Chapter 8: p-Block (Halogen & Inert Gas Family)

Chapter 8

p-block elements

(halogens family and noble gases)

Introduction :

Group 13 to 18 of the periodic table of elements constitute the p–block. The p–block contains metals, metalloids as well as non–metals.

The  p–block elements have general valence shell electronic configuration ns² np^1–6.

The first member of each group from 13–17 of the p–block elements differ in many respects from the other members of their respective groups because of small size, high electronegativity and absence of d–orbitals.

The first member of a group also has greater ability to form pp–pp multiple bonds to itself (e.g. C=C, C=C, N º N) and to element of second row (e.g C=O, C=N, CºN, N=O) compared to the other members of the same group.

The highest oxidation of p–block element is equal to the group number minus 10. Moving down the group, the oxidation state two less than the highest group oxidation state becomes more stable in groups 13 to 16 due to inert pair effect (reluctance of s-subshell electrons to participate in chemical bonding)

Group 17 Elements : The Halogen Family

Electronic Configuration : All these elements have seven electrons in their outermost shell (ns² np⁵) which is one electron short of the next noble gas.

Atomic and Ionic Radii : The halogens have the smallest atomic radii in their respective periods due to maximum effective nuclear charge . Atomic and ionic radii increase from fluorine to iodine due to increasing number of quantum shells.

Ionisation Enthalpy : They have little tendency to lose electron. Thus they have very high ionisation enthalpy. Due to increase in atomic size, ionisation enthalpy decreases down the group.

Electron Gain Enthalpy : Halogen have maximum negative electron gain enthalpy in the corresponding period. This is due to the fact that the atoms of these elements have only one electron less than stable noble gas configurations. 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. It is due to small size of fluorine atom. As a result, there are strong interelectronic repulsions in the relatively small 2p orbitals of fluorine and thus, the extra electron (incoming) does not experience much attraction.

Electronegativity : They have very high electronegativity. The electronegativity decreases down the group. Fluorine is the most electronegative element in the periodic table

Physical Properties : Fluorine and chlorine are  gases, bromine is a liquid whereas iodine is a solid. Their melting and boiling points steadily increase with atomic number. All halogens are coloured. This is due to absorption of radiations in visible region which results in the excitation of outer electrons to higher energy level. By absorbing different quanta of radiation, they display different colours. For example, F₂, has yellow, Cl₂, greenish yellow, Br₂, red and I₂, violet colour. Fluorine and chlorine react with water. Bromine and iodine are only sparingly soluble in water. But are soluble in organic solvents such as chloroform, carbon tetrachloride, carbon disulphide and hydrocarbons to give coloured solutions. Except the smaller enthalpy of dissociation of F₂ compared to that of Cl₂. The X-X bond disassociation enthalpies from chlorine onwards show the expected trend : Cl – Cl > Br – Br > F – F > I – I. The reason for the smaller enthalpy of dissociation of F₂ is the relatively larger electrons-electrons repulsion among the lone pairs in F₂ molecule where they are much closer to each other than in case of Cl₂.

Atomic & physical properties

 

Chemical Properties

Oxidation states and trends in chemical reactivity

All the halogens exhibit –1 oxidation state. However, chlorine, bromine and iodine exhibit + 1, + 3, + 5 and + 7 oxidation states also. The higher oxidation states of chlorine, bromine and iodine are realised mainly when the halogens are in combination with the small and highly electronegative fluorine and oxygen atoms e.g., in interhalogens, oxides and oxoacids.

The fluorine atom has no d orbitals in its valence shell and therefore cannot expand its octet. Being the most electronegative, it exhibits only – 1 oxidation state.

All the halogens are highly reactive. They react with metals and non-metals to form halides. The reactivity of the halogens decreases down the group.

