P block elements

Group 18 Elements Group 18 consists of six elements: helium, neon, argon, krypton, xenon and radon. _____________________________________________________________________________________________

Occurrence

All the noble gases except radon occur in the atmosphere. ARGON is the major constituent. HELIUM AND NEON are found in radioactive minerals e.g., pitchblende, monazite, cleveite. The main commercial source of HELIUM is

natural gas. Xenon and radon is

rare and Radon is obtained as a radioactive decay product of 226Ra. _____________________________________________________________________________________

Why are the elements of Group 18 known as noble gases? The elements present in Group 18 have their valence shell orbitals completely filled, Due to stable electronic configuration it is difficult to remove electrons or accept the electron and therefore these gases exhibit very high ionization enthalpy and have large positive values of electron gain enthalpy. Therefore, react with a few elements only under certain conditions ______________________________________________________________________________

Physical Properties All the noble gases are monatomic. They are colorless, odorless and tasteless. They are sparingly soluble in water. They have very low melting and boiling points because the only type of interatomic

weak dispersion forces. Helium has the lowest known substance. It has an unusual Property of

Interaction in these elements is boiling point (4.2 K) of any

diffusing

through most commonly used laboratory materials Such as rubber, glass or

plastics. _____________________________________________________________________________________

Chemical Properties In general, noble gases are least reactive. Their chemical inertness is attributed to the following reasons: (i) except helium (1s2) all have completely filled ns2np6 electronic configuration in their valence shell. (ii) They have high ionization enthalpy.

INVESTIGATING THE REACTIVITY OF NOBLE GASES Though noble gas elements have generally very low reactivity (due to their very high ionization and electron gain enthalpies), the heavier noble gases, Krypton and Xenon tend to form some compounds. (REASON: These heavier elements have more electron shells than the lighter ones. Hence, the outermost electrons experience a shielding effect from the inner electrons that makes them more easily ionized. This results in an ionization energy low enough to form stable compounds with the most electronegative elements like fluorine and oxygen.) Ever since their discovery, the reactivity of the noble gases has been investigated sporadically ever, but all early attempts to coerce them into compound formation were unsuccessful. Until the 1960s the only known compounds were the unstable diatomic species such as He2-and Ar2-,

In March 1962, Neil Bartlett, then at the University of British Columbia, observed the reaction of Xenon, a noble gas. Bartlett’s motivation for studying xenon was based on the observations that the highly oxidizing compound, PtF6, can oxidize O2, to give the red solid O2+PtF6–. As the ionisation energy of O2 to O+2 (1165 kJ mol−1) is nearly equal to the ionisation energy of Xe to Xe+ (1170 kJ mol−1), he tried the reaction of Xe with PtF6. This yielded a red crystalline product, xenon hexafluoroplatinate, Xe+[PtF6]−. The compounds of other noble gases are fewer. Only the difluoride (KrF2) and (RnF2) has been identified. No true compounds of Ar, Ne or He are yet known.

Xenon fluorides

XeF2, XeF4 and XeF6 are colorless crystalline solids and sublime readily at 298 K. They are powerful fluorinating agents.

The structures of the three xenon fluorides

can be deduced

from VSEPR. XeF2 is linear and XeF4 is square planar. XeF6 has 7 electron pairs (6 bonding pairs and 1 lone pair) and would, thus, have a distorted octahedral structure.

