The p-Block Elements
Group 15: Occurrence and Atomic Properties
The differentiating electron in group 15 elements tends to enter the p orbital, resulting in a valence shell electronic configuration of ns 2np 3.
The atomic and ionic radii of group 15 elements increase in size down the group due to the addition of a new principal energy level in each successive element.
There is a considerable increase in the covalent radius from nitrogen to phosphorus, which is due to the effective shielding of the s and p electrons present in the penultimate shell of phosphorus.
Group 15 elements show a higher value of ionisation enthalpy as compared to group 14 elements.
As we move down the group, the ionisation energy gradually decreases.
As we move down the group, the electronegativity gradually decreases.
Group 15: Physical Properties and Oxidation States
Group 15 elements are less metallic as compared to group 14 elements, but their metallic character increases down the group.
The melting point of the elements increases from nitrogen to arsenic and decreases from arsenic to bismuth.
The boiling point gradually increases from nitrogen to antimony.
The density of the elements increases regularly from nitrogen to bismuth.
All the elements, except for bismuth, show allotropy.
Nitrogen shows an oxidation state of -3 in nitrides by gaining electrons.
The elements of this group form covalent bonds and show a formal oxidation state of -3.
The elements show positive oxidation states of +3 and +5.
In group 15 elements, the covalent character decreases from nitrogen to bismuth.
Nitrogen exhibits various oxidation states from -3 to +5.
Group 15: Anomalous Properties
The unique properties of nitrogen are attributed to its
• Small atomic size • High electro-negativity or high ionisation enthalpy • Non availability of d-orbitals • Propensity to form multiple bonds
The tendency of nitrogen to form pp-pp bonds is one of the major features that distinguishes it from phosphorous and other group 15 elements.
The other elements do not form pp-bonds because of their relatively larger size.
The catenation tendency is less for nitrogen when compared to the other elements of the group.
Nitrogen doesn’t form dp-pp or dp-dp bonds due to the absence of d-orbitals.
Group 15: Chemical Reactivity
The stability of hydrides decreases from NH 3 to BiH 3. This is due to an increase in the size of the central atom down the group.
Hydrides of group 15 elements are good reducing agents. The reducing character of hydrides increases from NH 3 to BiH 3.
NH 3 is the strongest base among all the hydrides. Nitrogen forms five oxides with oxidation state ranging from +1 to +5, while the other elements form oxides only in +3 and +5 oxidation states.
The lower oxides of nitrogen are neutral, while the higher oxides are acidic.
All the elements of group +15 form trihalides and pentahalides.
All the elements of group 15 react with metals to form their binary compound showing
-3 oxidation state.
Group 15: Nitrogen – Dinitrogen
Commercially, dinitrogen is mainly obtained by the fractional distillation of liquid air.
In the laboratory, dinitrogen is generally prepared by gently heating equimolar aqueous solutions of ammonium chloride and sodium nitrite.
NH 4Cl(aq) + NaNO 2(aq) → NaCl(aq) + NH 4 NO 2(aq) Ammonium Chloride Sodium Nitrite Sodium Chloride Ammonium Nitrite
NH 4NO 2(aq) → 2H 2O(vap) + N 2(g)
Sodium azide or barium azide, when heated carefully to about 573K, undergoes thermal decomposition to produce dinitrogen.
2NaN 3 → 2Na + 3N2 Sodium Azide Sodium Dinitrogen Gas
Ba(N 3)2 → Ba + 3N 2 Barium Azide Barium Dinitrogen Gas
Dinitrogen is a colourless, odourless and tasteless non-toxic gas. A molecule of dinitrogen consists of a triple bond that has a very high bond dissociation energy of 945.4 kJmol -1. Hence, dinitrogen is inert at room temperature. At high temperatures, dinitrogen reacts directly with metals such as magnesium, calcium and aluminium to form the respective nitrides.