The ready acceptance of an electron is the reason for the strong oxidising nature of halogens. F₂ is the strongest oxidising halogen and it oxidises other halide ions in solution or even in the  solid phase. The decreasing oxidising ability of the halogens in aqueous solution down the group is evident from their standard electrode potentials. Fluorine oxidises water to  oxygen whereas chlorine and bromine react with water to form corresponding hydrohalic and hypohalous acids. The reactions of iodine with water is non- spontaneous . I^– can be oxidised by oxygen in acidic medium; just the reverse of the reaction observed with fluorine.

2F₂(g) + 2H₂O(l)  ® 4H+ (aq) + 4F– (aq) + O₂(g)

 X₂(g) + H₂O (l) ® HX(aq) + HOX (aq)   ;          (where X = Cl or Br)

4I^– (aq) + 4H+ (aq) + O₂(g) ® 2 I₂  (s) + 2H₂O (l)

Standard Reduction Potential (SRP) 

X₂ + 2e^– ® 2X^– 

F2 + 2e^– ® 2F^–           [E° =  + 2.87 V ]

Cl₂ + 2e^– ® 2Cl^–        [E° =  + 1.36 V]

Br2 + 2e ^– ® 2Br ^–      [E° =  + 1.09 V ] 

I₂ + 2e^– ® 2I^–         [E° =  + 0.54 V]

More the value of the SRP, more powerful  is the oxidising agent. Hence the order of oxidising power is
F₂ > Cl₂ > Br₂ > I₂ 

Since SRP is the highest for F₂ (among all elements of periodic table), it is a strongest oxidising agent.

Hydration energy of X– 

Smaller the ion, higher is the hydration energy.      

F ^–        Cl ^–         Br ^–          I^–

515       381         347          305    in kJ/mol

Anomalous behaviour of fluorine

The anomalous behaviour of fluorine is due to its small size, highest electronegativity, low F- F bond dissociation enthalpy, and non availability of d orbitals in valence shell. Most of the reactions  of fluorine are exothermic (due to the small and strong bond formed by it with other elements). It forms only one oxoacid while other halogens form a number of oxoacids. Hydrogen fluoride is liquid (b.p. 293 K) due  to strong hydrogen bonding. Other hydrogen halides are gases.

(i) Reactivity towards hydrogen : They all react with hydrogen to give hydrogen halides but affinity for hydrogen decreases from fluorine  to iodine with increasing atomic number. They dissolve in water to form hydrohalic acids.  The acidic strength of these acids increases in the order : HF < HCl < HBr < HI. The stability of these halides decreases down the group due  to decrease in bond (H–X) dissociation enthalpy in the order :

H – F > H – Cl > H –Br > H – I .

(ii) Reactivity towards oxygen : Halogens form many oxides with oxygen but most of them are unstable. Fluorine forms two oxides OF₂ and O₂F₂. However, only OF₂ is the thermally stable at 298 K. These oxide are essentially oxygen fluorides because of the higher electronegativity of flurorine than oxygen . Both are strong fluorinating agents. O₂F₂ oxidises plutonium to PuF₆ and the reaction is used in removing plutonium as PuF6 from spent nuclear fuel.

Chlorine, bromine and iodine form oxides in which the oxidation states of these halogen vary from + 1 to + 7. A combination of kinetic and thermodynamic  factors lead to the generally decreasing  order of stability of oxides formed by halogens, I > Cl > Br. The higher oxides of halogens tend to be more stable than the lower ones.

Chlorine oxides, Cl₂O, ClO₂, Cl₂O₆ and Cl₂O₇ are highly reactive oxidising agents and tend to explode. ClO₂ is used as a bleaching agent for paper pulp and textiles and in water treatment.

The bromine oxides, Br₂O, BrO₂, BrO₃ are the least stable halogen oxides and exist only at low temperature. They are very powerful oxidising agents.

The iodine oxides, I₂O₄, I₂O₅, I₂O₇ are insoluble solids and decompose on heating. I₂O₅ is very good oxidising agent and is  used in the estimation of carbon monoxide.

(iii)  Reactivity towards metals : Halogen react with metals to form metal halides. For e.g.,  bromine reacts with magnesium to give magnesium bromide.

