oxidising and reducing ability in group 7

8/9/2019 Oxidising and Reducing Ability in Group 7

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AS Information Sheet 2.5(2)

OXIDISING AND REDUCING ABILITYIN GROUP 7

This information sheet is intended to try to prevent misunderstandings in questions on redox

and to help you answer exam questions in the right way. Oxidation and reduction will be

discussed in terms of electrons in the following paragraphs.

SOME DEFINITIONS

♦ Oxidation Loss of electrons (oxidation number increases)

♦ Reduction Gain of electrons (oxidation number decreases)

♦ Oxidising Agent substance which can oxidise another

substance. !n doing this" the oxidising agent taes a!a"

e#ect$ons from the other substance. This means that it gainselectrons and is itse#% $educed . Oxidising agents are" therefore"

easily reduced.

♦ Oxidising Po!e$ The oxidising power (or ability)

describes how easily one substance can ta#e away electrons

from another. powerful oxidising agent is #een to gain

electrons $ it must be highly electronegative. %luorine is the

best example.

♦ Reducing Agent substance which can reduce another

substance. !n doing this" the reducing agent gi&es e#ect$ons to

the other substance. This means that it loses electrons and is

itse#% oxidised . &educing agents are" therefore" easily oxidised.

♦ Reducing Po!e$ The reducing power (or ability) describes

how easily one substance can give electrons to another.

powerful reducing agent is #een to donate electrons. l#ali

metals are good examples (caesium is the best).

TYPICAL EXAM !ESTIONS

'uestions which as# you to compare oxidising abilitypower or reducing abilitypower areoften in the context of trends in Group (the *alogens).

It i" im#ortant to $e %er& 'ear a$ot the *ifferen'e

$et+een ,-aoen"/ an* ,-ai*e ion"/.

♦ 'a#ogens These are the e#e(ents" fluorine" chlorine"

bromine and iodine. They consist of diatomic molecules (F2"

C2" 0r2 and I2). +otice that the names end in )INE .

♦ 'a#ide Ions These are the negati&e ions (,- oxidation state)

fluoride (F1)" chloride (C1)" bromide (0r1) and iodide (I1)

-




formed when the halogen atoms gain an electron each. The

ions have noble gas electron configurations. +otice that their

names end in )IDE .




TENDS 3IT-IN 4O!P

Ato(ic * Ionic Radii

The radii of both the atoms and the ions inc$ease down the Group from % to !. This is

because there is an additional electron shell added from one period to the next. /achsuccessive shell has a larger average radius than the previous one. &emember that the

ionic radius for a negative ion (li#e the halide ions) is larger than the corresponding atom

because the nucleus has to hold in an extra electron.

♦ E#ect$onegati&it"

/lectronegativity is a measure of how much an atom attracts electrons to itself in a

(covalent) bond. The electronegativity of the elements dec$eases down the Group.

%luorine has the highest value and iodine has the lowest. lthough the atomic number

increases (and therefore the actual positive charge on the nucleus) down the Group" the

valence electrons are further away from the nucleus. There are also more inner shells toshield the nucleus from the valence electrons. This means that the outer electrons

experience a smaller force of attraction to the nucleus as we go down the Group. %luorine

atoms are small with a high effective nuclear charge and can" therefore" attract bonding

electrons easily. !odine has large atoms with a much smaller effective nuclear charge"

ma#ing it less good at attracting other electrons.

Oxidising Po!e$ o% t+e E#e(ents

The ability of the elements to oxidise other substances dec$eases down the Group. This is

due to a combination of the increasing atomic radius and decreasing electronegativity

described above. %luorine0s small si1e and high electronegativity ma#e it a powerfuloxidising agent. !odine has large atoms and a low electronegativity" so is a much less

powerful oxidising agent.

Reducing A,i#it" o% t+e 'a#ide Ions

This is concerned with the ability of the negative ions to give away their extra electron

and act as reducing agents" as shown in the equations below2,

2F1 F2 6 2e1 2I1 I2 6 2e1

%er& *iffi't 7ite ea"&

This inc$eases down the Group (it is the opposite of the oxidising power of the element).

/xam questions relating to redox reactions in Group are typically of two types2

-. *alogen displacement reactions.

/. The reaction of concentrated sulfuric acid with solid sodium or potassium halides (e.g."

sodium chloride or potassium iodide).

These are discussed in more detail on the next two pages.

3




-ALO4EN DISPLACEMENT EACTIONS

This is where a more reactive halogen element displaces a less reactive one from its

compound in solution. The halogen in the compound is in oxidation state $- (halide ion).