N2 + 3Mg → Mg 3 N2 Dinitrogen Magnesium Magnesium Nitride
N2 + 3Ca → Ca 3 N2 Dinitrogen Calcium Calcium Nitride
3N 2 + 6Al → 6AlN Dinitrogen Aluminium Aluminium Nitride
It also reacts with hydrogen at high temperature and pressure in the presence of a catalyst.
ℎ
N2 + 3H 2 2NH 3 Dinitrogen Hydrogen Ammonia
It is widely used as an inert atmosphere in iron, steel and other metallurgical industries.
A large amount of dinitrogen is used to prepare ammonia and nitrolim.
Liquid dinitrogen is used as a refrigerant to preserve biological material.
Lsn 6: Group 15: Nitrogen - Ammonia
Ammonium chloride, when heated with lime or caustic soda, evolves ammonia gas. ∆ 2NH 4Cl + Ca(OH) 2 2NH 3 + 2H 2O + CaCl 2 Ammonium Chloride Slaked Lime Ammonia Water Calcium Chloride
In accordance with Le Chatelier’s principle, high pressure and low temperature favours the better yield of ammonia.
A high pressure of 200-300 atmospheres, an optimum temperature about 700K and the catalyst iron oxide along with promoter alumina are used to obtain better yields of ammonia.
Fe 3O4, Al 2O3 200-300 atm -1 N2 + 3H 2 2NH 3 ∆H = -92.2 KJ mol Nitrogen Hydrogen 700K Ammonia
On account of the presence of one lone pair in ammonia, the geometry of the molecule is trigonal pyramid.
An aqueous solution of ammonia precipitates hydroxides of metal from their aqueous salt solutions.
An aqueous solution of ammonia combines with copper sulphate solution to form a deep blue coppertetraammine complex.
CuSO 4 + 2NH 4OH Cu(OH) 2 + (NH4)2SO 4 Blue ppt.
Cu(OH) 2 + (NH 4)2SO 4 + 2NH 4 OH [ Cu(NH 3)4 ] SO 4 + 4H2O Deep Blue
Ammonia is used in the manufacture of certain industrially important chemicals such as nitric acid, baking soda and washing soda.
Group 15: Nitrogen – Oxides
N2O is prepared by heating NH 4NO 3 gently upto 240 °C.
NH NO + N O + 2H O 4 3 240 ° 2 2 Ammonium nitrate Nitrous oxide
NO is prepared by the reduction of nitrite salt with FeSO 4.
2NaNO 3 + 2FeSO 3 + 3H 2 SO 4 Fe(SO 4)3 + 2NaHSO 4 + 2H 2O + 2NO Sodium nitrite Ferrous sulphate Nitric oxide
N2O and NO are both neutral in nature.
Dinitrogen trioxide (N 2O3) is obtained as a blue liquid when a mixture of NO and NO 2
or N 2O4 is cooled to below -20 °C.
NO 2 is obtained by heating Pb(NO 3)2 at about 673 K.
2Pb(NO ) 4N O + 2PbO + O 3 2 673 2 2 Lead nitrate Nitrogen dioxide
At low temperatures, NO 2 associates to a solid dimer called N 2O4.
N2O5 is prepared by the dehydration of HNO 3 with P 4O10.
4HNO 3 + P4O10 4HPO 3 + 2N 2O5 Nitric acid Phosphorus Metaphosphoric Dinitrogen pentoxide acid pentoxide
N2O3, NO 2 or N 2O4 and N2O5 are acidic in nature.
Group 15: Nitrogen – Nitric Acid
In the laboratory, HNO 3 is prepared by heating NaNO 3 with concentrated H 2SO 4.
On a commercial scale, HNO 3 is prepared by Ostwald’s process. The steps are:
• Catalytic oxidation of NH 3 into NO
• Oxidation of NO to NO 2
• Formation of HNO3 from NO 2
+ − An aqueous solution of HNO 3 undergoes ionisation to produce H 3O and NO 3.
Hot and concentrated HNO 3 acts as a powerful oxidising agent.