Preparation: Xenon forms three binary fluorides, XeF2, XeF4 and XeF6 by the direct reaction of elements under appropriate experimental conditions. Xe (g) + F2 (g)  XeF2(s) (excess) Xe (g) + 2F2 (g XeF4(s) (1:5 ratio) Xe (g) + 3F2 (g)  XeF6(s) (1:20 ratio) XeF6 can also be prepared by the interaction of XeF4 and O2F2. XeF4 + O2F2XeF6 + O2

Key points: Xenon fluorides are strong oxidizing agents Xenon fluorides react with strong Lewis acids They are readily hydrolyzed even by traces of WATER. For example, XeF2 is hydrolyzed to give Xe, HF and O2. 2XeF2 (s) + 2H2O (l)  2Xe (g) + 4 HF + O2 As with the interhalogens, the xenon fluorides react with strong Lewis acids to form xenon fluoride cations: An example is the reaction with FLUORIDE ION ACCEPTORS to form cationic species and fluoride ion donors to form fluoroanions. XeF2 + PF5  [XeF]+[PF6]– XeF4 + SbF5  [XeF3]+[SbF6]– XeF6 + MF  M+[XeF7]– (M = Na, K, Rb or Cs)

Xenon-oxygen compounds Preparation: Xenon oxides are endergonic compounds and cannot be prepared by direct interaction of the elements, so we need to look for an indirect method. The oxides and oxofluorides are prepared by the

hydrolysis of xenon fluorides:

Hydrolysis of XeF4 and XeF6 with water gives Xe03. 6XeF4 + 12 H2O  4Xe + 2Xe03 + 24 HF + 3 O2 XeF6 + 3 H2O XeO3 + 6 HF

Partial hydrolysis of XeF6 gives oxyfluorides, XeOF4 and XeO2F2.

XeF6 + H2O  XeOF4 + 2 HF XeF6 + 2 H2O  XeO2F2 + 4HF XeO3 is a colorless explosive solid and has a pyramidal molecular structure XeOF4 is a colorless volatile liquid and has a square pyramidal molecular structure XeOF5 is a pentagonal pyramid.

USES Helium

is a non-inflammable and light gas. Hence, it is used in filling

balloons for meteorological observations. It is also used in gas-cooled nuclear reactors. Liquid helium (b.p. 4.2 K) finds use as cryogenic agent for carrying out various experiments at low temperatures. It is used to produce and sustain powerful superconducting magnets which form an essential part of modern NMR spectrometers and Magnetic Resonance Imaging (MRI) systems for clinical diagnosis. It is used as a diluent for oxygen in modern diving apparatus because of its very low solubility in blood.

Neon is used in discharge tubes and fluorescent bulbs for advertisement display Argon is used mainly to provide an inert atmosphere in high temperature

Metallurgical processes arc welding and for filling electric bulbs. It is also used in the laboratory for handling substances that are air-sensitive.

Group 17 Elements

Fluorine, chlorine, bromine, iodine and astatine are members of Group 17. They are called

halogens, meaning: salt producers). There is a regular gradation in their physical and chemical properties. _____________________________________________________________________________________________

OCCURRENCE Generally, these occur as metal halides in rocks and sea water. Fluorine is present as insoluble fluorides (fluorspar

fluoroapatite 3Ca3 (PO4)2.CaF2)

CaF2, cryolite Na3AlF6 and

carnallite, KCl.MgCl2.6H2O. Iodine is present in sea weeds and Chile saltpetre (sodium iodate) Chlorine is seen as sodium chloride and

_____________________________________________________________________________________________ Halogen Electronegativity Electron Affinity Bond Enthalpy

Fluorine 4 328

Chlorine 3.2 349

Bromine 3 325

Iodine 2.6 295

158

242

192

151

The electronic configuration (ns2np5) shows that the elements are one electron short of the next noble gas. This is the reason for their maximum effective nuclear charge in the periods, which further renders them these properties:   

Smallest atomic radii in their respective periods Very high ionization enthalpy Maximum negative electron gain enthalpy

(NOTE: Electron gain enthalpy of the elements of the group becomes less negative down the group. However, the negative

electron gain enthalpy of fluorine is less

than that of chlorine. Due to small size of fluorine atom, there are strong interelectronic

repulsions in the relatively small 2p orbitals of fluorine and thus, the incoming electron does not experience much attraction. This anomaly is also reflected in the

enthalpies, where the same reason applies.)

