Uk Chemistry in Perspective for Bored and Confused Senior Schoo

8/8/2019 Uk Chemistry in Perspective for Bored and Confused Senior Schoo

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hembook.co.uk: CHEMISTRY IN PERSPECTIVE FOR BORED AND CONFUSED SENIOR SCHOOL STUDENTS#chap15

5.1. INTRODUCTION

5.1.1. The order of these next three chapters may seem unusual: groups I and II, then groupII, then group IV. It is quite intentional. The chemistry of groups I and II is predominated by t

ow effective nuclear charge of the elements they contain. The chemistry of group VII isredominated by the high effective nuclear charge of the elements it contains. It is useful to loot these two extremes before considering group IV as an example of a group where the effectiv

uclear charge is intermediate, and the chemistry changes down the group from being more likeroup VII to being more like groups I and II.

5.1.2. The elements in groups I and II are all very similar. Their low effective nuclear chargemeans that their chemistry is dominated by the loss of electrons to form positive ions (section.4.2.). Group I elements form 1+ ions and group II elements form 2+ ions (section 4.7.4.).

here are some differences between groups I and II, mainly due to the higher effective nuclearharge in group II. There are also changes in behaviour down both groups due to the increase i

umber of shells. Since the dominating effect of low effective nuclear charge is less marked inroup II, the changes are more noticeable down group II.

he changes down groups I and II, and the differences between them, form the basis of thishapter, but do not get them out of proportion. The overwhelming characteristic of these elemes their similarity to each other. They are all electropositive metals which nearly all form basicxides and hydroxides.

5.2. THE ELEMENTS

5.2.1. Atomic and ionic radius increase down both groups as can be predicted from thencreasing number of shells. The higher effective nuclear charge in group II enables us to predichat the radii of both atoms and ions will be lower in group II.

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n addition, the ions are smaller than the parent atoms, not only because they have one less sheontaining electrons, but also because the ratio of protons to electrons is higher and thus the

ffective nuclear charge is higher in the ions. Note that this increase in effective nuclear charge reater when group II elements form ions.

5.2.2. Metallic bond strength decreases down both groups because the nuclei become furtherrom the metallic bonding orbital of electrons. In group II the higher effective nuclear charge, anhe greater number of electrons in the metallic bonding orbital, enables us to predict that the

metallic bond strength is generally greater than in group I.

Group I elements are soft silvery metals which become softer as the group is descended, reflect

he decreasing bond strength. Group II elements are harder, grey metals which also becomeofter as the group is descended.

he weaker the bonding, the less vigorously the ions in the metal must vibrate before the bondsreak. Thus lower temperatures are needed to melt and boil the metals as the groups areescended. In addition, melting and boiling points are generally higher in group II. (What is thexplanation?)

Other functions which follow the changes and differences in metallic bond strength are: enthalp



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f fusion, enthalpy of vapourisation, and atomisation enthalpy. These all decrease down theroups, and are generally greater in group II.

Remember that atomisation energy is the energy change during the process:

M(s)..........g ..........M(g)...............( metal lattice........g ........free gaseous atoms )

n contrast, electrical and thermal conductivity do not follow the pattern that might be predictedrom decreasing bond strength. You might predict that it would be easier to transfer electronshrough a lattice in which they are are not tightly held. In fact it is more difficult, a bit like tryingush a car made of candyfloss. Thus conductivity decreases down both groups and is generallyreater in group II.

5.2.3. Ionisation becomes even more able to occur down both groups as the outer electronsecome further from the attractive force of the nucleus. The higher effective nuclear charge

makes ionisation generally less likely in group II, though it is hardly relevant in practice.



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oth these facts are reflected in the values of the first ionisation energies which decrease downoth groups, and which are generally higher in group II: the more stongly an electron is attractey the nucleus, the greater the energy needed to remove it.

Remember that first ionisation energy is the energy change involved in the process:

M(g)..........g ..........M+(g)..........+..........e-

5.2.4. Flame tests are also a characteristic of the ease with which electrons are promoted toigher energy levels. The heat of a flame is enough to raise electrons, even in the ions, to highe

evels. When they return to lower levels, electromagnetic radiation is emitted (section 2.2.).

n the case of group I and II elements, some of these emissions are in the visible light range:

ABLE 15.1. Flame test colours.

Elem ent Flam e ColourElectron Transition(interest only)



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Lithium red 2p t 2s

Sodiumyellow(very strong)

3p t 3s

Potassium lilac 4d t 4p

Calcium brick red 5s t 4p

5.2.5. The above changes (sections 15.2.1. to 15.2.3.) do not occur smoothly down the groupshe differences are due to the different screening effects of different types of orbital (section3.2.).

