Properties of Period 3 oxides

8/11/2019 Properties of Period 3 oxides

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PHYSICAL PROPERTIES OF THE PERIOD 3

OXIDES

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These pages explain the relationship between the physical properties of the oxidesof Period elements (sodium to chlorine) and their structures! "rgon is obviously

omitted because it doesn#t form an oxide!

A quick summary of the trends

The oxides

The oxides we#ll be loo$ing at are%

Na&' g' "l&' i'& P*'+ ' l&'.

P*'/ '& l&'

Those oxides in the top row are $nown as the highest oxidesof the various

elements! These are the oxides where the Period elements are in their highestoxidation states! 0n these oxides, all the outer electrons in the Period elementare being involved in the bonding - from 1ust the one with sodium, to all seven ofchlorine#s outer electrons!

The structures

The trend in structure is from the metallic oxides containing giant structures ofions on the left of the period via a giant covalent oxide (silicon dioxide) in the

middle to molecular oxides on the right!

Melting and boiling points

The giant structures (the metal oxides and silicon dioxide) will have high meltingand boiling points because a lot of energy is needed to brea$ the strong bonds (ionic

or covalent) operating in three dimensions!

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The oxides of phosphorus, sulphur and chlorine consist of individual molecules -some small and simple2 others polymeric!

The attractive forces between these molecules will be van der 3aals dispersion anddipole-dipole interactions! These vary in si4e depending on the si4e, shape and

polarity of the various molecules - but will always be much wea$er than the ionic orcovalent bonds you need to brea$ in a giant structure!

These oxides tend to be gases, li5uids or low melting point solids!

Electrical conductivity

None of these oxides has any free or mobile electrons! That means that none ofthem will conduct electricity when they are solid!

The ionic oxides can, however, undergo electrolysiswhen they are molten! They canconduct electricity because of the movement of the ions towards the electrodes

and the discharge of the ions when they get there!

The metallic oxides

The structures

odium, magnesium and aluminum oxides consist ofgiant structures containingmetal ions and oxide ions!


There are strong attractions between the ions in each of these oxides and theseattractions need a lot of heat energy to brea$! These oxides therefore have high

melting and boiling points.


None of these conducts electricity in the solid state, but electrolysisis possible if

they are molten! They conduct electricity because of the movement and dischargeof the ions present! The only important example of this is in the electrolysis ofaluminum oxide in the manufacture of aluminum!

3hether you can electroly4e molten sodium oxide depends, of course, on whether itactually melts instead of subliming or decomposing under ordinary circumstances! 0fit sublimes, you won#t get any li5uid to electroly4e6

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agnesium and aluminum oxides have melting points far too high to be able toelectroly4e them in a simple lab!

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Silicon dioxide (silicon(I! oxide!

The structure

The electronegativity of the elements increases as you go across the period, and bythe time you get to silicon, there isn#t enough electronegativity difference betweenthe silicon and the oxygen to form an ionic bond! ilicon dioxide is agiant covalentstructure"

There are three different crystal forms of silicon dioxide! The easiest one to

remember and draw is based on the diamond structure!


ilicon dioxide has a high melting point - varying depending on what the particularstructure is (remember that the structure given is only one of three possiblestructures), but they are all around +.7! 8ery strong silicon-oxygen covalentbonds have to be bro$en throughout the structure before melting occurs! ilicon

dioxide boils at &&7!

9ecause you are tal$ing about a different form of bonding, it doesn#t ma$e sense totry to compare these values directly with the metallic oxides! 3hat you can safelysay is that because the metallic oxides and silicon dioxide have giant structures,the melting and boiling points are all high!


ilicon dioxide doesn#t have any mobile electrons or ions - so it doesn#t conductelectricity either as a solid or a li5uid!

The molecular oxides

Phosphorus, sulphur and chlorine all form oxides which consist of molecules! omeof these molecules are fairly simple - others are polymeric! 3e are 1ust going to

loo$ at some of the simple ones!

elting and boiling points of these oxides will be much lower than those of the

metal oxides or silicon dioxide! The intermolecular forces holding one molecule toits neighbors will be van der 3aals dispersion forces or dipole-dipole interactions!The strength of these will vary depending on the si4e of the molecules!

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None of these oxides conducts electricity either as solids or as li5uids! None ofthem contains ions or free electrons!

The phosphorus oxides

Phosphorus has two common oxides, phosphorus (000) oxide, P*'/, and phosphorus(8) oxide, P*'+!