X-X bond dissociation

Why is fluorine a better oxidizing agent than chlorine, in spite of its lower electron gain enthalpy? It is due to (i) low enthalpy of dissociation of F-F bond (ii) high hydration enthalpy of F–

____________________________________________________________________________________________________________________________________________________________

Oxidation states and trends in chemical reactivity __________________________________________________________________

All the halogens exhibit –1 oxidation state. Chlorine, bromine and iodine exhibit + 1, +3, + 5 and + 7 oxidation states also. (Fluorine doesn’t exhibit other oxidation states primarily because it is the most electronegative element and also because it can’t expand its octet due to lack of d orbitals.)

The higher oxidation states of chlorine, bromine and iodine are observed in their compounds with the small and highly electronegative fluorine and oxygen atoms. e.g., in interhalogens, oxides and oxoacids. F2 is the strongest oxidizing halogen and it oxidizes other halide ions in solution. Their highly

+ standard electrode (red.) potentials also illustrate this.

Their reactions with water also show this:  Fluorine oxidizes water to oxygen  Chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids.  Iodine is too weak an oxidizer to react with water. In fact, I– can be oxidized by oxygen in acidic medium which is just the reverse of the reaction observed with fluorine. ______________________________________________________________________________

Anomalous behavior of fluorine   

Ionization enthalpy Electronegativity, Enthalpy of bond dissociation and electrode potentials are all HIGHER

  

Ionic and covalent radii, Melting and boiling points Electron gain enthalpy are LOWER

This anomaly is due to:

   

Small size, highest Electronegativity, low F-F bond dissociation enthalpy, Non availability of d orbitals in valence shell.



Form exothermic compounds (due to the small and strong bond formed by it with other elements). Forms only one oxoacid: HOF (hypofluorous acid), while other halogens form a number of oxoacids.



_____________________________________________________________________________ HYDROGEN HALIDES HF is a liquid due to strong hydrogen bonding. Others are gases. They all react with hydrogen to give hydrogen halides The ACIDIC STRENGTH (Ka) of the acids is in the order: HF < HCl < HBr < HI. So, STABILITY is in the reverse order (ΔdissH and bond length). ______________________________________________________________________________ HALOGEN OXIDES Halogens form many oxides with oxygen but most of them are unstable. Stability of oxides formed by halogens, is in the order: I > Cl > Br. The higher oxides of halogens tend to be more stable than the lower ones.  

Fluorine forms two oxides OF2 and O2F2. Only OF2 is thermally stable. O2F2 is used to remove plutonium from spent nuclear fuel by oxidizing it to PuF6

 

Chlorine oxides, Cl2O, ClO2, Cl2O6 and Cl2O7 are explosively oxidizing. ClO2 is used as a bleaching agent in paper industry and in water treatment.



Bromine oxides, Br2O, BrO2, and BrO3 are the least stable halogen oxides.

 Iodine oxides, I2O4, I2O5, and I2O7 decompose on heating.  I2O5 is used in the estimation of carbon monoxide. ______________________________________________________________________________ METAL HALIDES The ionic character of the halides: MF >MCl > MBr > MI The halides in higher oxidation state will be more covalent than the one in lower oxidation state. For e.g., SnCl4, UF6 are more covalent than SnCl2, UF4. _____________________________________________________________________________________________________________________

CHLORINE Preparation 

By heating HCl on MnO2.

4NaCl + MnO2 + 4H2SO4 → MnCl2 + 4NaHSO4 + 2H2O + Cl2

 By the action of HCl on KMnO4.

2KMnO4 + 16HCl → 2KCl + 2MnCl2 + 8H2O + 5Cl2

Manufacture of chlorine  

Deacon’s process: oxidation of HCl in the presence of CuCl2 Electrolysis of brine: Chlorine is liberated at anode.

