IB QUESTIONS - ACIDS AND BASES

1. Which statement about hydrochloric acid is false?

A. It can react with copper to give hydrogen

B. It can react with sodium carbonate to give carbondioxide

C. It can react with ammonia to give ammonium chloride

D. It can react with copper oxide to give water

2. 1.00 ern! of a solution has a pH of 3. 100 crn ' of the samesolution will have pH of:

A. 1 C. 5

B. 3 D. Impossible to calculate from the data given.

3. Which statement(s) is/are true about separate solutions of astrong acid and a weak acid both with the sameconcentration?

I. They both have the same pH.

II. They both have the same electrical conductivity.

A. I and II B. lonly C. II only D. Neither I nor II

4. Identify the correct statement about 25 ern! of a solution of0.1 mol drrr-' ethanoic acid CH3COOH.

A. It will contain more hydrogen ions than 25 crn ' of0.1 mol drn=' hydochloric acid.

B. It will have a pH greater than 7.C. It will react exactly with 25 crn ' of 0.1 mol drrr-'

sodium hydroxide.

D. It is completely dissociated into ethanoate andhydrogen ions in solution.

5. Which species cannot act as a Lewis acid?

A. NH3 B. BF3 C. Fe2+

6. NH3(aq), HCI(aq), NaOH(aq), CH3COOH(aq)

When 1.0 mol drrr-' solutions of the substances above arearranged in order of decreasing pH the order is:

A. NaOH(aq), NH3(aq), CH3COOH(aq), HCI(aq)

B. NH3(aq), NaOH(aq), HCI(aq), CH3COOH(aq)

C. CH3COOH(aq)' HCI(aq)' NaOH(aq), NH3(aq)

D. HCI(aq), CH3COOH(aq), NH3(aq), NaOH(aq)

7. A solution with a pH of 8.5 would be described as:

A. very basic. C. slightly acidic.

B. slightly basic. D. very acidic.

8. Which statement is true about two solutions one with a pHof 3 and the other with a pH of 6?

A. The solution with a pH of 3 is twice as acidic as thesolution with a pH of 6

B. The solution with a pH of 6 is twice as acidic as thesolution with a pH of 3

C. The hydrogen ion concentration in the solution with apH of 6 is one thousand times greater than that in thesolution with a pH of 3

D. The hydrogen ion concentration in the solution with apH of 3 is one thousand times greater than that in thesolution with a pH of 6

9. Which of the following is not a conjugate acid-base pair?

A. HNO/N03-

B. H2SO/HS04-

C. NH/NH2-

D. H30+/OH-

10.During the titration of a known volume of a strong acidwith a strong base:

A. there is a steady increase in pH

B. there is a sharp increase in pH around the end point

C. there is a steady decrease in pH

D. there is a sharp decrease in pH around the end point.

11. Three acids, HA, HB, and HC have the following Ka values

Ka(HA) = 1 x 10-5 Ka(HB) = 2 x 10-5 Ka(HC) = 1 x 10-6

What is the correct order of increasing acid strength(weakest first)?

A. HA, HB, HC

B. HC HB, HA

C. HC HA, HB

D.HB,HA,HC

12.Which of the following reagents could not be addedtogether to make a buffer solution?

A. NaOH(aq) and CH3COOH(aq)

B. NaCH3COO(aq) and CH3COOH(aq)

C. NaOH(aq) and NaCH3COO(aq)

D. NH4CI(aq) and NH3(aq)

13.When 1.0 cm ' of a weak acid solution is added to100 crn ' of a buffer solution:

A. the volume of the resulting mixture will be 100 crn '

B. there will be almost no change in the pH of the solution

C. the pH of the solution will increase noticeably

D. the pH of the solution will decrease noticeably.

14.What is the pH of a buffer solution in which theconcentration of the acid HX and the salt NaX are both0.1 mol dm-3 (Ka = 1 x 10-5)?