Concentrated HNO 3 reacts with “Cu” and “Zn” to give NO 2.
Dilute HNO 3 reacts with “Cu” to give NO, and with “Zn” to give N 2O.
Noble metals dissolve in aqua regia, of which HNO 3 is a constituent.
Concentrated HNO 3 oxidises non-metals to their corresponding higher oxoacids. The brown ring test confirms the presence of nitrate ions in a given salt.
Group 15: Nitrogen – Allotropes
The existence of an element in more than one physical form is called allotropy. The different physical forms of the same element are called allotropes.
White phosphorus is a soft, waxy and translucent solid.
The angular strain in the molecule makes white phosphorus unstable and highly reactive.
Red phosphorus is an iron-grey lustrous crystalline solid that is odourless, non-toxic and insoluble in water.
Black phosphorus is highly polymerised and thermally the most stable allotrope of phosphorus.
Group 15: Phosphorus – Preparation of Phosphine
Phosphine is a hydride of phosphorus analogous to ammonia.
The hydrolysis of metallic phosphides with water or dilute mineral acids like hydrochloric acid or sulphuric acid produces phosphine.
On a laboratory scale, phosphine is prepared by hydrolyzing white phosphorus with a concentrated solution of sodium hydroxide.
Pure phosphine can be obtained by using alcoholic potassium hydroxide instead of an aqueous sodium hydroxide solution.
Phosphine has pyramidal geometry with a bond angle of 93.5 °.
Group 15: Phosphorus – Properties and Uses of Phosphine
Phosphine is a colourless and extremely poisonous gas that smells of rotten fish.
Phosphine is slightly soluble in water, and more soluble in carbon disulphide and other organic solvents.
The bond angle in the phosphonium ion is greater than in phosphine because the extent of repulsive interaction in it is less than in phosphine.
An aqueous solution of phosphine in the presence of light decomposes to red phosphorus and hydrogen.
It burns with an explosion when it comes in contact with even small amounts of oxidising agents.
It precipitates metal phosphides from their aqueous metal salt solutions.
It finds use in the manufacture of Holme’s signals and smoke screens.
Group 15: Phosphorus – Halides
In the laboratory, PCl 3 is prepared by heating white phosphorus in a current of dry chlorine.
PCl3 has pyramidal geometry. It acts as a Lewis base due to its ability to denote its lone pair of electrons.
In the laboratory, PCl 5 is prepared by passing excess chlorine into a solution of PCl 3 in
CCl 4.
PCl 3 reacts violently with water and gets hydrolysed to form H 3PO 3.
PCl 5 reacts with excess and gets hydrolysed to H 3PO 4.
Both PCl 3 and PCl 5 react with organic compounds that contain hydroxyl groups.
PCl 5 dissociates into PCl 3 and chlorine on strong heating.
Trigonal bipyramidal geometry is observed for PCl 5 only in its liquid and gaseous state.
It exists as a salt in the solid state.
Group 15: Phosphorus – Preparation of Oxoacids
Phosphorus usually exhibits an oxidation state of +3 or +5 in its oxoacids.
The oxidation state of Phosphorus in Hypophosphorus acid is +1 and +4 in
Pyrophosphorous acid.
The oxoacids in which Phosphorus exhibits lower oxidation state of +1 or +3constitute the “Phosphorous acid series”.
The oxoacids in which Phosphorus exhibits higher oxidation state of +4 or +5constitute the “Phosphoric acid series”.
Hypophosphorous acid is prepared by the alkaline hydrolysis of white Phosphorus. ∆ 2P4 + 3Ba(OH) 2 + 6H 2O 3Ba(H 2PO 2)2 + 2PH3 White phosphorus Barium hydroxide Water Barium hypophosphite Phosphine
Ba(H 2PO 2)2 + H 2SO 4(dil) BaSO 4 + 2H 3PO 2 Barium hypophosphite Sulphuric acid Barium sulphate Hypophosphorous acid
The hydrolysis of Phosphorous trioxide or Phosphorous trichloride gives Orthophosphorous acid.