(iv) Reactivity of halogen towards other halogens : Halogens combine amongst themselves to form a number of compounds known as interhalogen of the types AB, AB₃, AB₅ and AB₇ where A is a larger size halogen and B is smaller size halogen.

CHLORINE (Cl₂)

Preparation :

(i) By heating MnO₂ with concentrated hydrocloric acid.

MnO₂ + 4HCl ® MnCl₂ + Cl₂ + 2H₂O

(ii) By heating chloride with concentrated H₂SO₄ in presence of MnO₂.

4H⁺  + MnO₂ + 2X^– ® X₂ + Mn⁺² + 2H₂O

(iii) 2KMnO₄ + 16 HCl ® 2 KCl + 2 MnCl₂ + 5 Cl₂ + 8 H₂O

(iv) Manufacture of chlorine :

(a) Deacon’s process : By oxidation of hydrogen chloride gas by  atmospheric oxygen in the presence of CuCl₂ (catalyst) at 723 K.

4 HCl + O₂  2 Cl₂ + 2 H₂O

(b) Electrolytic process : Chlorine is obtained by the electrolysis of brine (concentrated NaCl solution). Chlorine is liberated at anode. It is obtained as a by–product in many chemical industries e.g.; in mfg. of sodium hydroxide.

NaX (aq) ® Na+ (aq) + X^– (aq)

Anode : 2X^– ® X₂ + 2e^– 

Properties :

(i) It is a greenish–yellow gas with pungent and suffocating odour. It is about 2–5 times heavier than air. It can be  liquefied into greenish–yellow liquid which boils at 239 K. It is soluble in water.

(ii) At low temperature it forms a hydrate with water having formula Cl₂ . 8H₂O which is infact a clathrate compound.

(iii) Reaction with metals :

2Al + 3Cl₂ ® 2AlCl₃   

2Fe + 3Cl₂ ® 2FeCl₃

Reaction with non-metals :

P₄ + 6Cl₂ ® 4PCl₃ 

S₈ + 4Cl₂ ® 4S₂Cl₂

(iv) Affinity for hydrogen : It reacts with compounds containing hydrogen and form HCl.

H₂ + Cl₂ ® 2HCl  

H₂S + Cl₂ ® 2HCl + S

(v)  Reaction with NaOH :

(a) 2 NaOH ( cold & dilute) + Cl₂ ® NaCl + NaClO + H₂O

(b) 6 NaOH (hot & concentrated) + 3 Cl₂ ® 5 NaCl + NaClO₃ + 3 H₂O

(vi) Reaction with dry slaked lime, Ca(OH)₂ : To give bleaching powder.

2 Ca(OH)₂ + 2 Cl₂ ® Ca(OCl)₂ + CaCl₂ + 2 H₂O

The composition of bleaching powder is Ca(OCl)₂.CaCl₂.Ca(OH)₂ .2H₂O

(vii)   Oxidising & bleaching properties :

Chlorine dissolves in water (Cl₂ water is yellow) giving HCl (colourless) and HOCl (colourless). Hypochlorous acid (HOCl) so formed, gives nascent oxygen which is responsible for oxidising and bleaching properties of chlorine.

(a) It oxidises ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid and iodine to iodic acid.

2 FeSO₄ + H₂SO₄ + Cl₂ ® Fe₂(SO₄)₃ + 2 HCl

Na₂SO₃ + Cl₂ + H₂O ® Na₂SO₄ +  2 HCl

SO₂ + 2 H₂O + Cl₂ ® H₂SO₄ + 2 HCl

I₂ + 6 H₂O + 5 Cl₂ ® 2 HIO₃ + 10 HCl

(b) It is a powerful bleaching agent ; bleaching action is due to oxidation.

Cl2 + H₂O ® 2 HCl + O

Coloured substance + O ® Colourless substance

It bleaches vegetable or organic matter in the presence of moisture. Bleaching effect of chlorine is permanent.

Note :

The bleaching action of SO₂ is temporary because it takes place through reduction.