Properties

Greenish yellow gas Heavier than air Soluble in water Pungent smelling

Ammonia

With excess ammonia, chlorine gives N2 and NH4Cl 8NH3 + 3Cl2 → 6NH4Cl + N2 With excess chlorine, NCl3 (explosive) forms NH3 + 3Cl2 → NCl3 + 3HCl

Alkali

With cold and dilute alkalis: chloride and hypochlorite 2NaOH + Cl2 → NaCl + NaOCl + H2O With hot and concentrated alkalis: gives chloride and chlorate. 6 NaOH + 3Cl2 → 5NaCl + NaClO3 + 3H2O With dry slaked lime it gives bleaching powder. 2 Ca(OH)2 + 2Cl2 → Ca(OCl)2 + CaCl2 + 2H2O The composition of bleaching powder is Ca(OCl)2.CaCl2.Ca(OH)2.2H2O.

Chlorine water

On standing loses its yellow color due to the formation of HCl and HOCl. Bleaching properties is due to the release of nascent (O) as HOCl decomposes Oxidizing properties: (i) ferrous to ferric (ii) sulphite to sulphate, (iii) sulphur dioxide to sulphuric acid (iv) iodine to iodic acid (HIO3) (v) Colored substance → Colorless substance

USES  bleaching wood pulp (manufacture of paper)  bleaching cotton  extraction of gold and platinum  preparation of poisonous gases :

phosgene (COCl2), tear gas (CCl3NO2), mustard gas __________________________________________________________________

HYDROGEN CHLORIDE Preparation

NaCl + H2SO4 ⎯→ NaHSO4 + HCl NaHSO4 + NaCl ⎯→ Na2SO4 + HCl Aqua regia (3:1 of HCl and HNO3) dissolves gold, platinum. Except copper, silver, gold and those metals below hydrogen in the reactivity series, all other metals react with HCl Iron in HCl forms FeCl2 and not FeCl3 Hydrochloric acid decomposes salts of weaker acids NaHCO3 + HCl → NaCl + H2O + CO2 Na2SO3 + 2HCl → 2NaCl + H2O + SO2 USES: Manufacture of glucose (from corn starch) Extracting glue from bones and purifying bone black,

OXOACIDS Halic (I) acid (Hypohalous acid) HOF HOCl HOBr HOI

Halic (III) acid – (Halous acid) -----------HOCIO -----------------------

Halic (V) acid – (Halic acid) ------------HOCIO2 HOBrO2 HOIO2

Halic (VII) acid – (Perhalic acid) -----------HOCIO3 HOBrO3 HOIO3

INTERHALOGEN COMPOUNDS XX′, XX3 ′, XX5′ and XX’7

Preparation

 By direct combination or  By the action of halogen on lower interhalogens compounds. I2 + Cl2 → 2ICl (equimolar) I2 + 3Cl2 → 2ICl3 (excess) Cl2 + F2 ⎯→ 2ClF (equal volume) Cl2 + 3F2⎯→2ClF3 (excess) Br2 + 3F2 → 2BrF3 (diluted with water) Br2 + 5F2 → 2BrF5 (excess)

Properties   

Are diamagnetic in nature. They are volatile solids or liquids except CIF which is a gas. Interhalogens compounds are more reactive than halogens (except fluorine). (This is because X–X′ bond in interhalogens is weaker than X–X bond in halogens except F–F bond.)

Hydrolysis

These undergo hydrolysis giving halide ion derived from the smaller halogen and hypohalite ( when XX′), halite ( whenXX′3), halate (when XX′5) or perhalate (when XX′7) derived from the larger halogen. e.g. 2 XX' + H O → HX' + HOX

USES: These compounds can be used as non aqueous solvents. Interhalogen compounds are very useful fluorinating agents. ClF3 and BrF3 are used for the production of UF6 in the enrichment of 235U.