Pcl 3 + 3H2O H3PO 3 + 3HCl Phosphorous trichloride Water Orthophosphorous acid Hydrochloric acid
Orthophosphoric acid is industrially the most important oxoacid of Phosphorus. It is prepared by the hydrolysis of Phosphorous pentoxide. ∆ P4O10 + 6H 2O 4H 3PO 4 Phosphorous pentoxide Water Orthophosphoric acid
Crystalline Phosphorous acid, on being heated with Bromine in a sealed tube, gives Metaphosphoric acid. ∆ H3PO 3 + Br 2 HPO 3 + 2HBr Phosphorous acid Bromine Metaphosphoric acid Hydrobromic acid
Group 15: Phosphorus – Structural Aspect of Oxoacids
In all oxoacids of phosphorus, the central phosphorus atom is surrounded tetrahedrally by other atoms.
At least one P = O and P – OH bond are present in all of these oxoacids.
P – H bonds are present only in the oxoacids belonging to the phosphorous acid series.
Only the hydrogen atoms that are attached to oxygen in P –OH bonds are ionisable and impart acidic character to the compound.
The P – H bonds found in the oxoacids belonging to the “phosphorous acid series” have reducing properties.
Hypophosphorous acid reduces silver nitrate to metallic silver.
Orthophosphoric acid is a tribasic acid.
The oxoacids in lower oxidation states of phosphorus +1 and +3 tend to disproportionate to higher and lower oxidation states.
Group 16: Occurrence and General Characteristics
Group 16 elements include oxygen, sulphur, selenium, tellurium and polonium.
The ionisation enthalpy decreases down the group.
Electron gain enthalpy is less negative for oxygen and increases from sulphur to polonium.
Oxygen is the most electronegative element in the group.
The elevtronegativity decreases down the group.
The general electronic configuration of these elements is ns 2np 4.
The atomic size increases down the group.
The ionisation enthalpy decreases down the group.
Electron gain enthalpy is less negative for oxygen and increases from sulphur to polonium.
Oxygen is the most electronegative element in the group.
The electronegativity decreases down the group.
Group 16: Chemical Properties
The common oxidation states exhibited by the elements of group 16 include -2, +2, +4 and +6.
Oxygen shows +2 oxidation state in oxygen difluoride and +1 in dioxygen difluoride.
Water has an abnormally high boiling point because its molecules are associated with each other by means of hydrogen bonds.
Hydrogen bond formation involving sulphur atoms is rare and extremely weak, and hence, hydrogen sulphide remains unassociated and exists as a gas.
Acidic character increases from H 2O to H 2Te due to the decrease in the H-E bond dissociation enthalpy.
Hexafluorides undergo sp 3d2 hybridisation and have octahedral geometry.
Tetrafluorides undergo sp 3d hybridisation and have see-saw geometry.
Dihalides are formed by sp 3 hybridisation and have tetrahedral geometry.
The structure of the monohalides is similar to that of hydrogen peroxide.
Group 16: Oxygen – Dioxygen
Commercially, dioxygen is prepared by either the fractional distillation of liquid air or by the electrolysis of water.
Oxygen is slightly denser than air and is sparingly soluble in water.
The small amount of oxygen dissolved in water supports aquatic life.
In spite of having an even number of electrons, dioxygen exhibits paramagnetic behavior.
The reaction of oxygen with other elements is highly exothermic.
Most metals burn in dioxygen to form oxides that are mostly basic in nature.
Most non-metals burn in dioxygen to form acidic oxides.
Oxygen has many important uses in our life.
Group 16: Oxygen – Simple Oxides
A binary compound of oxygen with another element is called an oxide.
An oxide that contains just as much oxygen as permitted by the normal valency of its metal is known as a simple oxide.
Oxides of non-metals or metals of higher oxidation states are acidic in nature.