A. 3 B.4 C.5 0.6

15. Which salt does not form an acidic solution in water?

A. MgCI2

16. An indicator changes colour in the pH range 8.3-10.0.This indicator should be used when titrating a knownvolume of:

A. a strong acid with a weak base

B. a weak acid with a weak base

C. a weak base with a strong acid

D. a weak acid with a strong base.

53 18 questions - Acids and bases

M e 9W eEp

Redox reactions (1)DEFINITIONS OF OXIDATION AND REDUCTIONOxid~tion used to be narrowly defined as the addition of oxygen to a substance. For example, when magnesium is 'burned'i!lair--the magnesium is oxidized to magnesium oxide.

2Mg(s) + 02(g) -> 2MgO(s)

The electronic configuration of magnesium is 2.8.2. During the oxidation process it loses two electrons to form the Mg2+ ionwith the electronic configuration of 2.8. Oxidation is now defined as the loss of one or more electrons from a substance. This is amuch broader definition, as it does not necessarily involve oxygen. Bromide ions, for example, are oxidized by chlorine to formbromine.

2B,(aq) + CI2(aq) -> Br2(aq) + 2CI-(aq)

If a substance loses electrons then something else must be gaining electrons. The gain of one or more electrons is called reduction.In the first example oxygen is reduced as it is gaining two electrons from magnesium to form the oxide ion 02-. Similarly, in thesecond example chlorine is reduced as each chlorine atom gains one electron from a bromide ion to form a chloride ion.

Since the processes involve the transfer of electrons oxidation and reduction must occur simultaneously. Such reactions areknown as redox reactions. In order to distinguish between the two processes half-equations are often used:

2Mg(s) -> 2Mg2+(S)+ 4e-

02(g) + 4e- -> 202-(s)

2Mg(s) + 02(g) -> 2MgO(s)

-OXIDATION -

- REDUCTION -

OVERALL REDOX EQUATION

2B,(aq) -> Br2(aq) + 2e-

CI2(aq) + 2e- -> 2CI-(aq)

Understanding that magnesium must lose electrons and oxygen must gain electrons when magnesium oxide MgO is formed fromits elements is a good way to remember the definitions of oxidation and reduction. Some students prefer to use the mnemonicOILRIG: Oxidation Is the Loss of electrons, Reduction Is the Gain of electrons.

RULES FOR DETERMINING OXIDATIONNUMBERSIt is not always easy to see how electrons have beentransferred in redox processes. Oxidation numbers can be auseful tool to identify which species have been oxidized andwhich reduced. Oxidation numbers are assigned according toa set of rules:

1. In an ionic compound between two elements the oxidationnumber of each element is equal to the charge carried bythe ion, e.g.

Na+CI-(Na=+1;CI=-1)

Ca2+CI-2(Ca = +2; CI = -1)

4. The algebraic sum of all the oxidation numbers in an ion =the charge on the ion, e.g.

SOl- [(+6) + 4 x (-2) = -2];Mn04- [(+7) + 4 x (-2) = -1]; NH/ [(-3) + 4 x (+1) = +1]

5. Elements not combined with other elements have anoxidation number of zero, e.g. 02; p4; 58'

6. Oxygen when combined always has an oxidation numberof -2 except in peroxides (e.g. H202) when it is -1.

7. Hydrogen when combined always has an oxidationnumber of +1 except in certain metal hydrides (e.g. NaH)when it is -1.

Many elements can show different oxidation numbers indifferent compounds, e.g. nitrogen in:

NH3(-3)

NO(+2)

2. For covalent compounds assume the compound is ionicwith the more electronegative element forming thenegative ion, e.g.

CCI4(C = +4; CI =-1)

NH3(N = -3; H= +1)

3. The algebraic sum of all the oxidation numbers in acompound = zero, e.g.

CCI4 [(+4) + 4 x (-1) = 0];H2S04 [2 x (+1) + (+6) + 4 x (-2) = 0]

54 Oxidation and reduction

When elements show more than one oxidation state theoxidation number is represented by using Roman numeralswhen naming the compound,

e.g. FeCI2 iron(11) chloride; FeCI3 iron(lll) chlorideK2Cr207 potassium dichromate(VI); KMn04 potassiummanganate(VII)Cu20 copper(l) oxide; CuO copper(ll) oxide.