SO₂ + 2 H₂O ® H₂ SO₄  + 2 H

SO₃^2– + Coloured material  SO₄^2– + Reduced colourless material.

Reduced Colourless material   Coloured material.

Uses :  Cl₂  is used

1. for bleaching wood pulp (required for the manufacture of paper and rayon), bleaching cotton and textiles,

2. in the manufacture of dyes, drugs and organic compounds such as CCl₄, CHCl₃, DDT, refrigerants, etc.

3. in the extraction of gold and platinum.

4. in sterilising drinking water and

5. preparation of poisonous gases such as phosgene (COCl₂), tear gas (CCl₃NO₂), mustard gas (ClCH₂CH₂SCH₂CH₂Cl).

HYDROGEN CHLORIDE (HCl)

Preparation :

By heating a halide with concentrated acid (Laboratory method) :

NaCl + H₂SO₄  NaHSO₄ + HCl 

NaHSO₄ + NaCl  Na₂SO₄ + HCl

This method is called as salt cake method as it involves the formation of  NaHSO₄ (salt cake).

HCl cannot be dried over P₂O₅ (P₄O₁₀) or quick lime since they react with gas chemically.

CaO + 2HCl  ® CaCl₂  + H₂O  

P₄O₁₀  + 3HCl ® POCl₃  + 3HPO₃ 

HCl is, hence dried by passing through concentrated H₂SO₄ .

Properties :

(i) It is a colourless, pungent smelling gases with acidic tastes.

(ii) It is easily liqueflied to a colourless liquid (b.p. 189 K) and freezes to a white crystalline solid (f.p. 159 K).

(iii)    It is quite soluble in water.

HCl ionises as below : HCl(g) + H₂O (l) ® H₃O+ (aq) + Cl^– (aq) ;  Ka = 107

It aqueous solution is called hydrochloric acid. High value of dissociation constant (Ka) indicates that it is a strong acid in water.

When three parts of concentrated HCl and one part of concentrated HNO₃ are mixed, aqua regia is formed which is used for dissolving noble metals, e.g., gold, platinum.

Au + 4 H⁺ + NO3^– + 4 Cl^– ® [AuCl₄]^– + NO + 2 H₂O

3 Pt + 16 H⁺ + 4 NO^3­– + 18 Cl^– 3 [PtCl₆]^2– + 4 NO + 8 H₂O

It reacts with ammonia forming white fumes of NH₄Cl

NH₃ + HCl ® NH₄Cl

(iv)     It decomposes salt of weaker acids.

Na₂CO₃ + 2HCl ® 2NaCl + H₂O + CO₂

NaHCO₃ + HCl ® NaCl + H₂O + CO₂

Na₂SO₃ + 2HCl  ® 2NaCl + H₂O+ SO₂

Uses :

1. HCl is used in preparation of Cl₂, chlorides, aqua regia, glucose,  (from corn starch),

2. It is used in medicines, laboratory as reagents, cleaning metal surfaces before soldering or electroplating.

3. It is used for extracting glue from bones and purifying bone black.

OXY-ACIDS OF  Halogens

Fluorine forms only one oxoacid, HOF due to high electronegativity and small size. Other halogens form a number of oxoacids which are stable only in aqueous solutions or in the form of their salts. They can not be isolated in pure form.

Some important order

(a) Acid strength

(i) HI > HBr > HCl > HF

(ii) HOCl > HOBr > HOI  

(iii) HClO₄ > HClO₃ >HClO₂ > HClO

(b) Oxidising powder

(i) F₂ > Cl₂ > Br₂ > I₂

(ii) BrO₄^– > IO₄^– > ClO4^–  (According to electrode potential)

(c) Order of disproportionations

3 XO^–  ® 2X^– + XO3^– (hypohalite ion) ;

IO^– > BrO^– > ClO^–

Interhalogen compounds

We know that halogen atoms have different electronegativity. Due to this difference in electronegativity the halogen atoms combine with each other and give rise to the formation of binary covalent compounds, which are called interhalogen compounds. These are of four types.