5.2.6. The strength of hydration of the ions decreases down both groups. This is predictableecause the increase in size of the ions down the groups means that their surface charge densitecreases, and the - oxygen atoms of water are therfore less strongly attracted.

n group II, the generally smaller size, together with the double positive charge, enables us toredict that the water molecules will be generally attracted more strongly. In some casesydrolysis occurs (section 14.4.2.ii.).

hese facts are reflected in the hydration energies which decrease down the groups, and areenerally higher in group II. Remember, hydration energy is the energy change during theollowing process:

Eg: Group I: ..........M+(g)..........g ..........M+(aq)



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5.2.7. Standard electrode potentials numerically increase (become more negative) down bothroups. The values in group II are generally lower (less negative). Remember, electrode potentre quoted for reduction reactions i.e.

M+(aq)..........+..........e-..........g ..........M(s)

M2+(aq)..........+..........2e-..........g ..........M(s)

ABLE 15.2. Standard electrode potentials for elements in groups I and II.

Group Ielements

Group IIelements

Element S.E.P. Element S.E.P.

Li -3.05V Be -1.85V

Na -2.71V Mg -2.38V

K -2.93V Ca -2.87V



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Rb -2.95V Sr -2.89V

Cs -2.92V Ba -2.90V

he increasing negative values mean the the reverse process (oxidation) becomes even morekely to occur as the groups are descended. Thus the metals become even better reducing agens the groups are descended. The generally less negative values in group II mean that thexidation is less likely in group II and the elements are slightly poorer reducing agents.

ote that the redox potential for H2O(l) + e-g OH-(aq) + H2(g) is 0.41V, so it can be predict

hat all these elements reduce water to hydrogen.

ontinuing the emphasis on the reverse process, it can be seen that the change from metal toqueous metal ions can be broken down into three simpler processes: atomisation, ionisation, aydration (though it should not be imagined that these three processes occur one after the othe

Down the groups, it is the increasing ability of the elements to be atomised (section 15.2.2.) andheir increasing ability to undergo ionisation (15.2.3.) which account for the increasing liklihood he change from metal to aqueous metal ions. These two increases counteract the decreasingtrength with which the ions are hydrated (section 15.2.6.). This also gives a more detailedccount of the numerical increase in S.E.P. values.

imilarly, the lower ability of group II metals to undergo atomisation and ionisation (sections5.2.2. and 15.2.3.) outweigh the greater strength of their hydration (section 15.2.6.). Thisccounts for the lower liklihood of their change from metal to aqueous metal ions, and their

umerically lower redox potentials.

he idiosyncratic S.E.P value for lithium is a charateristic of the particularly small size of its ionsnd the consequent strength with which water molecules are attracted to them. This outweighshe greater attraction of lithium's nuclei for their outer electrons i.e. the lower ability of the metao undergo atomisation/ionisation.

eryllium ions are even smaller and, moreover, doubly charged. However, this does not outweighe strength of beryllium's metallic bonding and its relative inability to ionise, a process which, it

hould be remembered, involves removal of two electrons in group II.

5.2.8. Reactivity

Reactivity with water generally follows the pattern that might be predicted from the order of tedox potentials shown above i.e. reactivity increases down both groups and is generally lower iroup II. Indeed the series in which elements are placed in order of redox potentialelectrochemical series) is often known as the reactivity series.



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owever, some of the rather clumsy language in section 15.2.7. was used in order to avoid usinhe concept of reactivity. Redox potentials give a measure of how likely a reaction is, that isnformation about equilibrium, or thermodynamic information (sections 11.13. and 12.5.3.).

n contrast, reactivity is a term most often used in connection with observable rate of reaction,inetic information. Thermodynamic reactivity is more closely reflected in the amount of heatenerated in a reaction, though in table 15.3., even this is confused with the rate at which heat

enerated.

n support of this distinction, lithium is much less reactive than the other elements in group Iespite the extremely high value of its redox potential.

ABLE 15.3. Reactivity of group I and II elements with water

Group I elements Group II elements

Li: slow reaction with cold water Be: slow reaction with steam

Na: molten bead skims on surface,ignition possible

Mg: very slow reaction with cold water,steady reaction with steam

K: ignition, possible explosionCa: steady reaction with cold water,fast reaction in hot water

Rb: explosive Sr: reacts readily with cold water

Cs: explosive Ba: vigorous reaction with cold water

he reason for lithium's low reactivity is the strength of its metallic lattice (section 15.2.2.).ithium atoms/ions must be pulled away from the lattice for reaction to occur. This, the rate-etermining step, is a difficult and slow process. In energy terms, the activation energy is high.