Redox reactions (2)OXIDATION AND REDUCTION IN TERMS OF OXIDATION NUMBERSWhen an element is oxidized its oxidation number increases,

e.g. Mg(s) ~ Mg2+(aq) + 2e-

(0) (+2)

When an element is reduced its oxidation number decreases,e.g. 80~-(aq) + 2H+(aq) + 2e- ~ 80§-(aq) + H20(I)

(+6) (+4)

The change in the oxidation number will be equal to the number of electrons involved in the half-equation.

Using oxidation numbers makes it easy to identify whether or not a reaction is a redox reaction.

Redox reactions (change in oxidation numbers)CuO(s) + H2(g) ~ Cu(s) + H20(I)(+2) (0) (0) (+1)

SFe2+(aq) + Mn04-(aq) + BH+(aq) ~ Fe3+(aq) + Mn2+(aq) + 4H20(I)(+2) (+7) (+3) (+2)

Not redox reactions (no change in oxidation numbers)precipitation Ag+(aq) + crtaq) --...AgCI(s)

(+1) (-1) (+1) (-1)

neutralization HCI(aq) + NaOH(aq) -- NaCI(aq) + HP(I)(+1)(-1) (+1)(-2)(+1) (+1)(-1) (+1)(-2)

Note: reactions where an element is uncombined on one side of the equation and combined on the other side must be redoxreactions since there must be a change in oxidation number,

e.g. Mg(s) + 2HCI(aq) -- MgCI2(aq) + H2(g)

OXIDIZING AGENTS AND REDUCING AGENTSA substance that readily oxidizes other substances is known as an oxidizing agent. Oxidizing agents are thus substancesthatreadily accept electrons. Usually they contain elements that are in their highest oxidation state,

e.g. 02' C12,F2, 803 (80/- in solution), Mn04-, and Cr20/-.

Reducing agents readily donate electrons and include H2, Na, C, CO, and 802 (80l- in solution),

e.g. Cr20/-(aq) + 38032-(aq) + BH+(aq) -- 2Cr3+(aq) + 380/-(aq) + 4H20(I)

(+6) (+4) (+3) (+6)(orange) (green)

(oxidizing agent) (reducing agent)

CI2(aq)(0)

(oxidizing agent)

+ 2Br-(aq)(-1)

(reducing agent)

2CI-(aq) + Br2(aq)(-1) (0)

Oxidation and reduction 55

Reactivity seriesREACTIVITYLithium, sodium, and potassium all react with cold water to give similar products but the reactivity increases down the group.

2M(s) + 2H20(I) --...2M+(aq) + 20H-(aq) + H2(g) (M = Li, Na, or K)

Slightly less reactive metals react with steam and will give hydrogen with dilute acids, e.g.

Mg(s) + 2H20(g) --...Mg(OH)2 + H2(g)

Mg(s) + 2HCI(aq) --...Mg2+(aq) + 2CI-(aq) + H2(g)

In all of these reactions the metal is losing electrons - that is, it is being oxidized and in the process it is acting as a reducingagent. A reactivity series of reducing agents can be deduced by considering the reactivity of metals with water and acids, andthe reactions of metals with the ions of other metals.

