PHYSICAL PROPERTIES OF THE PERIOD 3 OXIDES

PHYSI CALPROPERTI ESOF OFTHEPERI OD3 OXI DES These pages explain the relationship between the physical properties of the oxides of Period 3 elements (sodium to chlorine) and their structures. Argon 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!

"a#$

%g$

Al#$3

&i$#

P$

&$3

*l#$+

P$,

&$#

*l#$

Those oxides in the top row are nown as the t he highest oxides  of  of the various elements. These are the oxides where the Period 3 elements are in their highest oxidation states. -n these oxides all the outer electrons in the Period 3 element are being involved in the bonding / from 0ust the one with sodium to all seven of chlorine's outer electrons. The structures

The trend in structure is from the metallic oxides containing giant structures of ions 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 melting and boiling points because a lot of energy is needed to brea the strong bonds (ionic or covalent) operating in three dimensions. dimensions. The oxides of phosphorus sulphur and chlorine consist of individual molecules / some small and simple1 others polymeric.

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The attractive forces between these molecules will be van der 2aals dispersion and dipole/dipole interactions. These vary in sie depending on the sie shape and polarity of the various molecules / but will always be much weaer than the ionic or covalent bonds you need to brea in a giant structure. These oxides tend to be gases li4uids or low melting point solids. Electrical conductivity

"one of these oxides has any free or mobile electrons. That means that none of them will conduct electricity when they are solid. The ionic oxides can however undergo electrolysis  when  when they are molten. They can conduct 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 of giant structures containing metal ions and oxide ions . Melting and boiling points

There are strong attractions between the ions in each of these oxides and these attractions need a lot of heat energy to brea. These oxides therefore have high melting and boiling points. Electrical conductivity

"one of these conducts electricity in the solid state but electrolysis  is  is possible if they are molten. They conduct electricity because of the movement and discharge of the ions present. The only important example of this is in the electrolysis of aluminum oxide in the manufacture of aluminum. 2hether you can electrolye molten sodium oxide depends of course on whether it actually melts instead of subliming or decomposing decomposing under ordinary circumstances. -f it sublimes you won't get any li4uid to electrolye5 %agnesium and aluminum oxides have melting points far too high to be able to electrolye 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 by the time you get to silicon there isn't enough electronegativity difference between the silicon and the oxygen to form an ionic bond. &ilicon dioxide is a giant covalent structure" There are three different crystal forms of silicon dioxide. The easiest one to remember and draw is based on the diamond structure. Melting and boiling points

&ilicon dioxide has a high melting point / varying depending on what the particular structure is (remember that the structure given is only one of three possible structures) but they are all around +6*. 7ery strong silicon/oxygen covalent bonds have to be broen throughout the structure before melting occurs. &ilicon dioxide boils at ##36*. 8ecause you are taling about a different form of bonding it doesn't mae sense to try to compare these values directly with the metallic oxides. 2hat you can safely say is that because the metallic oxides and silicon dioxide have giant structures the melting and boiling points are all high. Electrical conductivity

&ilicon dioxide doesn't have any mobile electrons or ions / so it doesn't conduct electricity either as a solid or a li4uid.

The molecular oxides Phosphorus sulphur and chlorine all form oxides which consist of molecules. &ome of these molecules are fairly simple / others are polymeric. 2e are 0ust 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 to its neighbors9 will be van der 2aals dispersion forces or dipole/dipole interactions. The strength of these will vary depending on the sie of the molecules.

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"one of these oxides conducts electricity either as solids or as li4uids. "one of them contains ions or free electrons. The phosphorus oxides

Phosphorus has two common oxides phosphorus (---) oxide P $, and phosphorus (7) oxide P$. #hosphorus (III! oxide (tetraphosphorus hexoxide! 

Phosphorus (---) oxide is a white solid melting at #6* and boiling at +36*. The phosphorus is using only three of its outer electrons (the 3 unpaired p electrons) to form bonds with the oxygens. #hosphorus (! oxide (tetraphosphorus decoxide! 