Metal oxides are basic in nature.
Some metal oxides like zinc oxide or aluminium oxide are amphoteric oxides.
Neutral oxides do not show any tendency to form salts when treated with acids or bases.
Group 16: Oxygen – Ozone
Ozone is an unstable triatomic allotropic form of oxygen.
Ozone protects the earth’s surface from excessive exposure to ultra violet radiation.
Chlorofluorocarbons and oxides of nitrogen, particularly nitric oxide, deplete the ozone layer.
Ozone is thermodynamically unstable and decomposes to oxygen.
Ozone acts as a powerful oxidising agent.
Ozone is a pale blue gas with a characteristic fishy smell.
Ozone has many important uses.
Group 16: Sulphur – Allotropes
The two common crystalline forms of sulphur or rhombic sulphur, and b or monoclinic sulphur.
Rhombic sulphur is stable at room temperature, while monoclinic sulphur is stable above 369 K.
Both the forms of sulphur are stable at 369 K. This temperature is called transition temperature.
Both the forms contain puckered S8 rings with a crown conformation.
∗ Like O 2, the S 2 molecule contains two unpaired electrons in the anti-bonding π orbitals and exhibits paramagnetism.
Group 16: Sulphur – Sulphur Dioxide
Sulphur dioxide may be prepared in the laboratory by the action of dilute sulphuric acid on sulphites.
Large volumes of sulphur dioxide are prepared by roasting sulphide ores.
Levels of sulphur dioxide above 5 parts per million are poisonous.
It is an acidic oxide that is highly soluble in water.
In the presence of moisture, it can liberate nascent hydrogen and act as a reducing agent.
Sulphur dioxide may be detected in the laboratory:
• By the decolourisation of a potassium permanganate solution. • By the change in the colour of a filter paper moistened with acidified potassium dichromate solution to green. • By starch iodate paper turning blue.
Group 16: Sulphur – Oxoacids
Oxoacids with S – S linkages are called thio acids, while those with peroxo linkage are called peroxo acids.
Commercially, sulphuric acid is manufactured by the contact process.
Low temperature and high pressure are the conditions favourable for maximum yield of sulphuric acid.
The high boiling point and viscous nature of sulphuric acid is due to hydrogen bonding.
Sulphuric acid is always diluted by adding it to water slowly with constant stirring and not by adding water to it.
Sulphuric acid is a strong diabasic acid.
The acid neutralises alkalies and forms two series of salts: bisulphates and sulphates.
Concentrated sulphuric acid is a powerful dehydrating agent.
Sulphuric acid, because of its wide applications, is referred as the king of chemicals.
Group 17: Occurrence
The elements in group 17 are F, Cl, Br, I and At.
These elements are collectively referred to as the “halogens,” which means “salt- producing.”
The valence shell electronic configuration of these elements is ns 2np 5.
Owing to their high reactivity, the halogens do not occur in free state, but in combined state in nature.
Fluorine occurs widely as insoluble fluorides.
Chlorine, bromine and iodine are present in sea water in the form of chlorides, bromides and iodides.
The principal source of iodine is sea weeds and crude chile saltpeter.
Group 17: Atomic Properties
The atomic radii and ionic radii increase due to the addition of a new principal energy level in each successive element.
These elements have the least atomic radii when compared to other elements in the corresponding periods.
Group 17 elements show very high values of ionisation enthalpy.
Halogens have the maximum negative electron gain enthalpy in the respective periods.
Fluorine has less negative electron gain enthalpy than chlorine because of the small size and compact 2p sub-shell of its atom.
Fluorine is the most electronegative element in the periodic table.
Group 17: Physical Properties and Oxidation States
Group 17 elements are also known as halogens. These are highly electro-negative elements.
Fluorine is a pale yellow gas, chlorine is a greenish yellow gas, bromine is a reddish brown liquid and iodine is a dark violet solid.
The common negative oxidation state of these elements is -1.