Reactivity series of reducing agents

K(s) ¢ e- + K+(aq)Na(s) ¢ e- + Na+(aq)Li(s) ¢ e- + Li+(aq)Ca(s) ¢ 2e- + Ca2+(aq)Mg(s) ¢ 2e- + Mg2+(aq)AI(s) ¢ 3e- + AI3+(aq)Zn(s) ¢ 2e- + Zn2+(aq)Fe(s) ¢ 2e- + Fe2+(aq)Pb(s) ¢ 2e- + Pb2+(aq)~H2(g) ¢ e- + W(aq)Cu(s) ¢ 2e- + Cu2+(aq)Ag(s) ¢ e- + Ag+(aq)

Increasingreactivity

The series can be extended for oxidizingagents. The most reactive oxidizing agentwill be the species that gains electrons themost readily. For example, in group 7

I-(aq) = e+ ~ 12(aq)

Bi(aq) = e- + ~Br2(aq)

Crtaq) = e- + ~CI2(aq)

F-(aq) = e- + ~F2(aq)

The more readily the metal loses its outer electrons the more reactive it is. Metalshigher in the series can displace metal ions lower in the series from solution, e.g.zinc can react with copper ions to form zinc ions and precipitate copper metal.

Zn(s) + Cu2+(aq) --...Zn2+(aq) + Cu(s)

increasingoxidizingability

Zn(s) --...Zn2+(aq) + 2e-Cu2+(aq) + 2e- --...Cu(s)

Zn loses electrons in preference to CuCu2+ gains electrons in preference to Zn2+

Oxidizing agents lower in the series gainelectrons from species higher in the

This also explains why only metals above hydrogen can react with acids (displace series, e.g.

hydrogen ions) to produce hydrogen gas, e.g. CI2(aq) + 2Bi(aq) --...2CI-(aq) + Br2(aq)

Zn(s) + 2W(aq) --...Zn2+(aq) + H2(g)

SIMPLE VOLTAIC CELLSA half-cell is simply a metal in contact with an aqueous solution of its own ions. A voltaic cell consists of two different half-cells,connected together to enable the electrons transferred during the redox reaction to produce energy in the form of electricity. Thecells are connected by an external wire and by a salt bridge, which allows the free movement of ions.

A good example of a voltaic cell is a zinc half-cell connected to a copper half-cell. Because zinc is higher in the reactivity seriesthe electrons will flow from the zinc half-cell towards the copper half-cell. To complete the circuit and to keep the half-cellselectrically neutral, ions will flow through the salt bridge. The voltage produced by a voltaic cell depends on the relativedifference between the two metals in the reactivity series. Thus the voltage from a Mg(s)/Mg2+(aq) half-cell connected to aCu(s)/Cu2+(aq) half-cell will be greater than that obtained from a Zn(s)/Zn2+(aq) half-cell connected to a Fe(s)/Fe2+(aq) half-cell.

voltmeter

negative electrode ~-~----iH

oxidation occurs here

positive electrode(+)

reduction occurs heresaltbridge

(saturated solution of KN03)

CllM Zn

voltaic cellZn'"(aq)

half-cell this half-cell produces electrons this half-cell receives electrons

Zn(s)-Zn'"(aq) + ze- Cu'"(aq) + ze- ~ Cuts)

56 Oxidation and reduction

Electrolysis (1)ELECTROLYTIC CELLIn a voltaic cell electricity is produced by the spontaneous redox reaction taking place. Electrolytic cells are used to makenon-spontaneous redox reactions occur by providing energy in the form of electricity from an external source. In an electrolyticcell electricity is passed through an electrolyte and electrical energy is converted into chemical energy. An electrolyte is asubstance which does not conduct electricity when solid, but does conduct electricity when molten or in aqueous solution and ischemically decomposed in the process. A simple example of an electrolytic cell is the electrolysis of molten sodium chloride.

During the electrolysis: • sodium is formed at the negative electrode (cathode)Na+(I) + e- -;. Na(l) reduction

• chlorine is formed at the positive electrode (anode)2CI-(I) --> CI2(g) + 2e- oxidation

• The current is due to the movement of electrons inthe external circuit and the movement of ions in theelectrolyte.

