the more active metal is m .. -idized, and the less active metal is reduced. If the oxidation and reduction half reactions are physically separated and attached by a wire, electrons v.~ flow through the wire duri11g the reaction and can be used to power our portable electronics. This is done by putting electrolytes, usually aqueous acids, bases, or salts, into separate containers. The separate containers are called half cells because the half reactions are isolated in them. They are connected by a salt bridge whieh lets ions travel between half cells. Electrodes are immersed into the electrolytes. The electrodes are merely metals '"ith differing activity. Completing the circuit by connecting the electrodes enables electrons to flow from the more active metal to the less active metal. reducing it. 111e electrode where reduction occurs is called the cathode. The electrode where oxidation occurs is called the anode. The device that produces electric current from a chemical reaction is called a ~·oltaic cell. Several voltaic cells attached togetherfom1 a battery of cells. A battery, produces a higher voltage than a single cell.

+------\l<:~tn1~tt~.r + w"'ir~----

Salt bridg~~ ----+~

;-----Anode

fkcfH}lyle

Hodt1di<lt1 .h:M cdi O:<ldatfon half l~H

· .............. CATHODE rMnO~ and grapnit;:}

.~:::::::: ... ·.; .. ELECTROLYTE

...... ······ (f..OH{aqj)

............... SALT BRtDGE . xiro~s mat~rial)

·········-···ANODE {pow"5ere.d Zn)

Answer the questions below based on your reading above and on your know led .-~e of chemistry.

Answer questionsi-4 by referring to !he diagrcnn to the ·right showing an electrochemical cell. The metal at electrode A is silver. The metal at electrode C is lead The e/ectro~vtes at locations B, D, and E are potassium nitrate, silver nitrate, and lead nitrate respecrNely.

1. In what direction do electrons flow in the electrochemical cell pictured to the

right (A to C or C to A)?----------------

2. Wh?t type of chernjcal change is taking place in the half-cell contained in the

beaker at location Fl---------------~----'--

3. At which location are-electrons being gained? _________ _

4. Which metal is being replaced during the reaction in this electrochemical cell?

Continue~

~.-~ B t~a;1

::. . ... ..: .. . :· ~ . :.:

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1 . ·-q l 1--P ~ L- : ~-i ·• \,~ ., I t --n .:· ~ .... , _ ...... ·~ , --~ J·--r~ ...... L <-.~ ( _ ___::} < .... ___ ,)

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REDOX AND ELECTROCHEMISTRY Page"'

Answer questions 5-16 by referring to Table J. For each. of the electrode pairs, which would be the anode ;n an electrochemical cell? ·

5. Cu/Zn ........ 9 ... Au/Pb ........

6. Pb/Sn ........ 10. ~n .......

7. K/Al ......... IL Fe/Zn ........

8. Ba/Li ........ 12. Co/Ca ........

Answer questionsll-19 by referr;ng tot he setup shown to the r;ghtus;nga lemon and metal strips. It actually produces measurable electridty.

17. Explain how the lemon battery \X.'orks? ------------'

13. Co/Ni ........ -----

14. H/Ag ........ ___ _

15. Cu/Mg ........ ___ _

16. Zn/Al .. ......... - 4

_,_, .. ---·-··--------·-·····-V•:.ltmeim:

Metu1 s~nr. ... .:··

tt:!2:2!~:: . .. Lemon

·...... . .. .......

. \....__/

18. What parts of a typical voltaic cell are missing in the lemon battery? What effect does this have on how well it fimctions?

E~']Jlali1. _______________________________________ _

19. If the metal strip oD the right is iron and the metal strip OD the left is aluminum, m what direction will electriCity flow?