#hosphorus (III! oxide (tetraphosphorus hexoxide!

Phosphorus (000) oxide is a white solid, melting at &*7 and boiling at +.7!

The phosphorus is using only three of its outer electrons (the unpaired pelectrons) to form bonds with the oxygens!

#hosphorus (! oxide (tetraphosphorus decoxide!

Phosphorus (8) oxide is also a white solid, subliming (turning straight from solid tovapour) at 7! 0n this case, the phosphorus uses all five of its outer electrons inthe bonding!

olid phosphorus(8) oxide exists in several different forms - some of them

polymeric! 3e are going to concentrate on a simple molecular form, and this is alsopresent in the vapour!

This is most easily drawn starting from P*'/! The other four oxygens are attachedto the four phosphorus atoms via double bonds!

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The sulphur oxides

ulphur has two common oxides, sulphur dioxide (sulphur (08) oxide), '&, and

sulphur trioxide (sulphur (80) oxide), '!

Sulphur dioxide

ulphur dioxide is a colourless gas at room temperature with an easily recogni4edcho$ing smell! 0t consists of simple '&molecules!

The sulphur uses * of its outer electrons to form the double bonds with theoxygen, leaving the other two as a lone pair on the sulphur! The bent shape of '&is due to this lone pair!

Sulphur trioxide

Pure sulphur trioxide is a white solid with a low melting and boiling point! 0t reactsvery rapidly with water vapour in the air to form sulphuric acid! That means that ifyou ma$e some in the lab, you tend to see it as a white sludge which fumesdramatically in moist air (forming a fog of sulphuric acid droplets)!

:aseous sulphur trioxide consists of simple 'molecules in which all six of thesulphur#s outer electrons are involved in the bonding!

There are various forms of solid sulphur trioxide! The simplest one is a trimer,

';, where three 'molecules are 1oined up and arranged in a ring!

The fact that the simple molecules 1oin up in this way to ma$e bigger structures iswhat ma$es the sulphur trioxide a solid rather than a gas!

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The chlorine oxides

hlorine forms several oxides!


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PROPERTIES OF THE PERIOD 3 CHLORIDES

This page loo$s at the structures of the chlorides of the Period elements (sodium

to sulphur>), their physical properties and their reactions with water!

hlorine and argon are omitted - chlorine because it is meaningless to tal$ about?chlorine chloride?, and argon because it doesn#t form a chloride!

A quick summary of the trends

The chlorides

The chlorides we#ll be loo$ing at are%

Nal gl& "ll il* Pl@ &l&

Pl

"s you will see later, aluminum chloride exists in some circumstances as a dimer,

"l&l/!

The structures

odium chloride and magnesium chloride are ionic and consist of giant ionic latticesat room temperature

"luminum chloride and phosphorus (8) chloride are tric$y6 They change theirstructure from ionic to covalent when the solid turns to a li5uid or vapour! There is

much more about this later on this page!

The others are simple covalent molecules!


odium and magnesium chlorides are solids with high melting and boiling points

because of the large amount of heat (energy) which is needed to brea$ the strongionic attractions!

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The rest are li5uids or low melting point solids! Aeaving aside the aluminum chlorideand phosphorus (8) chloride cases where the situation is 5uite complicated, the

attractions in the others will be much wea$er intermolecular forces such as van der3aals dispersion forces! These vary depending on the si4e and shape of the

molecule, but will always be far wea$er than ionic bonds!


odium and magnesium chlorides are ionic and so will undergo electrolysiswhen theyare molten! Blectricity is carried by the movement of the ions and their discharge

at the electrodes!

0n the aluminum chloride and phosphorus (8) chloride cases, the solid doesn#tconduct electricity because the ions aren#t free to move! 0n the li5uid (where itexists - both of these sublime at ordinary pressures), they have converted into a

covalent form, and so don#t conduct either!

The rest of the chlorides don#t conduct electricity either solid or molten because

they don#t have any ions or any mobile electrons!

%eactions &ith &ater

"s an approximation, the simple ionic chlorides (sodium and magnesium chloride)1ust dissolve in water!

The other chlorides all react with water in a variety of ways described below foreach individual chloride! The reaction with water is $nown as hydrolysis!

The Individual $hlorides

Sodium chloride' a$l

odium chloride is a simple ionic compound consisting of a giant array of sodium and

chloride ions!