Group 16 Elements

Chalcogens: ore-forming

Oxygen, sulphur, selenium, tellurium and polonium constitute Group 16. _________________________________________________________________________________________

Occurrence Order of elements in the earth’s crust: O Si Al Fe Ca Na K Mg Most copper minerals contain either oxygen or sulphur Oxygen is the most abundant of all the elements on earth (46.6% by mass) Combined sulphur exists primarily as Sulphates:  gypsum CaSO4.2H2O,  Epsom salt MgSO4.7H2O,  baryte BaSO4

Sulphides:  galena PbS,  zinc blende ZnS,  Copper pyrites CuFeS2.

Selenium and tellurium are also found as metal selenides and tellurides in sulphide ores. Polonium occurs in nature as a decay product of thorium and uranium minerals. _________________________________________________________________________________________

Physical properties

The elements of Group16 have ns2np4 general electronic configuration. Elements of this group have lower ionization enthalpy values compared to those of Group15 or Group17 [reason: This is due to the fact that Group 15 elements have extra stable half-filled p orbitals electronic configurations.] Oxygen has the highest electronegativity value; the metallic character increases from oxygen to polonium. The size of oxygen atom is very small and has the least negative electron gain enthalpy.

Order: O H2Po Reducing property and this character increases from H2S to H2Te. Bond angle (°) H2O > H2S > H2Se > H2Te > H2Po a

1.8×10–16 1.3×10–7 1.3×10–4 2.3×10–3

Oxides of the type EO

2

and EO3

Both types of oxides are acidic in nature. Ozone (O3) and sulphur dioxide (SO2) are gases Selenium dioxide (SeO2) is solid. Reducing property of dioxide decreases from SO2 to TeO2; SO2 is reducing while TeO2 is an oxidising agent

Halides of the type, EX6, EX4 and EX2

The stability of the halides decreases in the order F– > Cl– > Br– > I Hexahalides Hexafluorides are the only stable halides. All are gaseous. SF6 is exceptionally stable for steric reasons. Tetrafluorides SF4 is a gas, SeF4 a liquid and TeF4 a solid.

sp3d hybridization and trigonal bipyramidal structure and see-saw geometry. Dihalides Selenium cannot form dihalides

sp3 hybridization and tetrahedral structure. Monohalides Dimeric in nature. Egs: S2F2, S2Cl2, S2Br2, Se2Cl2 and Se2Br2. Undergo disproportionation: 2Se2Cl2 SeCl4 + 3Se

Dioxygen Preparation

in the laboratory:  By heating oxygen rich salts such as chlorates, nitrates and permanganates.  By the thermal decomposition of the oxides of metals low in the electrochemical series and higher oxides of some metals. o 2Ag2O(s) 4 Ag(s) + O2(g) o 2HgO(s) 2 Hg(l) + O2(g) o 2PbO2(s) 2 PbO(s) + O2(g)   

Decomposition of Hydrogen peroxide by catalysts such as finely divided metals and manganese dioxide. 2H2O2(aq) 2H2O(1) + O2(g) Electrolysis of water Fractional distillation

Properties

Dioxygen is a colourless and odourless gas. Its solubility in water is to the extent of 3.08 cm3 in 100 cm3 water at 293 K which is just sufficient for the vital support of marine and aquatic life. It liquefies at 90 K and freezes at 55 K. It has three stable isotopes: 16O, 17O and 18O. Molecular oxygen, O2 is unique in being paramagnetic (in spite of having even no. of e-) bond dissociation enthalpy of oxgyen-oxygen double bond is high (493.4 kJ mol–1). Oxides can be simple (e.g., MgO, Al2O3 ) or mixed (Pb3O4, Fe3O4). acidic oxide (e.g., SO2, Cl2O7, CO2, N2O5 ). non-metal oxides are acidic but oxides of some metals in high oxidation state also have acidic character (e.g., Mn2O7, CrO3, V2O5). basic oxides (e.g., Na2O, CaO, BaO). In general, metallic oxides are basic. amphoteric oxides or neutral oxides are CO, Al2O3, NO and N2O.