The positive oxidation states of these elements are observed in several compounds.
Group 17: Chemical Properties
Group 17 elements are more reactive than all the other elements in the periodic table. They react with metals and non-metals to form halides.
The reactivity of these elements decreases down in the group.
They are good oxidising agents as they are ready to accept an electron from other elements. Their oxidising capacity decreases down the group.
Fluorine forms only two oxides that are thermally stable at room temperature, and are fluorinating agents.
Chlorine forms oxides that are powerful bleaching agents.
Bromine forms oxides that are less stable and act as powerful oxidising agents.
Iodine forms oxides that are solids and insoluble in water.
The anomalous behavior of fluorine is due to its high electro-negativity, small atomic size, low bond dissociation energy and no availability of d-orbital in the valence shell.
Group 17: Chlorine
Chlorine is a greenish-yellow gas.
It is prepared by heating manganese dioxide with concentrated hydrochloric acid.
Commercially, chlorine gas is manufactured by Electrolysis and Deacon’s process.
Chlorine combines with hydrogen in light with an explosion to form hydrochloric acid.
Chlorine is good oxidising and a powerful bleaching agent. Its bleaching action is due to oxidation.
Chlorine forms bleaching powder with slaked lime.
Chlorine is used as a bleaching agent in the wood pulp, cotton and textile industries.
Chlorine is used in the preparation of poisonous gases such as phosgene, tear gas and mustard gas.
Group 17: Chlorine – Hydrogen Chloride
Hydrogen chloride is prepared in the laboratory by treating sodium chloride with concentrated sulphuric acid.
420 NaCl + H2SO 4 NaHSO 4 + HCl Common salt Conc. Sulphuric acid Sodium bisulphate Hydrogen chloride (Sodium chloride)
Hydrogen chloride is a colourless gas with a pungent odour.
An aqueous solution of hydrogen chloride is also called hydrochloric acid.
+ Because of its ability to dissociate almost 100% to produce hydronium ions (H 3O ), hydrochloric acid is considered a strong acid.
+ - ( )( ) 7 HCl + H O H O + Cl Ka = = 10 (g) 2 (l) 3 (aq) (aq) ( ) Hydrogen chloride Water Hydrochloric acid
Hydrochloric acid:
• Reacts with ammonia to form dense fumes of ammonium chloride.
NH3 + HCl NH 4Cl Ammonia Hydrochloric acid Ammonium chloride • Decomposes salts of weaker acids.
Na2 CO 3 + 2HCl 2NaCl + CO 2 + H2O Sodium Carbonate Hydrochloric acid Sodium chloride Carbon dioxide Water
Na2 SO3 + 2HCl 2NaCl + SO2 + H2O Sodium sulphite Hydrochloric acid Sodium chloride Sulphur dioxide Water
Hydrochloric acid:
• Is used in the manufacture of chlorine, and chlorides like ammonium chloride.
NaOCl + 2HCl Cl2 + H2O + NaCl Sodium hypochlorite Hydrochloric acid Chlorine Water Sodium chloride
NH3 + HCl NH 4Cl Ammonia Hydrochloric acid Ammonium chloride
Group 17: Oxoacids of Halogens
Halogens usually form four series of oxoacids, namely:
• Hypohalous acids • Halous acids • Halic acids and • Perhalic acids
The central halogen atom is sp 3 hybridised in all the oxoacids.
One X-OH bond is essentially present in every oxoacid. In most of these oxoacids, “X=O” bonds are present.
The oxidation state of the halogen is +1 in hypohalous acids, +3 in halous acids, +5 in halic acids and +7 in perhalic acids.
The hypohalite ion has a linear shape, while chlorite ion has a V shape.
The halate ion is pyramidal in shape, while the perhalate ion is tetrahedral.
The acidic strength of an oxoacid increases with an increase in the oxidation number of the halogen.
Perchloric acid is the strongest acid among all the oxoacids of halogens.