) Reactivity with air increases down the groups, and is generally lower in group II. They allarnish in air, though beryllium and magnesium form only a surface coating.

ll the elements burn in air to form oxides. There are three possible oxides, the normal oxide, th

eroxide, and the superoxide. Not all the elements form all three oxides on combustion in andequate supply of air, indeed some of the oxides do not even exist for certain of the elements.he explanation is contained within section 15.3.3.

ABLE 15.4. Oxides of group I and II elements.

Oxide ionElements which form it oncombustion in adequate supply of air

Elements which form it at all



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Normal O2-Gp I: Li (&Na)GpII: All

Gp I: AllGp II: All

Peroxide O22- Gp I: Na, (K, Rb & Cs)

Gp II: (Sr) & BaGp I: AllGp II: All except Be

Superoxide O2- Gp I: (Na), K, Rb, Cs

Gp II: NoneGp I: All except LiGp II: All except Be & Mg

) Reactivity with nitrogen decreases down the groups and is generally greater in group II. Thiseflected by the fact that, when they burn in air, group II elements form some nitride, theroportion of which decreases down the group. Moreover, lithium is the only group I element toombine directly with nitrogen.

he reason that group II elements combine directly with nitrogen is that the small size and doubharge of their ions enables them to form a very strong lattice with nitride ions. This is strong

nough to outweigh the strong triple bond which needs to be broken in nitrogen, as well as themetallic bonding which is broken in the elements themselves.

hese predictions are supported by a study of the energy changes involved.

itride ions, like hydride ions (section 14.3.2.i.) and oxide ions (sections 14.5.2.i. and 15.3.5.)ave such a high surface charge density that they are hydrolysed in water, producing an alkalineolution:

5.3. THE COMPOUNDS

5.3.1. Most of the compounds are ionic, though the small size of lithium and beryllium ions givehe compounds of these two elements a high degree of covalent character (section 4.6.2.i.).

5.3.2. The strength of the lattices decreases down both groups, and is generally greater in groI. This is because (as a rough initial guide) the larger the metal ion, the less tightly it can pack



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with the anions. The forces of attraction are inversely proportional to (r+ + r-)2, where r+ and rre the radii of the positive and negative ions respectively. Ions of similar size form particularlytrong lattices.

lso, the more highly charged the metal ion, the more strongly it attracts the anions i.e. the forcf attraction are proportional to the charge difference between the ions, z.

he changes in lattice strength are reflected by the overall decrease in lattice enthalpies down troups, and the generally higher lattice enthalpies in group II. Remember that lattice enthalpy ishe enthalpy change for the process:

MX(s)..........g ..........M+(g)..........+..........X-(g)

hough often it is quoted for the reverse process.

ote that lattice enthalpy will be proportional to z/(r+ + r-)

5.3.3. Solubility of compounds shows somewhat complex changes down the group.

he overall process for MX(s) is represented by:

MX(s)..........g ..........M+(aq)..........+..........X-(aq)

n other words, breaking of the lattice and hydration of the ions.

Down both groups:

The lattice becomes weaker (section 15.3.2.) tending to favour solubility;

) Atraction for water molecules (hydration) becomes weaker (section 15.2.6.), tending toiscourage solubility.

he key to the balance between these counteracting effects is the size of the anion:

ince the strength of the lattice is inversely proportional to (r+ + r-)2

, then:

for small anions, the decreasing strength of the lattice will be considerable and outweigh theecreasing strength of hydration, thus solubility will increase (e.g. fluorides and hydroxides).

for large anions, the decreasing strength of the lattice will be insignificant compared with theecreasing strength of hydration, and solubility will decrease (e.g. sulphates(VI) and carbonates

for intermediate sized anions, the solubility at first decreases, and then increases (e.g. chloride



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nd nitrates(V)).

hese group changes are summarised in FIG. 15.5.

or a given anion, the group II compounds are generally less soluble because the effect of the

ouble charge on the strength of the lattice is greater than the effect on hydration.

Moreover, doubly charged cations (group II) form particularly strong lattices with doubly charge

nions (e.g. CO32-) so these are generally less soluble than the compounds of group II elements

with singly charged anions. This, together with the similarity in size of barium and sulphate ionssection 15.3.2.), accounts for the extreme insolubility of barium sulphate.

attice enthalpy and hydration enthalpy data can be used to check predictions. Note also that inorder line cases, entropy changes must be considered.