This is oxidised to the element (oxidation state 4). t the same time" the incoming halogen

element is reduced from oxidation state 4 to $-. The example shown below is for chlorine

displacing bromine from a solution of potassium bromide2,

5l (aq) 6 78r (aq) → 75l (aq) 6 8r (aq)

5l (aq) 6 8r , (aq) → 5l, (aq) 6 8r (aq)

0 1- 1- 0

It is i(2o$tant to $ead t+e !a" t+e 3uestion is 2+$ased4 !t often says 9explain the reactions

in terms of the oxidising ability of the halogens:. !n this case" the chlorine oxidises the

bromide because it is a better oxidising agent than bromine. /xplain in terms of atomic

radius and electronegativity discussed on the previous page.

&emember that the halide ions (which have a noble gas electron configuration) are colourless" but the elements (diatomic molecules) have characteristic colours.

'a#ogenCo#ou$ o% 5a2ou$ o$ in

Ine$t So#&ent 6e4g4 +exane.Co#ou$ o% A3ueous So#ution

chlorine green,yellow very pale yellow,green

bromine red,brownyellow to orange (depending on

concentration)

iodine purple brown

(blue,blac# on adding starch)

EACTIONS OF SOLID -ALIDES 3IT- CONC. S!LP-!IC ACID

5oncentrated sulfuric acid behaves differently to the dilute acid. s well as behaving li#e a

normal acid (donating *6 ions) and it can act as an oxidising agent. (!t can also act as a

dehydrating agent" but this feature is not involved in these reactions.) ;ou must learn the

following observations2,

-ALIDE O0SE8ATIONS EXPLANATIONS

+a5l

or

75l

<isty fumes of strongly acidic *5l gas

are formed.

This is an acid,base reaction.

5l is still in oxidation state $-"

so no redox reactions areinvolved.

+a8r

or

78r

<isty fumes of strongly acidic *8r gas

are formed. &ed,brown bromine vapour

and colourless =O gas (pungent and

acidic fumes) are also formed" but =O is

disguised by the presence of *8r.

8r $ ions have been oxidised

to 8r . They have reduced the

= in *=O> from oxidation

state 6? to 6> in =O gas.

+a!

or

7!

little *! gas and dar# brown solid

iodine are formed (gives purple vapour if

warmed). There is a strong smell of bad

eggs from *= gas. @ale yellow solidsulphur is also formed and some =O gas.

! $ ions have been oxidised to

!. They have reduced the = in

*=O> from oxidation state 6?

to 6> in =O gas" 4 in solidsulphur and $ in *=.

>




The equations for the reactions are as follows2,

- +a5l(s) 6 *=O>(l) → +a*=O>(aq) 6 *5l(g)

(*=O> acts as an acid and donates an *6 ion to the 5l $ ion to form *5l gas)

. +a8r (s) 6 *=O>(l) → +a*=O>(aq) 6 *8r (g)

8r , 6 *6 6 *=O> → 8r (g) 6 =O(g) 6 *O(l)

8. +a!(s) 6 *=O>(l) → +a*=O>(aq) 6 *!(g)

! , 6 *6 6 *=O> → !(s) 6 =O(g) 6 *O(l)

?! , 6 ?*6 6 *=O> → 3!(s) 6 =(s) 6 >*O(l)

A! , 6 A*6 6 *=O> → >!(s) 6 *=(g) 6 >*O(l)

It is i(2o$tant to $ead t+e 2+$asing o% 3uestions on t+is4 They are often of the form

“explain the observations in terms of the reducing ability of the halide ions:.

5oncentrated *=O> can act as an oxidising agent and is sufficiently powerful to be able to

oxidise iodide to iodine easily and bromide to bromine partially" but not chloride or fluoride

(for these *5l or *% gas is formed" which is not oxidation" but an acid,base reaction" as

explained above).

nother way of loo#ing at this is to thin# in terms of the +a#ide ions ,eing a,#e to $educe t+e

su#%u$ic acid . % , is least willing to give away its extra electron" so its reducing ability is very

poor. 5l $ similarly does not readily part with its extra electron. 8r $ can give its extra

electron away relatively easily and the S in the acid is reduced from oxidation state 69 to 6:in SO2. ! $ very readily gives away its extra electron to something else" so its reducing ability

is good. !n these reactions" the S in the acid is reduced from oxidation state 69 to a variety of

lower states" e.g." 6: in SO2" ; in solid S and <2 in -2S. The reducing ability of the halide

ions increases down the group.

The explanation is that the larger ions with the valence electrons far from the nucleus and

shielded by the inner shells can lose the extra electron more easily. =o" the halogen element

which is the most powerful oxidising agent (fluorine) has an ion (fluoride) which is the least

powerful reducing agent.

B

oxidising and reducing ability in group 7

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