Group 17: Interhalogen Compounds
Interhalogens are highly reactive.
Interhalogens exist in all three different physical states.
Interhalogens are named as Halogen Halides.
Interhalogens are prepared in two methods that is by direct combination of halogens and by reaction of halogens with lower interhalogens.
Halogens are used as halogenating agents in preparation of poly halogens and in enrichment of uranium.
Group 18: Occurrence and Atomic Properties
Helium, Neon, Argon, Krypton, Xenon and Radon are collectively referred to as noble gases.
Due to the duplet or octet configuration in their valence shells, the elements of this group are highly stable.
As these elements do not have a tendency to loose, gain or share electrons with other atoms of elements, their valency is 0.
The noble gases make up approximately 1% by volume of the atmosphere, in which Argon alone constitutes 0.93% by volume of the atmosphere.
Argon can be obtained from liquid air by fractional distillation.
Natural gas deposits are the most important commercial and economical source of Helium.
The atomic radii increase on descending the group with an increase in the atomic number.
Due to their stable electronic configuration, these elements have very high ionisation potentials.
As these elements do not have a tendency to accept electrons, they have very high positive values of electron gain enthalpy.
Group 18: Physical and Chemical Properties
The elements of group eighteen are colourless, tasteless and odourless monoatomic gases.
Noble gases are sparingly soluble in water.
The magnitude of the van der waals forces increases on moving down the group with an increase in the polarisability of atoms.
The melting and boiling points of group eighteen elements are very low because of very weak van der Waals forces.
The inertness to chemical reactivity of noble gases is attributed to their stable electronic configuration, high ionisation enthalpy and higher positive electron gain enthalpy.
In 1962, Neil Bartlett first prepared a compound of xenon, called xenon hexafluoridoplatinate five.
Group 18: Xenon –Fluorine Compounds
Xe and F 2 in the ratio of 2:1, heated at 673 K at 1 bar pressure give XeF 2.
Xe and F 2 in the ratio of 1:5, heated at 873 K at 7 bar pressure give XeF 4.
Xe and F 2 in the ratio of 1:20, heated at 573 K at 60-70 bar pressure give XeF 6.
The structure of XeF 2 is linear owing to the presence of the 3 lone pairs of electrons on the central Xe atom.
The structure of XeF 4 is square planar owing to the presence of the 2 lone pairs of electrons on the central Xe atom.
The structure of XeF 6 is distorted octahedral owing to the presence of the single lone pair of electrons on the central Xe atom.
Xenon fluorides are white crystalline solids, and are strong oxidising and fluorinating agents.
The fluorides of Xe form complexes with covalent penta fluorides.
Group 18: Xenon –Oxygen Compounds
Xenon tetrafluoride and xenon hexafluoride undergo complete hydrolysis to give xenon trioxide as the main product.
The structure of xenon trioxide is pyramidal owing to the presence of a lone pair of electrons on the central xenon atom.
The structure of xenon oxytetrafluoride is square pyramidal owing to the presence of a lone pair of electrons on the central xenon atom.
The structure of xenon dioxydifluoride is see saw owing to the presence of a lone pair of electrons on the central xenon atom.
Xenon trioxide is a colourless, explosive and hygroscopic crystalline solid, while xenon oxytetrafluoride is a colourless liquid.
Group 18: Uses
“He” is used to fill weather balloons and airships because it is not inflammable and has a low density.
Liquid Helium is used as a cryogenic agent to perform several experiments at very low temperatures.
Liquid Helium is used to cool the superconducting magnets used in nuclear magnetic resonance spectrometers.
“He” is used to dilute O 2 in the cylinders carried by sea divers because it is only slightly soluble in blood even at high pressures.
“Ne” is used in discharge tubes and fluorescent bulbs that give the familiar red orange glow of “neon” signs.
“Ar” is used to provide inert atmosphere in laboratory apparatus and for metallurgical processes.
“Ra” is used in the treatment of cancer.