Electrolysis is an important industrial process used to obtain reactive metals, such assodium, from their common ores.

battery ammeter tomeasure current

electrolysis ofmolten sodiumchloride

(+i electrodeoxidation

occurs here2CC-Q+2e

(-) electrodereductionocrurs hereNa++e--Na

FACTORS AFFECTING THE DISCHARGE OF IONS DURING ELECTROL YSISDuring the electrolysis of molten salts there are only usually two ions present, so the cation will be discharged at the negativeelectrode (cathode) and the anion at the positive electrode (anode). However, for aqueous electrolytes there will also behydrogen ions and hydroxide ions from the water present. There are three main factors which influence which ions will bedischarged at their respective electrodes.

1. Position in theelectrochemical series

The lower the metal ionis in the electrochemicalseries the more readily itwill gain electrons (bereduced) to form themetal at the cathode.Thus in the electrolysisof a solution of sodiumhydroxide, hydrogenwill be evolved at thenegative electrode inpreference to sodium,whereas in a solution ofcopper(lI) sulfate, copperwill be deposited at thenegative electrode inpreference to hydrogen.

For negative ions theorder of discharge followsOH- > CI- > 50/-.

3. The nature of the electrode

It is normally safe to assume that theelectrode is inert, i.e. does not play anypart in the reaction. However, if copperelectrodes are used during the electrolysisof a solution of copper sulfate then thepositive electrode is itself oxidized torelease electrons and form copper(lI) ions.Since copper is simultaneously depositedat the negative electrode theconcentration of the solution will remainconstant throughout the electrolysis.

2. Concentration

If one of the ions is much more concentratedthan another ion then it may bepreferentially discharged. For example, whenelectricity is passed through an aqueoussolution of sodium chloride both oxygen andchlorine are evolved at the positiveelectrode. For dilute solutions mainly oxygenis evolved, but for concentrated solutions ofsodium chloride more chlorine than oxygenis evolved.

inert electrodes(Ptor graphite)

copper electrodes

(+i (-) (+i H

oxygen evolved2H,!O - 4Ht + 4e-

- +°2

copperdepositedCu'" + 2e- -- Cu

copper dissolvesCu - Cu2+ + 2e-

copper depositedCu2t + 2e- -v Cu

solution loses blue colourand forms sulfuric acid

solution retains blue colour

ELECTROPLATINGElectrolysis can also be used in industry to coat one metal with a thin layer of another metal. This process is known as'electroplating,

For example, in copper plating the negative electrode (cathode) is made from the metal to be copper plated. It is placed into asolution of copper(lI) sulfate together with a positive electrode (anode) made from a piece of copper. As electricity is passedthrough the solution the copper anode dissolves in the solution to form Cu2+(aq) ions and the Cu2+(aq) ions in solution aredeposited onto the cathode. By making the anode of impure copper and the cathode from a small piece of pure copper thisprocess can also be used to purify impure copper. This is an important industrial process as one of the main uses of copper isfor electrical wiring, where purity is important since impure copper has a much higher electrical resistance.

Oxidation and reduction 57

£m Electrolysis (2) and standard electrode potentials

FACTORS AFFECTING THE QUANTITY OF PRODUCTS DISCHARGED DURING ELECTROL YSISThe amount of substance deposited will depend on:

2. The charge on the ion. To form one mole of sodium in the electrolysis of moltensodium chloride requires one mole of electrons to flow through the cell. However, theformation of one mole of lead during the electrolysis of molten lead(11) bromide requirestwo moles of electrons.

1. The number of electronsflowing through the system,i.e. the charge passed. This inturn depends on the currentand the time for which itflows. If the current isdoubled then twice as manyelectrons pass through thesystem and twice as muchproduct will be formed.Similarly if the time isdoubled twice as manyelectrons will pass throughthe system and twice as muchproduct will be formed.

charge = current x time(1 coulomb = 1 ampere x1 second)

Na+(I) + e--..Na(l)Pb2+(I) + 2e- ~ Pb(l)

If cells are connected in series ..

o~ ~oo~ ~o000 0000 0

molten lead(lI) chloride dilute sulphuric acid

molar ratios of products evolved 2C12: 2Pb : °2: 2H2

If cells are connected in series then the same amount of electricity will pass throughboth cells and the relative amounts of products obtained can be determined.