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5.3.4. The thermal stability of compounds with large anions increases down both groups, and ienerally lower in group II.

hermal instability of the compounds with large anions is a function of the polarising power of tation. This in turn is greatest for small ions (section 15.2.1.), and highly charged ions i.e. ions

with a high surface charge density. (Hence the group changes and differences described in the l

aragraph.)

hus ions with a high surface charge density tend to polarise large anions, pulling electrons intoonds with themselves, and breaking off a smaller anion which forms a stronger lattice.

.g. polar ising e ffect



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hus the carbonates, hydroxides, and nitrates(III) of group II elements and of lithium in group end to decompose in the same way on heating. Thermal stability of the group II carbonates,

ydroxides, and nitrates(V), increases down the group. Beryllium sulphate(VI) is the only sulphaVI) to decompose at comparable temperatures.

he carbonates and hydroxides of group I elements other than lithium are relatively stable to hend their nitrates(V) decompose according to the fifth equation in the above list.

he polarising power of cations with a high surface charge density also accounts for the increastability down the groups of oxides with large anions, as well as the difficulty in forming group Iuperoxides at all (section 15.2.8.ii. and table 15.4.).

5.3.5. Acid/base properties of the oxides and hydroxides: With the exception of beryllium, all throup I and group II elements form basic oxides and hydroxides (sections 14.5.2.i. and 14.5.3.ieryllium forms amphoteric oxides and hydroxides (14.5.2.iv. and 14.5.3.iii.).

n general, the oxides react vigorously with water, though BeO is insoluble (section 14.5.2.iv.) amagnesium hardly reacts. There is, then, increasing reactivity down the groups and lowereactivity in group II, as may be predicted. The reaction is between the oxide ions and water, anhe vigour depends on the ease with which they are broken away from the lattice (section 15.3.



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ompare this with the reactions of the hydrides (section 14.3.2.i.) and nitrides (section 15.2.8.iin water.

) All the group II hydroxides except LiOH are very strong bases and very soluble in water. Inwater they clearly produce hydroxide ions i.e. an alkaline solution. Solubility increases down theroup (section 15.3.3.) and thus, so too, does maximum possible concentration of hydroxide ionand max. alkalinity).

he group II hydroxides are less soluble (15.3.3.), though solubility again increases down the

roup and thus, so too, does the maximum possible concentration of hydroxide ions (and max.lkalinity). Beryllium hydroxide is weakly basic and insoluble in water.

he thermal stability of hydroxides has already been discussed in section 15.3.4.

5.3.6. The important properties of other compounds like hydrides and chlorides can be predicterom the relevant parts of sections 14.3. and 14.4. together with an understanding of the changown and between groups I and II discussed in this chapter.

5.4. QUESTIONS

) Draw tables summarising the properties of the group I and II elements and their compoundsnd the changes down both groups plus the differences between the groups. You may need a biece of paper. E.g.

PROPERTY GROUP I group differences GROUP II



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Elements Li..................Cs Be..................Ba

E.g. meltingpoint

decreases down group greater in group II decreases down group

Compounds ..... ..... .....

..... ..... ..... .....

) Which will involve the greater contraction in size, the formation of ions by group I elements, he formation of ions by group II elements?

) Explain the relationship between changes in metallic bonding down groups I and II, and thenthalpy changes described in section 15.2.2.

) Describe the relationship between the increasing number of shells down groups I and II, andhe numerical increase in standard electrode potentials down the groups. Also account for thenomalous S.E.P. of lithium in terms of its size, and for the difference between groups I and II ierms of the higher effective nuclear charge of group II elements.

) Predict the differences in standard electrode potentials described in question 4, usingredictions about atomisation, ionisation, and hydration enthalpies.

) Head elements (those at the top of groups) are often said to show anomalous behaviour. Listnd explain any examples of anomalous behaviour shown by lithium and beryllium and theirompounds.

) Lithium and magnesium show similarities which are often said to be an example of a diagonaelationship. List any examples of such a diagonal relationship, and describe the basis of the effe

) If the redox potentials of group I and II elements are not true indicators of their reactivity wiwater, how should this really be predicted? Apply your method.

) Write balanced equations for the reactions which occur when each of the following elements eated in an adequate supply of air: lithium, potassium, beryllium and calcium.

0) Are sections 15.2.8. ii and iii talking about kinetic, or thermodynamic reactivity? Explain.

1) How can lattice enthalpy and hydration enthalpy data be used to check solubility predictionssection 15.3.3.)?

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Uk Chemistry in Perspective for Bored and Confused Senior Schoo