Uses  Oxyacetylene welding,  Manufacture of steel.  Rocket fuel: hydrazine in liquid oxygen

Ozone -

An allotrope of oxygen

Preparation Silent electric discharge of oxygen [the formation of ozone from oxygen is an endothermic process]

Properties

O-O bond lengths are identical (128 pm), bond angle is 117o. Pure ozone is a pale blue gas, dark blue liquid and violet-black solid. Ozone has a characteristic smell Above 100 ppm, breathing difficulty, headache and nausea results. Ozone is thermodynamically unstable with respect to oxygen (explosive). As it liberates atoms of nascent oxygen, it acts as a powerful oxidizing agent Lead sulphide to lead sulphate Iodide ions to iodine

Estimating

O3 gas:

Iodine liberated, when ozone reacts with potassium iodide solution can be titrated against a standard solution of sodium thiosulphate.

Depletion of ozone: Nitric oxide from the exhaust systems of supersonic jets Freons, which are used in aerosol sprays and as refrigerants

Uses     

Germicide, disinfectant For sterilizing water. bleaching oils, ivory, flour, starch, Oxidizing agent in the manufacture of potassium permanganate.

Sulphur Allotropic Forms The allotropes exist in equilibrium @ 369K (transition temperature)

Rhombic sulphur (α-sulphur)   

Monoclinic sulphur (β -sulphur)  

Yellow in color The stable form at rtp Prepared by by evaporating the solution of roll sulphur in CS2 Both rhombic and monoclinic sulphur have

It is soluble in CS2. prepared by melting rhombic sulphur and cooling, till crust is formed as needle shaped

S8 molecules. Shape-puckered crown.

Both are insoluble in water but soluble in organic solvents

S6, has chair form

Cyclo-

At high temperatures, sulphur exists as S2 and is paramagnetic like O2.

[Two unpaired electrons in the antibonding ∏* orbitals ]

Sulphur Dioxide Preparation   

Properties

Burn sulphur in air Treat sulphite with acid (protons from acid are oxidized to water) Roasting sulphide ores

  

Colorless gas Pungent smell Highly soluble in water

It reacts readily with sodium hydroxide solution, forming sodium sulphite, and then sodium hydrogen sulphite. Sulphur dioxide reacts with chlorine in the presence of charcoal (which acts as a catalyst) to give sulphuryl chloride, SO2Cl2.

When moist, sulphur dioxide behaves as a reducing agent. Converts iron (III) ions to iron (II) ions Decolorizes KMnO4 solution SO2

Uses:    

refining petroleum and sugar bleaching wool and silk anti-chlor Disinfectant and preservative.

Detection

Oxoacids of Sulphur Manufacture

Manufactured by the Contact Process (low temperature and high pressure) 2 bar /720 K. 

Burning of sulphur or sulphide ores in air to generate SO2.



conversion of SO2 to SO3 in the presence of (V2O5), Absorption of SO3 in H2SO4 to give Oleum (H2S2O7).



Properties

The acid forms two series of salts: normal sulphates and acid sulphates (hydrogen sulphate). Concentrated sulphuric acid is a strong dehydrating agent. moderately strong oxidizing agent.

Uses:      

Manufacture of fertilizers (e.g., ammonium sulphate, superphosphate). petroleum refining manufacture of pigments, paints and dyestuff detergent industry cleansing metals storage batteries



nitrocellulose products

Group 15 Elements

Group 15 includes nitrogen, phosphorus, arsenic, antimony and bismuth. N and P are non-metals, Ar and Sb are metalloids and Bi is a typical metal. _________________________________________________________________________________________

Occurence Nitrogen occurs as NaNO3 (Chile saltpetre) and KNO3 (Indian saltpetre) Phosphorus occurs in apatite minerals: Ca9(PO4)6. CaX2 (e.g., fluoroapatite Ca9 (PO4)6. CaF2) Arsenic, antimony and bismuth are found mainly as sulphide minerals.