STANDARD ELECTRODE POTENTIALSThere are two opposing tendencies in a half-cell. The metal maydissolve in the solution of its own ions to leave the metal with anegative potential compared with the solution, or the metal ionsmay deposit on the metal, which will give the metal a positivepotential compared with the solution. It is impossible tomeasure this potential, as any attempt to do so interferes withthe system being investigated. However, the electrode potentialof one half-cell can be compared against another half-cell. Thehydrogen half-cell is normally used as the standard. Understandard conditions of 1 atm pressure, 298 K, and 1.0 mol dm'

hydrogen ion concentration the standard electrode potential of the hydrogen electrode is assigned a value of zero volts.

H i+)

MM

Hi+)

MC'(aq)MC'iaq)

When the half-cell contains a metal above hydrogen in the reactivity series electrons flow from the half-cell to the hydrogenelectrode, and the electrode potential is given a negative value. If the half-cell contains a metal below hydrogen in the reactivityseries electrons flow from the hydrogen electrode to the half-cell, and the electrode potential has a positive value. The standardelectrode potentials are arranged in increasing order to form the electrochemical series.

298 K298 K G:Q]1],-----~---+~I+---~------

W§] e~

.~<; ©

(+) H zinc

1------00

I-- f.---- t---0 00 00 00 00 00 0

1.00 mol drrr" H''{aq) 1.00 mol drrr" Zn2+(aq)

(+)(-) copper

1.00 mol dm-3 Cu2+(aq)1.00 mol dm-3 H+(aq)

The platinum electrode is coated with finely -divided platinum, which serves as a catalystfor the electrode reaction.

pressure Just above 1 atm so H2ca n escape from electrode

58 Oxidation and reduction

Electrochemical series

ELECTROCHEMICAL SERIES(more complete series can be found inthe IB data booklet)

Couple E" I VK(s)!K+(aq) -2.92Ca(s)!Ca2+(aq) -2.87Na(s)!Na+(aq) -2.71Mg(s)/Mg2+(aq) -2.36AI(s)1 AI3+(aq) -1.66Zn(s)/Zn2+(aq) -0.76Fe(s)/Fe2+(aq) -0.44~H2(g)/W(aq) 0.00Cu(s)/Cu2+(aq) +0.34naq)/~I2(aq) +0.62Ag(s)1 Ag+(aq) +0.80Br-(aq)/~Br2(aq) +1.09

C1-(aq)/~12(aq) + 1.36F-(aq)/~F 2(aq) +2.87

SHORTHAND NOTA TlON FOR A CELLTo save drawing out the whole cell a shorthandnotation has been adopted. A half-cell is denotedby a I between the metal and its ions, and two -0. 76-+-...::Z::.;..n::..:/Z=n.:...2_+_vertical lines are used to denote the salt bridgebetween the two half cells,

Cu(s)/Cu2+(aq) II W(aq)/H2(g) andZn(s)/Zn2+(aq) II W(aq)/H2(g)

The standard electromotive force (emf) of anycell E,':II is simply the difference between thestandard electrode potentials of the twohalf-cells,

e.g.

E" IV 0

Cu/Cu2++0.34

E"+ O.34V

e.g. Cu(s)/Cu2+(aq) II Zn(s)/Zn2+(aq)

-0.76 V

ELECTRON FLOW AND SPONTANEOUSREACTIONSBy using standard electrode potentials it is easy todetermine what will happen when two half-cells areconnected together. The electrons will always flow from themore negative half-cell to the more positive half-cell, e.g.consider an iron half-cell connected to a magnesium half-cell: e