Preparation :

Cl₂ + F₂ (equal volumes)  2ClF

Cl₂ + 3F₂ (excess)  2ClF₃ 

  I₂ + Cl₂ ® 2ICl ;

 (equimolar)

Diluted with water :   

Br₂ (g) + 3F₂ ® 2BrF₃

F₂ is diluted with N₂ :

I₂ + 3F₂  2IF₃

F₂ is taken in freon :

Br₂ + 5F₂ (excess)  ® 2BrF₅

  Properties :

(i)  These compounds may be gases, liquids or solids.

Gases : ClF, BrF, ClF₃ , IF₇

Liquids : BrF₃, BrF₅ 

Solids : ICl, IBr, IF₃, ICl₃.

(ii)  All interhalogens are covalent molecules and are diamagnetic in nature since all the valence electrons present as bonding or non-bonding electrons are paired.

(iii) The boiling points increases with the increase in the electronegativity difference between A and B atoms.

(iv) Interhalogen compounds are more reactive than the parent halogens but less reactive than F₂.

(v) Hydrolysis : All these undergo hydrolysis giving halide ion derived from the smaller halogen and a hypohalite (when AB), halite (when AB₃), halate (when AB₅), and perhalate (when AB₇) anion derived from the larger halogen.

AB + H₂O ® HB + HOA

BrCl + H₂O ® HCl + HOBr 

ICl + H₂O ® HCl + HIO₂

 Uses :

(i) These compounds can be used as non aqueous solvents.

(ii) Interhalogen compounds are very useful fluorinating agents.

(iii) ClF₃ and BrF₃ are used for the production of UF₆ in the enrichment of 235U.

 U(s) + 3 ClF₃ (l) ® UF₆ (g) + 3 ClF (g)

Group 18 Elements : The Zero Group Family

Group 18 consists of six elements: helium, neon, argon, krypton , xenon and radon . All these are gases and  chemically unreactive. They form very few compounds . Because of this they are termed noble gases.

Occurrence : All the noble gases except radon occur in the atmosphere. Their atmospheric abundance in dry air is ~ 1% by volume of which argon is the major constituent. Helium and sometimes neon are found in minerals of radioactive origin e.g., pitchblende, monazite, cleveite. The main commercial source of helium is natural gas. Xenon and radon are the rarest elements of the group. Radon is obtained as a decay product of  ²²⁶Ra.

Most abundant element in air is Ar. Order of abundance in the air is Ar > Ne > Kr > He > Xe.

Electronic Configuration : All noble gases have general electronic configuration ns2np6 except helium which has 1s² . Many of the properties of noble gases including their inactive nature are ascribed to their closed shell structures.

Ionisation Enthalpy : Due to stable  electronic configuration these gases exhibit very high ionisation enthalpy . However, it decreases down the group with increases in atomic size.

Atomic Radii : Atomic radii increase down the group with increase in atomic number.

Electron Gain Enthalpy : Since noble gases have stable electronic configurations, they have no tendency to accept the electron and therefore, have larger positive values of electron gain enthalpy.

Physical properties : All the noble gases are mono-atomic. They are colourless, and tasteless. They are sparingly soluble in water. They have very low  melting and boiling points because the only  type of interatomic interaction in these elements is weak dispersion forces,. Helium has the lowest boiling point (4.2K) of any known substance. It  has a unusual property of diffusing through most commonly used laboratory materials such as rubber, glass or plastics.

Atomic & physical properties

Chemical Properties :

 In general, noble gases are least reactive. Their inertness to chemical reactivity is attributed to the following reasons :

(i) The noble gases except helium (1s²) have completely filled ns² np⁶ electronic configuration in their valence shell.

(ii)  They have high ionisation enthalpy and more positive electron gain enthalpy.