---c·e•--20. What happens at the anode of an electrochemical cell?---------,---------------

21. TI1ere are two voltaic cells pictured on the previous page. TI1e one on the left is called a ·wet cell, while the one at the left is called a dry cell. TI1e one at the rigl1t is also called an alkalin~ cell. What is f11e difference between these cells that

accounts for the difference in the 1x,.:ay they are named? _______________________ _

\__)

·~1 EYan P. Silberstein., 2003

~r

0ll-t.ttdtt1"':(: Form WSlO. 3. 2A Name

REDOX AND ELECTROCHEMISTRY Date

-t

F@rcin~ Ei~«;tr@.t13 t® ~©v~·

The energy to run most cars from gasoline. Except in diesel engines, the energy is released from the gasoline by e:\-ploding it with a tiny spark from a spark plug. The energy to rnake the spark comes from the car's battery. The battery in the car is called a "-wet cell." It contains sulfuric acid [H2SOiaq)], a liquid electrolyte. The electricity is generated by the =-,, fo~ow-ing chemical reaction: ._ .. ,, ..

Pb02 +Pb+ 2H:?S04 -+ 2PbS04 .+ 2H;i0 ;~ -.;;-~x····<·:-:.-....:-:.·:.:v:.-:-:-::-.·M-;;.:..:-;... ... :·:·~X«·z~:..:-:····.;:.;-:;.:.:,.:.; .... xv:...v:-:«.-.N:V;.; .. ""W"X·:.-·····:.:..::.;-:...::~ ....... :<A..::~$-t"'

Period -~~

Car batteries can last for several years. This is because they gerrecharged. As the engine spins, a moVing magnet in the alternator pushes electrons in a direction opposite to the way they nonnally flow from the battery. These electrons revei:-se the chemical reaction that generated electricity in the battery. A cell that uses electricity to produce a chemical reaction in this way is called an electrolytic cell. 'When the car battery is generating electricity it is au electrochemical cell. When it is being recharged, it is au electrolytic cell. ·

Answer the questions below based on the reading above and on your knowledge of chemistry.

1. \Vrite the chemical reaction that occurs when a car battery generates electricity.---------------

a. · Write the half reactions: -----------------------------------

b. What is o:-..-idizeci. and what is reduced? ________________________ _

2. Write the chemical reaction that occurs wl1en a car battery is rechafged. ----------------­

a. Write the half reactions:--------~-----~------------------

b. What is oxidized, and what is reduced? _________________________ _

3. Aluminum is found in the mineral bamcite (Al20,). To get pure almuinum, the aluminum needs to be separated from

o:\-ygen.

a. Imagine bauxite forms by the following reaction: 4Al + 302 -+- 2Al,03. Write the half reactions. _____ _

b. During the formation of bau.xite from its elements, what is oxidized, and what is reduced? Does this make sense

considering that alumimnn is a metal? E::"l..--plain. ------------------------

c. Write the reaction for the purification of aluminum from bauxite. ------------------

~ Continue on the next page.

.~-

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REDOX AND ELECTROCHEMISTRY Pag~

d. Write the half reactions for the purification of alwnllJ.um. During the purification,. what is oxidized., and what is

reduced?-----------------------------------~

e. ·considering that aluminum is a metal,. suggest a method to purify it. Explain. ------------

4. Iron· is often protected from rusting by a process called galvanizing. When a metal is gah•anized, it is coated with zinc. · One way to coat iron with zinc is through a single replacement reaction: Fe + Zn(N03) 2 -+ Fe(N03)3 +Zn. Since the reaction occurs at the surface of the iron. the iron becomes plated ·with zinc. ·

a. Write the h?lfreactions for this reaction. What was o:-..-idized. and what was reduced? __________ _

b. Consult the actiYity series on Chart J. How likely is this reaction to occur? E"'-J>Iain. ------------'--

c. Suggest a method to plate iron \vith zinc.-----------------------------

5. What is an electrolytic cell? What are some of its functions? ____________________ _

6. What type of cell is represented by the following reaction: Cu+ AgNO~ -+ Ag+ Cu(N03) 2?. Write the half reactions

associated ;vi th it. Identi:f'.y the oxidation and reduction half reactions. ------------------

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©Evan P. Silberstein, 2003

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