Phosphorus (7) oxide is also a white solid subliming (turning straight from solid to vapour) at 36*. -n this case the phosphorus uses all five of its outer electrons in the bonding. &olid phosphorus(7) oxide exists in several different forms / some of them polymeric. 2e are going to concentrate on a simple molecular form and this is also present in the vapour. This is most easily drawn starting from P$,. The other four oxygens are attached to the four phosphorus atoms via double bonds.

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The sulphur oxides

&ulphur has two common oxides sulphur dioxide (sulphur (-7) oxide) &$# and sulphur trioxide (sulphur (7-) oxide) &$3. Sulphur dioxide 

&ulphur dioxide is a colourless gas at room temperature with an easily recognied choing smell. -t consists of simple &$# molecules.

The sulphur uses  of its outer electrons to form the double bonds with the oxygen 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. -t reacts very rapidly with water vapour in the air to form sulphuric acid. That means that if  you mae some in the lab you tend to see it as a white sludge which fumes dramatically in moist air (forming a fog of sulphuric acid droplets). :aseous sulphur trioxide consists of simple &$ 3 molecules in which all six of the sulphur's outer electrons are involved in the bonding.

There are various forms of solid sulphur trioxide. The simplest one is a trimer &3$; where three &$3 molecules are 0oined up and arranged in a ring. The fact that the simple molecules 0oin up in this way to mae bigger structures is what maes the sulphur trioxide a solid rather than a gas.

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The chlorine oxides

*hlorine forms several oxides.
*hlorine (-) oxide is a yellowish/red gas at room temperature. -t consists of simple small molecules.

There's nothing in the least surprising about this molecule and it's physical properties are 0ust what you would expect for a molecule this sie.

$hlorine (II! oxide (dichlorine heptoxide! 

-n chlorine (7--) oxide the chlorine uses all of its seven outer electrons in bonds with oxygen. This produces a much bigger molecule and so you would expect its melting point and boiling point to be higher than chlorine (-) oxide. *hlorine (7--) oxide is a colourless oily li4uid at room temperature. The diagram has been drawn as a standard structural formula for simplicity. -n fact the shape is tetrahedral around both chlorines and 7/shaped around the central oxygen.

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PROPERTI ESOFTHEPERI OD3CHLORI DES This page loos at the structures of the chlorides of the Period 3 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!

"a*l

%g*l#

Al*l3

&i*l

P*l@

&#*l#

P*l3 As you will see later aluminum chloride exists in some circumstances as a dimer Al#*l,. The structures

&odium chloride and magnesium chloride are ionic and consist of giant ionic lattices at room temperature Aluminum chloride and phosphorus (7) chloride are tricy5 They change their structure from ionic to covalent when the solid turns to a li4uid or vapour. There is much more about this later on this page. The others are simple covalent molecules. Melting and boiling points

&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 strong ionic attractions.

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The rest are li4uids or low melting point solids. eaving aside the aluminum chloride and phosphorus (7) chloride cases where the situation is 4uite complicated the attractions in the others will be much weaer intermolecular forces such as van der 2aals dispersion forces. These vary depending on the sie and shape of the molecule but will always be far weaer than ionic bonds. Electrical conductivity

&odium and magnesium chlorides are ionic and so will undergo electrolysis  when they are molten. Blectricity is carried by the movement of the ions and their discharge at the electrodes. -n the aluminum chloride and phosphorus (7) chloride cases the solid doesn't conduct electricity because the ions aren't free to move. -n the li4uid (where it exists / 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

As an approximation the simple ionic chlorides (sodium and magnesium chloride)  0ust dissolve in water. The other chlorides all react with water in a variety of ways described below for each 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. A 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 heat energy to brea and so sodium chloride has high melting and boiling points. -t doesn't conduct electricity in the solid state because it hasn't any mobile electrons and the ions aren't free to move.
%agnesium chloride is also ionic but with a more complicated arrangement of the ions to allow for having twice as many chloride ions as magnesium ions. Again lots of heat energy is needed to overcome the attractions between the ions and so the melting and boiling points are again high. &olid magnesium chloride is a non/conductor of electricity because the ions aren't free to move.
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-ons of this sort are acidic / the degree of acidity depending on how much the electrons 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 more easily pulled off by a base. -n the magnesium case the amount of distortion is 4uite small and only a small proportion of the hydrogen atoms are removed by a base / in this case by water molecules in the solution.