" small representative bit of a sodium chloride lattice loo$s li$e this%

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This is normally drawn in an exploded form as%

The strong attractions between the positive and negative ions need a lot of heatenergy to brea$, and so sodium chloride has high melting and boiling points!

0t doesn#t conduct electricity in the solid state because it hasn#t any mobile

electrons and the ions aren#t free to move!


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0ons of this sort are acidic - the degree of acidity depending on how much theelectrons in the water molecules are pulled towards the metal at the centre of the

ion! The hydrogens are made more positive than they would otherwise be, and moreeasily pulled off by a base!

0n the magnesium case, the amount of distortion is 5uite small, and only a smallproportion of the hydrogen atoms are removed by a base - in this case, by watermolecules in the solution!

The presence of the hydronium ions in the solution causes it to be acidic! The factthat there aren#t many of them formed (the position of e5uilibrium lies well to theleft), means that the solution is only wea$ly acidic!

Gou may also find the last e5uation in a simplified form%


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The coordination of the aluminum changes at these temperatures! 0t becomes *-coordinated - each aluminum now being surrounded by * chlorines rather than /!

3hat happens is that the original lattice has converted into "l&l/molecules!

This conversion means, of course, that you have completely lost any ionic character

- which is why the aluminum chloride vapori4es or melts (depending on thepressure)!

There is an e5uilibrium between these dimers and simple "llmolecules! "s the

temperature increases further, the position of e5uilibrium shifts more and more tothe right!

)ummary

"t room temperature, solid aluminum chloride has an ionic lattice with a lotof covalent character!

"t temperatures around +H - +;7 (depending on the pressure), aluminumchloride coverts to a molecular form, "l&l/! This causes it to melt orvapori4e because there are now only comparatively wea$ intermolecularattractions!

"s the temperature increases a bit more, it increasingly brea$s up intosimple "llmolecules!

olid aluminum chloride doesn#t conduct electricity at room temperature becausethe ions aren#t free to move! olten aluminum chloride (only possible at increasedpressures) doesn#t conduct electricity because there aren#t any ions any more!

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The aluminum chloride reacts with the water rather than 1ust dissolving in it! 0n thefirst instance, hexaa5uaaluminum ions are formed together with chloride ions!

Gou will see that this is very similar to the magnesium chloride e5uation given above- the only real difference is the charge on the ion!

That extra charge pulls electrons from the water molecules 5uite strongly towards

the aluminum! That ma$es the hydrogens more positive and so easier to removefrom the ion! 0n other words, this ion is much more acidic than in the corresponding

magnesium case!

These e5uilibria (whichever you choose to write) lie further to the right, and so thesolution formed is more acidic - there are more hydroxonium ions in it!

or, more simply%

Silicon tetrachloride' Si$l*

ilicon tetrachloride is a simple no-messing-about covalent chloride! There isn#t

enough electronegativity difference between the silicon and the chlorine for thetwo to form ionic bonds!

ilicon tetrachloride is a colourless li5uid at room temperature which fumes inmoist air! The only attractions between the molecules are van der 3aals dispersion

forces!

0t doesn#t conduct electricity because of the lac$ of ions or mobile electrons!

0t fumes in moist air because it reacts with water in the air to produce hydrogenchloride! 0f you add water to silicon tetrachloride, there is a violent reaction toproduce silicon dioxide and fumes of hydrogen chloride! 0n a large excess of water,the hydrogen chloride will, of course, dissolve to give a strongly acidic solution

containing hydrochloric acid!

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The phosphorus chlorides

There are two phosphorus chlorides = phosphorus (000) chloride, Pl, andphosphorus (8) chloride, Pl@!

#hosphorus (III! chloride (phosphorus trichloride!' #$l)

This is another simple covalent chloride - again a fuming li5uid at room temperature!

0t is a li5uid because there are only van der 3aals dispersion forces and dipole-dipole attractions between the molecules!

0t doesn#t conduct electricity because of the lac$ of ions or mobile electrons!

Phosphorus (000) chloride reacts violently with water! Gou get phosphorous acid,


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Phosphorus (8) chloride has a violent reaction with water producing fumes ofhydrogen chloride! "s with the other covalent chlorides, if there is enough water

present, these will dissolve to give a solution containing hydrochloric acid!

The overall e5uation in boiling water is%

,-isulphur dichloride' S.$l.,

Jisulphur dichloride is a simple covalent li5uid - orange and smelly6

The shape is surprisingly difficult to draw convincingly6 The atoms are all 1oined upin a line - but twisted%

The reason for drawing the shape is to give a hint about what sort ofintermolecular attractions are possible! There is no plane of symmetry in the

molecule and that means that it will have an overall permanent dipole!