Property

The valence shell e.c. is ns2np3. The p orbitals are half-filled, giving it extra stability. Except nitrogen, all the elements show allotropy. The melting point increases up to arsenic and then decreases up to bismuth

Oxidation states and trends in chemical reactivity The common oxidation states are –3, +3 and +5. –3 Oxidation state becomes less common down the group due to decreased electronegativity. Bismuth doesn’t have compounds with –3 oxidation state.

The stability of +5 oxidation state decreases and that of +3 state increases (due to inert pair effect) down the group. In the case of nitrogen, all oxidation states from +1 to +4 tend to disproportionate in acid solution. 3HNO2 HNO3 + H2O + 2NO In case of phosphorus nearly all intermediate oxidation states disproportionate into +5 and –3 both in alkali and acid.

Anomalous properties of nitrogen  

Nitrogen has unique ability to form p∏-p∏ multiple bonds with itself and others. Thus, nitrogen exists as a diatomic molecule with a triple bond 941.4 kJ mol–1. The single N–N bond is weaker than the single P–P bond because of high interelectronic repulsion of the non-bonding electrons, owing to the small bond length.

Due to the absence of d orbitals in its valence shell:  covalency is restricted to 4  Cannot form d∏ –p∏ bond as the heavier elements can e.g., R3P = O or R3P = CH2 or d∏ –d∏ bond as in P(C2H5)3 and As(C6H5)3.

Hydrides

The stability of hydrides decreases from NH3 to BiH3 (bond dissociation enthalpy) So, the reducing character of the hydrides increases. BiH3 is the strongest reducing agent. Basicity also decreases in the order NH3 > PH3 > AsH3 > SbH3 > BiH3.

Oxides:

E2O3 and E2O5.

The oxide in the higher oxidation state is more acidic. N2O3 and P2O3 are acidic; As2O3 and Sb2O3 are amphoteric;

Halides:

BiO3 is basic

EX3 and EX5.

Pentahalides are more covalent than trihalides. Only stable trihalide of nitrogen is NF3 Only ionic halide is BiF3

Dinitrogen Preparation    

Fractional distillation of air. Treating an aqueous solution of ammonium chloride with sodium nitrite. NH4CI+ NaNO2 N2+ 2H2O+ NaCl Thermal decomposition of ammonium dichromate. (NH4)2Cr2O7 ⎯⎯ N2 + 4H2O + Cr2O3 Very pure nitrogen -thermal decomposition of sodium or barium azide. Ba(N3)2 Ba + 3N2

Properties Forms ionic nitrides with metals and covalent nitrides with non metals.

Ammonia Preparation  

Treating ammonium salts with NaOH/CaO Haber’s process: 200 atm ~ 700 K Catalyst such as iron oxide with small amounts of K2O and Al2O3

Properties

Its aqueous solution is weakly basic due to the formation of OH– ions. It IS a Lewis base. It forms linkage with metal ions in complex compounds Used in the detection of metal ions such as Cu2+, Ag+: Cu2+ (blue) + 4 NH3 [Cu(NH3)4]2 (deep blue)

Oxides of Nitrogen

Why does NO2 dimerise ? NO2 contains odd number of valence electrons. It behaves as a typical odd molecule. On dimerisation, it is converted to stable N2O4 molecule with even number of electrons.

Nitric Acid Preparation  heating Nitrate and concentrated H2SO4  Ostwald’s process: Platinum-Rhodium catalyst, NH3NONO2HNO3 Properties Planar molecule Concentrated nitric acid is a strong oxidising agent: I2 + 10 HNO3  2 HIO3 + 10 NO2 + 4 H2O C + 4 HNO3  CO2 + 2 H2O + 4 NO2 S8 + 48 HNO3 (conc.)  8 H2SO4 + 48 NO2 + 16 H2O P4 + 20 HNO3 (conc.)  4 H3PO4 + 20 NO2 + 4 H2O The products of oxidation depend upon the concentration of the acid, temperature and the nature of the material undergoing oxidation.