The presence of the hydronium ions in the solution causes it to be acidic. The fact that there aren't many of them formed (the position of e4uilibrium lies well to the left) means that the solution is only wealy acidic. Gou may also find the last e4uation in a simplified form!


Blectronegativity increases as you go across the period and by the time you get to aluminum1 there isn't enough electronegativity difference between aluminum and chlorine for there to be a simple ionic bond. Aluminum chloride is complicated by the way its structure changes as temperature increases. At room temperature the aluminum in aluminum chloride is ,/coordinated. That means that each aluminum is surrounded by , chlorines. The structure is an ionic lattice / although with a lot of covalent character. At ordinary atmospheric pressure aluminum chloride sublimes (turns straight from solid to vapour) at about H6*. -f the pressure is raised to 0ust over # atmospheres it melts instead at a temperature of ;#6*. 8oth of these temperatures of course are completely wrong for an ionic compound / they are much too low. They suggest comparatively wea attractions between molecules / not strong attractions between ions.

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The coordination of the aluminum changes at these temperatures. -t becomes / coordinated / each aluminum now being surrounded by  chlorines rather than ,. 2hat happens is that the original lattice has converted into Al #*l, molecules.

This conversion means of course that you have completely lost any ionic character / which is why the aluminum chloride vapories or melts (depending on the pressure). There is an e4uilibrium between these dimers and simple Al*l3 molecules. As the temperature increases further the position of e4uilibrium shifts more and more to the right.

&ummary  •

•

•

At room temperature solid aluminum chloride has an ionic lattice with a lot of covalent character. At temperatures around H / ;6* (depending on the pressure) aluminum chloride coverts to a molecular form Al#*l,. This causes it to melt or vaporie because there are now only comparatively wea intermolecular attractions. As the temperature increases a bit more it increasingly breas up into simple Al*l3 molecules.

&olid aluminum chloride doesn't conduct electricity at room temperature because the ions aren't free to move. %olten aluminum chloride (only possible at increased pressures) 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 0ust dissolving in it. -n the first instance hexaa4uaaluminum ions are formed together with chloride ions.

Gou will see that this is very similar to the magnesium chloride e4uation given above / the only real difference is the charge on the ion. That extra charge pulls electrons from the water molecules 4uite strongly towards the aluminum. That maes the hydrogens more positive and so easier to remove from the ion. -n other words this ion is much more acidic than in the corresponding magnesium case. These e4uilibria (whichever you choose to write) lie further to the right and so the solution 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 the two to form ionic bonds. &ilicon tetrachloride is a colourless li4uid at room temperature which fumes in moist air. The only attractions between the molecules are van der 2aals dispersion forces. -t doesn't conduct electricity because of the lac of ions or mobile electrons. -t fumes in moist air because it reacts with water in the air to produce hydrogen chloride. -f you add water to silicon tetrachloride there is a violent reaction to produce silicon dioxide and fumes of hydrogen chloride. -n 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 (---) chloride P*l3 and phosphorus (7) chloride P*l@. #hosphorus (III! chloride (phosphorus trichloride!' #$l ) 

This is another simple covalent chloride / again a fuming li4uid at room temperature. -t is a li4uid because there are only van der 2aals dispersion forces and dipole/ dipole attractions between the molecules. -t doesn't conduct electricity because of the lac of ions or mobile electrons. Phosphorus (---) chloride reacts violently with water. Gou get phosphorous acid <3P$3 and fumes of hydrogen chloride (or a solution containing hydrochloric acid if lots of water is used).