The reactivity of noble gases has been investigated occasionally  ever since their discovery, but all attempt to force them to react to form the compounds were unsuccessful for quite  a few years. In March 1962, Neil Bartlett, then at the University of British Columbia, observed the reaction of a noble  gas. First , he prepared a red compound which is formulated as O₂₊ PtF₆^–. He , then realised that  the first ionisation enthalpy of molecular oxygen (1175 kJ mol ^–1) was almost identical with that xenon (1170 kJ mol^–1). He made efforts to prepare same type of compound with Xe⁺ PtF₆^–  by mixing Pt F₆ and Xenon. After this discovery, a number of xenon  compounds mainly with most electronegative elements like fluorine and oxygen, have been synthesised.

The compounds of krypton are fewer. Only the difluoride (KrF₂) has been studied in detail. Compounds of radon have not been isolated but only identified (e.g., RnF₂) by radiotracer technique. No true compounds of Ar, Ne or He are yet known .

COMPOUNDS OF XENON

Xenon–FLUORINE compounds :

All fluorides of xenon are white solids. They are volatile, readily subliming at room temperature   (298 K)

They can be stored indefinitely in nickel or monel metal containers.

Xenon difluoride

 Xe + F₂   (2 : 1 mixture)      XeF₂ .

It is soluble in water, giving solution 0.15 M at 0ºC. The solution which have pungent smell due to XeF₂ are powerful oxidizing agents.

XeF₂(aq) + 2H⁺ + 2e^–  ® Xe + 2HF(aq) ; Eº = + 2.64 V

 XeF₂ + H₂ ®2 HF + Xe

Hydrolysis is slow in water due to dissolution of XeF₂ in HF formed, but is rapid in basic solution.

2Xe F₂ + 2H₂O ®2 Xe + 4HF + O₂ 

XeF₂ + 2OH^– ® Xe + O₂ + 2F^– + H₂O

It acts as fluoride ion donor  XeF₂ (Lewis base) + PF₅ (Lewis acid)  [XeF]⁺ [PF6]^– 

Xenon tetrafluoride

Xe + 2F₂ (1 : 5 mixture)  XeF₄ 

XeF₄ reacts violently with water, giving dangerously explosive XeO₃.

6XeF₄ + 12H₂O  ® 4Xe + 2 XeO₃ + 24 HF + 3O₂ 

XeF₄ can act as F– acceptor as well as F– donors and thus form anionic species in reactions such as :

XeF₄ + NaF ® Na⁺ XeF₅^–

XeF₄ + SbF₅ ® [XeF₃]+ [SbF₆]^– 

Xenon hexafluoride

Xe + 3F₂ (1 : 20 mixture)   XeF₆                

XeF₄ + O₂ F₂ ®  XeF₆ + O₂ 

XeF₆ reacts violently with water but slow hydrolysis with atmospheric moisture giving XeO₃.

XeF₆ + 6 H₂O  ® XeO₃ + 6HF (complete hydrolysis)

With small quantities of water, partial hydrolysis takes place.

XeF₆ + H₂O ®  XeOF₄ (colourless liquid) + 2HF

XeF₆ + 2 H₂O ® XeO₂F₂ + 4HF

XeF₆ reacts with fluoride ion donors to form fluoroanions.

XeF₆ + MF®  M⁺ [XeF₇]^–   (M = Na, K, Rb or Cs)

Xenon–oxygen compounds :

Hydrolysis of XeF₄ and XeF₆ with water gives XeO₃.

6 XeF₄ + 12 H₂O  ® 4 Xe + 2 XeO₃ + 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 + 4 HF

XeO₃ is a colourless explosive solid and has a pyramidal molecular structure. XeOF₄ is a colourless volatile liquid and has a square pyramidal molecular structure.

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 purposes. Neon bulbs are used in botanical gardens and in green houses.

Argon is used mainly to provide an inert atmosphere in high temperature metallurgical processes (arc welding of metals or alloys) and for filling electric bulbs. It is also used in the laboratory for handlsing substances that are air–sensitive.

Xenon and Krypton are used in light bulbs designed for special purposes