Copper reacts with dilute nitric acid to give NO and with concentrated acid to give NO2 Zinc reacts with dilute nitric acid to give N2O and with concentrated acid to give NO2.

Cr and Al do not react due to the formation of a passive oxide film on the surface.

Brown Ring Test: Based on the ability of Fe2+ to reduce nitrates to nitric oxide, which reacts with Fe2+ to form a brown colored complex.

Uses:     

fertilizers and explosives nitroglycerin, trinitrotoluene

pickling of stainless steel, etching of metals oxidizer in rocket fuels

Phosphorus

Allotropic Forms-white, red and black

White phosphorus

Red phosphorus

Translucent white waxy solid It is poisonous Insoluble in water but soluble in CS2 Glows in dark It gives PH3 and Na2HPO2 with NaOH Less stable/more reactive because of angular strain in P4 molecule 60° Catches fire in air to give dense

Possesses iron grey lustre Nonpoisonous Insoluble in water as well as in CS2 It does not glow in the dark. When red phosphorus is heated, black phosphorus is formed. Red phosphorus is much less reactive than white phosphorus By heating white phosphorus @ 573K Polymeric-P4 tetrahedra linked together

white fumes of

P4O10.

It consists of discrete P4 molecules

Black phosphorus

has two forms α-black phosphorus and β-black phosphorus.

red phosphorus ---heat@803K  α-black phosphorus. white phosphorus –heat@473K  β-Black phosphorus

Phosphine Preparation  

reaction of calcium phosphide with water or dilute HCl.

heating white phosphorus with concentrated NaOH solution

It is non inflammable when pure, but becomes inflammable owing to the presence of

P2H4 or P4 vapors.

To purify it from the impurities, it is absorbed in HI to form (PH4I) treated with KOH

Properties     

Colorless gas with a rotten fish smell and is highly poisonous. It explodes in contact with oxidizing agents It is slightly soluble in water. PH3 in water decomposes in presence of light giving red phosphorus and H2. With copper sulphate or mercuric chloride solution, their phosphides form Phosphine is weakly basic like ammonia

Uses:  Holmes’s signals. calcium carbide CaC2 and calcium phosphide Ca3P2  It is also used in smoke screens.

Phosphorus Halides Phosphorus Trichloride Preparation 

passing dry chlorine over heated white phosphorus.

P4 + 6Cl2  4PCl3 

action of thionyl chloride with white phosphorus.

P4 + 8SOCl2  4PCl3 + 4SO2 + 2S2Cl2 Properties  

colourless PCl3 fume in moisture as it hydrolyses giving fumes of HCl and H3PO3

Phosphorus Pentachloride Preparation 

passing excess dry chlorine over heated white phosphorus.

P4 + 10Cl2  4PCl5 

action of SO2Cl2 with white phosphorus.

P4 + 10SO2Cl2  4PCl5 + 10SO2 Properties       

yellowish white powder hydrolyses to POCl3 and then to H3PO4 When heated, it gives PCl3 and Cl2 Finely divided metals on heating with PCl5 give corresponding chlorides In gaseous and liquid phases, it has a trigonal bipyramidal structure. The two axial bonds are longer than equatorial bonds. (Repulsion) In the solid state it exists as an ionic solid, [PCl4]+[PCl6]–, with tetrahedral cation, [PCl4]+ and octahedral anion, [PCl6]–

Oxoacids of Phosphorus

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The compositions of the oxoacids are in terms of loss or gain of H2O or O-atom In oxoacids phosphorus is tetrahedrally surrounded by other atoms. All these acids contain one P=O and at least one P–OH bond. The P–P and P–H (can’t be found together. Acids in +3 oxidation tend to disproportionate to higher and lower oxidation states. For example, orthophophorous acid (or phosphorous acid) on heating disproportionate to give orthophosphoric acid (or phosphoric acid) and phosphine. The acids which contain P–H bond have strong reducing properties. Compare Basicity of these acids