#hosphorus (! chloride (phosphorus pentachloride!' #$l + 

Infortunately phosphorus (7) chloride is structurally more complicated. Phosphorus (7) chloride is a white solid which sublimes at ,36*. The higher t he temperature goes above that the more the phosphorus (7) chloride dissociates (splits up reversibly) to give phosphorus (---) chloride and chlorine.

&olid phosphorus (7) chloride contains ions / which is why it is a solid at room temperature. The formation of the ions involves two molecules of P*l@. A chloride ion transfers from one of the original molecules to the other leaving a positive ion EP*l FD and a negative ion EP*l,F/. At ,36* the phosphorus (7) chloride converts to a simple molecular form containing P*l@ molecules. 8ecause there are only van der 2aals dispersion forces between these it then vapories. &olid phosphorus (7) chloride doesn't conduct electricity because the ions aren't free to move.

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Phosphorus (7) chloride has a violent reaction with water producing fumes of hydrogen chloride. As with the other covalent chlorides if there is enough water present these will dissolve to give a solution containing hydrochloric acid. The overall e4uation in boiling water is!

,-isulphur dichloride' S .$l.,

Jisulphur dichloride is a simple covalent li4uid / orange and smelly5 The shape is surprisingly difficult to draw convincingly5 The atoms are all 0oined up in a line / but twisted!

The reason for drawing the shape is to give a hint about what sort of intermolecular attractions are possible. There is no plane of symmetry in the molecule and that means that it will have an overall permanent dipole. The li4uid will have van der 2aals dispersion forces and dipole/dipole attractions. There are no ions in disulphur dichloride and no mobile electrons / so it never conducts electricity. Jisulphur dichloride reacts slowly with water to produce a complex mixture of things 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 e4uation 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 oxides depend on the type of bonding. "a#$ Al#$3 and %g$ are ionic oxides and hence have a high melting point. %g$ and Al#$3 have a higher melting point than "a#$ since the charges are higher resulting in a stronger attraction between the ions. &i$# has a giant covalent structure and hence a high melting point. There are strong covalent bonds between all the atoms and thus lots of energy is re4uired to brea them. P$ and &$3 are molecular covalent and so only intermolecular forces exist between the molecules. The melting points are thus much lower. P$ is a much bigger molecule than &$3 and so has a much higher melting point as the van der 2aal9s forces are stronger.

Blement

"a

%g

Al

&i

P

&

Kormulae of oxide

"a#$

%g$

Al#$3

&i$#

P$

&$3

&tructure of oxide

-onic

-onic

%ostly ionic

:iant covalent

%olecular covalent

%olecular covalent

%elting point of oxide L6*

#+@

#H@#

#+#

+3

3

/

#. Acid0base character of oxides -onic oxides contain the $#/ ion. This is a strongly basic ion which reacts with water to produce hydroxide ions! $#/(a4) D <#$(l)  #$
%$(s) D <#$(l)  %$($<)/(a4)D
"a#$(s) D <#$(l)  #"a$<(a4) "a#$(s) D #
%g$(s) D <#$(l)  %g($<)#(s)  %g($<)#(a4) %g$(s) D #
Al#$3(s) D ,
&i$#(s) D #$
P$(s) D ,<#$(l)  <3P$(a4)

p< C 3

S1. and S1) are acidic oxides . They dissolve in water to give acidic solutions and also react with alalis!

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&$#(g) D <#$(l) CC <#&$3(a4)

p< C #

&$3(g) D <#$(l)  <#&$(a4)

p< C 

&$# is a waste gas in many industrial processes. -t is harmful because it dissolves in rain water to give acid rain. -t can be removed from waste gases because it dissolves in alali and so it is passed through an alaline solution in waste gas outlets to minimie the amount which escapes into the atmosphere.

T4E $451%I-ES 16 #E%I1- ) E5EMETS

All the elements of Period 3 except argon combine directly with chlorine to give chlorides. #"a(s) D *l#(g)  #"a*l(s) a$l is an ionic chloride .

%g(s) D *l#(g)  %g*l#(s) Mg1 is also an ionic chloride.