The li5uid will have van der 3aals dispersion forces and dipole-dipole attractions!

There are no ions in disulphur dichloride and no mobile electrons - so it neverconducts electricity!

Jisulphur dichloride reacts slowly with water to produce a complex mixture ofthings including hydrochloric acid, sulphur, hydrogen sulphide and various sulphur-

containing acids and anions (negative ions)! There is no way that you can write a

single e5uation for this - and one would never be expected in an exam!

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Summary

+! #hysical properties of oxides/ The physical properties of these oxidesdepend on the type of bonding!

Na&', "l&'and g' are ionic oxides and hence have a high melting point! g' and

"l&'have a higher melting point than Na&' since the charges are higher, resultingin a stronger attraction between the ions!

i'&has a giant covalent structure and hence a high melting point! There are strongcovalent bonds between all the atoms and thus lots of energy is re5uired to brea$them!

P*'+and 'are molecular covalent and so only intermolecular forces exist

between the molecules! The melting points are thus much lower! P*'+is a muchbigger molecule than 'and so has a much higher melting point, as the van der3aals forces are stronger!

Blement Na g "l i P

Kormulae ofoxide

Na&' g' "l&' i'& P*'+ '

tructure ofoxide

0onic 0onic ostlyionic

:iantcovalent

olecularcovalent

olecularcovalent

elting point ofoxide L7

+&.@ &H@& &.& +. -+

&! Acid0base character of oxides

0onic oxides contain the '&-ion! This is a strongly basic ion which reacts with water

to produce hydroxide ions%

'&-(a5) D


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'(s) D


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'&(g) D


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The acid-base properties of the oxides of Period can be summari4ed in thefollowing table%

Blement Na g "l i P

Kormulae ofoxides

Na&' g' "l&' i'& P*'+

'&

'

"cid-basecharacter ofoxide

9asic 9asic "mphoteric "cidic "cidic "cidic

p< of solution

when dissolvedin water

+& - +* H - ;.

insoluble

.

insoluble

& - *& - *('&)

+ - (')

The oxides therefore become more acidic on moving from left to right in theperiodic table!

#hysical properties of chlorides

a$l and Mg$l.are ionic chlorides! ince a large amount of energy is re5uired to

separate the ions, the melting point is high!

Al$l)and Si$l*are molecular covalent chlorides, and so only intermolecular forces

exist between the molecules! The melting points are thus much lower than the ionicchlorides!

"llactually exists in polymeric form in the solid state, which is converted to adimeric form in the gas phase! "t high temperatures, it reverts to a simplemolecular structure%

Al

Cl

Cl

Cl

Al

Cl

Cl

Cln

Al Al

Cl Cl

Cl

Cl

ClCl

n/2 n

complex polymer dimer monomer

(solid) (gas - at low T) (gas - at high T)

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The aluminum atom is electron deficient = it has only of its four valence orbitals

occupied, so it has an empty orbital with which it can accept a lone pair of electronsfrom a l atom on an ad1acent monomer!

#$l+is ionicso its melting point is thus high! 'n heating, however, it reverts to asimple covalent structure and sublimes!

Blement Na g "l i P

Kormula ofchloride

Nal gl& "ll il* Pl@

tructure of

chloride

ionic ionic polymer molecular

covalent

0onic

elting point of

chloride L7

H+ .+ +H* @H +/&

%eaction of chlorides &ith &ater

The way in which chlorides react with water depends on the type of bonding

present in the chloride%

Ionic chloridesdissolve in water to give neutral solutions%Nal(s)NaD(a5) D l-(a5) p< C .

gl&(s)g&D(a5) D &l-(a5) p< C .

Aluminium chloridereacts with water to give hydrated aluminum ions and chlorideions! The hydrated aluminium ions undergo deprotonation to give an acidic solution%"ll(s) D /


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The water molecules attac$ the covalent chlorides by donating lone pairs ofelectrons into empty low-lying orbitals on the electropositive atoms! 0n the case of

"ll, there is an available p orbital, and in il*and Pl@there are available d-orbitals%

s pd

"ll ENeF

il* ENeF

Pl@ ENeF

0t the availability of these low-lying empty orbitals which enables these chloridesto react readily with water!

%eaction of $l.&ith &ater

l&(g) D

Properties of Period 3 oxides

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