Fe(S)/Fe2+(aWg2+(aq)/M9(S)EB -0.44 V -2.36 V

more positiveFe2+ + 2e- --;. Fe

more negativeMg --;. Mg2+ + 2e-

Spontaneous reaction: Mg(s) + Fe2+(aq) --;.Mg2+(aq) + Fe(s) E~ell = 1.92 V

"Positive E;ell values give negative tJ.C values as the reactioncan provide electrical energy, i.e. do work, and the reactionis spontaneous. The reverse reaction (Fe(s) + Mg2+(aq) --;.Fe2+(aq) + Mg(s)) has a negative E;ell value which gives apositive value for tJ.C"and the reaction will not bespontaneous, so can only proceed if an external voltagegreater than 1.92 V is applied to force the reaction in theopposite direction. Note: just because a reaction has apositive E;ell value does not mean that it will necessarilyproceed, as there may be a large activation energy whichneeds to be overcome.

E~ell = 1.10V

REDOX EQUATIONSStandard electrode potentials can be extended to cover anyhalf-equation. The values can then be used to determinewhether a particular redox reaction is spontaneous,e.g. Cr20/-(aq)/Cr3+(aq), E~ + 1.33 V, but this only takesplace in acid solution, so hydrogen ions and water arerequired to balance the half-equation. Consider the reactionbetween dichromate(VI) ions and sulfite ions 50/-. Usingstandard electrode potentials the cell becomes:

e-~

Cr2072-(aq)/Cr3+ (aq) II 8032-(aq)/80/-(aq)

+1.33V +O.17VEB

more positiveCr 20/- will gain electrons

Cr20/- + 14H+ + 6e---;.2Cr3++ 7Hp

more negative8°32- will lose electrons

80l-+ H2°--;.80/- + 2H+ + 2e-

1.10V

Since the number of electrons transferred must be the samefor both equations the overall equation is obtained bymultiplying the sulfite equation by three, adding the twoequations together, and simplifying the water and hydrogenion components which appear on both sides.

Cr20/-(aq) + 8W(aq) + 38032-(aq) --;.

2Cr3+(aq) + 380/-(aq) + 4HP(I) E~ell = 1.16 V

CONVENTION FOR WRITING CELLSBy convention in a cell diagram the half-cell undergoing oxidation is placed on the left of the diagram and the half-cellundergoing reduction on the right of the diagram. The two aqueous solutions are then placed either side of the salt bridgee.g.

Zn(s)/Zn2+(aq) II Cu2+(aq)/Cu(s)

Some text books then state that:

E~ell = E~ght hand side - E~ft hand side

This can cause much confusion. Essentially a cell consists of two half-cells connected by a salt bridge and it makes nodifference which is placed on the left or on the right. Provided you remember that the electron flow is always from the half-cellwith the more negative electrode potential to the half-cell with the more positive electrode potential the convention can besafely ignored.~ •• .m"na""""RRmD __~~ •• .unm __ •• am__"~"mB~" __~ __••••~ ••••nm .~'mm ,

Oxidation and reduction 59

IB QUESTIONS - OXIDATION AND REDUCTION

1. 5Fe2+(aq) + Mn04-(aq) + 8H+(aq) ~ 5Fe3+(aq) + Mn2+(aq)+ 4H20(I)

In the equation above:

A. Fe2+(aq) is the oxidizing agent

B. H+ (aq) ions are reduced

C. Fe2+(aq) ions are oxidized

D. Mn04-(aq) is the reducing agent

2. The oxidation numbers of nitrogen in NHy HN03, andN02 are, respectively

A. -3, -5, +4

B.+3, +5,+4

C. -3, +5,-4

D.-3, +5,+4

3. Which one of the following reactions is not a redoxreaction?

A. Ag+(aq) + Cltaqj -- AgCI(s)

B. 2Na(s) + Cl2(g) ~ 2NaCI(s)

C. Mg(s) + 2HCI(aq) ~ MgCI2(aq) + H2(g)

D. Cu2+(aq) + Zn(s) -- Curs) + Zn2+(aq)

4. Which substance does not have the correct formula?

A. iron(1I1) sulfate Fe2(S04)3

B. iron(11) oxide Fe20

C. copper(l) sulfate CU2S04

D. copper(ll) nitrate Cu(N03)2

5. For which conversion is an oxidising agent required?

A. 2H+(aq) ~ H2(g)

B. 2B,(aq) ~ Br2(aq)

C. S03(g) ~ SOl-(aq)

D. Mn02(s) -- Mn2+(aq)

6. Mn(s) + Hg2+(aq) ~ Mn2+(aq) + Hg(l)

Ni2+(aq) + Mn(s) -- Ni(s) + Mn2+(aq)

Hg2+(aq) + Ni(s) -- Hg(l) + Ni2+(aq)

From the above reactions the increasing reactivity of thethree metals is:

A. Hg, Ni, Mn

B. Mn, Hg, Ni

C. Ni, Mn, Hg

D. Mn, Ni, Hg

7. When an Fe(s)!Fe2+(aq) half-cell is connected to aCu(s)!Cu2+(aq) half-cell by a salt bridge and a currentallowed to flow between them

A. the electrons will flow from the copper to the iron.

B. the salt bridge allows the flow of ions to complete thecircuit.

C. the salt bridge allows the flow of electrons to completethe circuit.

D. the salt bridge can be made of copper or iron.

8. During the electrolysis of molten sodium chloride usingplatinum electrodes

A. sodium is formed at the negative electrode

B. chlorine is formed at the negative electrode

C. sodium is formed at the positive electrode

D. oxygen is formed at the postitive electrode

9. Which statement is true?

A. Lead chloride is ionic so solid lead chloride willconduct electricity.

B. When a molten ionic compound conducts electricityfree electrons pass through the liquid.

C. When liquid mercury conducts electricity mercury ionsmove towards the negative electrode.

D. During the electrolysis of a molten salt reduction willalways occur at the negative electrode.

10.The following information is given about reactions involvingthe metals X, Y and Z and solutions of their sulfates.

X(s) / YS04(aq) ~ no reaction

Z(s) + YS04(aq) -- Y(s) + ZS04(aq)

When the metals are listed in decreasing order of reactivity(most reactive first), what is the correct order?

A. Z> Y > X

B. X> Y > Z

C. Y>X>Z

D. Y> Z > X

11. Ethanol can be oxidised to ethanal by an acidic solution ofdichromate(VI) ions.

_C2HsOH(aq) + _H+(aq) + _Cr20/-(aq) --_CH3CHO(aq) + _Cr3+(aq) + _H20(I)

The sum of all the co-efficients in the balanced equation is:

A.24 C. 28 D.30B.26

12. Which statement(s) is/are true about the standard hydrogenelectrode?

I. The electrode potential is assigned a value of 0.00 volts.

II. The hydrogen ion concentration is 0.100 mol drrr '.

III. The temperature is 273 K.

A. I, II and III

B. I and II only

C. I only

D. II and III only

60 IB questions - Oxidation and reduction

Use the following information to answer questions 13 and 14.

Sn2+(aq) + 2e- ¢ Sn(s) E" = -0.14 V

Sn4+(aq) + 2e- ¢ Sn2+(aq) EB = +0.15 V

Fe2+(aq) + 2e- ¢ Fe(s) EB = -0.44 V

E" = +0.77 V

13. Under standard conditions which statement is correct?

A. Sn2+(aq) can reduce Fe3+(aq).

B. Fe(s) can oxidise Sn2+(aq).

C. Sn(s) can reduce Fe(s).

D. Fe3+(aq) can reduce Sn4+(aq).

14. When a half-cell of Fe2+(aq)!Fe3+(aq) is connected by a saltbridge to a half-cell of Sn2+(aq)/Sn4+(aq) under standardconditions and a current allowed to flow in an externalcircuit the total e.rn.f. of the spontaneous reaction will be:

A. +0.92 V C. +0.62 V D. -0.62 VB. -0.92 V