#Al(s) D 3*l#(g)

 #Al*l3(s)

Al$l) is covalent . -t forms a polymeric structure in the solid state turning 4uicly on heating into a dimeric gas (Al#*l,). -t thus behaves as a simple molecular chloride.

&i(s) D #*l#(g)  &i*l(s) Si$l* is a molecular covalent chloride.

P(s) D ,*l#(g)  P*l3(s) P*l@ is formed by reacting P*l3 with excess chlorine in a reversible reaction! P*l3(l) D *l#(g) CC P*l@(s) #$l+ is actually ionic in the solid state  / it exists as EP*l FDEP*l,F/ in the solid state.

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The acid/base properties of the oxides of Period 3 can be summaried in the following table!

Blement

"a

%g

Al

&i

P

& &$#

Kormulae of oxides

"a#$

%g$

Al#$3

&i$#

P$

Acid/base character of oxide

8asic

&$3

p< of solution when dissolved in water

# / 

8asic

Amphoteric

Acidic

+

+

H/; insoluble

Acidic

Acidic

#/

#/ (&$#)

insoluble /3 (&$3)

The oxides therefore become more acidic on moving from left to right in the periodic table. #hysical properties of chlorides a$l and Mg$l. are ionic chlorides . &ince a large amount of energy is re4uired 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 ionic chlorides.

Al*l3 actually exists in polymeric form in the solid state which is converted to a dimeric form in the gas phase. At high temperatures it reverts to a simple molecular structure! Cl

Cl Cl

 Al n

Cl

Cl

n

 Al

 Al

n/2 Cl

Cl

Cl

Cl

Cl

complex polymer 

dimer 

(solid)

(gas - at low T)

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Cl

 Al Cl

monomer  (gas - at high T)

The aluminum atom is electron deficient = it has only 3 of its four valence orbitals occupied so it has an empty orbital with which it can accept a lone pair of electrons from a *l atom on an ad0acent monomer. #$l+ is ionic  so its melting point is thus high. $n heating however it reverts to a simple covalent structure and sublimes.

Blement Kormula of chloride &tructure of chloride %elting point of chloride L6*

"a "a*l

%g %g*l#

Al Al*l3

&i &i*l

P P*l@

ionic

ionic

polymer

-onic

H

+

H

molecular covalent @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 chlorides  dissolve in water to give neutral solutions! "a*l(s)  "aD(a4) D *l/(a4) p< C + #D / %g*l#(s)  %g (a4) D #*l (a4) p< C + Aluminium chloride reacts with water to give hydrated aluminum ions and chloride ions. The hydrated aluminium ions undergo deprotonation to give an acidic solution! Al*l3(s) D ,<#$(l)  Al(<#$),3D(a4) D 3*l/(a4) Al(<#$),3D(a4) D <#$(l)  EAl(<#$)@($<)F#D(a4) D <3$D(a4) The other covalent chlorides  react readily with water at room temperature to form the oxide or hydroxide and <*l(g). The <*l is formed as white misty fumes and the observance of these fumes is a good indication that the chloride is covalent.

&i*l(l) D #<#$(l)  &i$#(s) D <*l(g) P*l@(s) D <#$(l)  <3P$(a4) D @<*l(g)

p< C  / # p< C  / #

$ovalent chlorides  thus react with water to give acidic solutions. The acidity is due to dissolved <*l.

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The water molecules attac the covalent chlorides by donating lone pairs of electrons into empty low/lying orbitals on the electropositive atoms. -n the case of Al*l3 there is an available 3p orbital and in &i*l  and P*l@ there are available d/ orbitals! 3s

3p 3d

Al*l3

E"eF

↓↑

↓↑

↓↑

&i*l

E"eF

↓↑

↓↑

↓↑

↓↑

P*l@

E"eF

↓↑

↓↑

↓↑

↓↑

↓↑

 -t the availability of these low/lying empty orbitals which enables these chlorides to react readily with water. %eaction of $l. &ith &ater

*l#(g) D <#$(l)  <*l$(a4) D <*l(a4) The resulting solution contains <*l(a4) and is thus acidic (p< C #).

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