The P-Block Elements:Class 12 Chemistry NCERT Chapter 7

Key Features of NCERT Material for Class 12 Chemistry Chapter 7 – The P-Block Elements

In the previous chapter 6:General Principles And Processes of isolation of elements, you have studied about the occurence or extraction of metals.In this chapter: The p-block elements you will study about the elements of group 13, 14, 15, 16, 17 and 18 of periodic table.

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(Chapter 7:The p-block elements)

What are the p – Block Elements? 

We have demonstrated the p square components in the chart beneath. These components are generally acceptable channels of power. Since they tend to lose their electrons. Additionally, they have a sparkling shine. 

In this square, you will discover probably the most astonishing and assorted properties of different components, similar to Gallium. It is a p square metal that can really soften in your palms. Then again p square components additionally have silicon that is a metalloid. It is a significant part really taking shape of glass.

More on P Block Elements 

A conspicuous trait of these p square components is that the last electron of every one of these components enters the peripheral p-subshell. P square components involve the different families that include: 

• Boron family 
• Nitrogen family 
• Oxygen family 
• Fluorine family and 
• Neon family, or the group of the inactive gases. 

In this way, we see that the P square beginnings from the thirteenth gathering and goes to eighteenth gathering in the intermittent table.

1. The p-Block components: Elements having a place with bunches 13 to 18 of the occasional table are called p-square components. 
2. General electronic arrangement of p-square components: The p-square components are described by the ns2np1-6 valence shell electronic design. 

• Agent components: Elements having a place with the s and p-hinders in the intermittent table are known as the delegate components or primary gathering components. 
• Inactive pair impact: The inclination of ns2 electron pair to take part in bond development diminishes with the expansion in nuclear size. Inside a gathering the higher oxidation state turns out to be less steady regarding the lower oxidation state as the nuclear number increments. This pattern is called ‘latent pair impact’. At the end of the day, the vitality required to unpair the electrons is more than vitality delivered in the development of two extra bonds.

(Chapter 7:The p-block elements)

GROUP 15 ELEMENTS:

• Nitrogen family: The components of gathering 15 – nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi) have a place with arrangement is . 

Nuclear and ionic radii: 

• Covalent and ionic radii increment down the gathering. 
• There is considerable increment in covalent radii from N to P. 
• There is little increment from As to Bi because of quality of totally filled d or f orbitals in substantial components. 

Ionization vitality: 

• It continues diminishing down the gathering because of increment in nuclear size. 
• Gathering 15 components have higher ionization vitality than bunch 14 components because of littler size of gathering 15 components. 
• Gathering 15 components have higher ionization vitality than bunch 16 components since they have stable electronic setup i.e., half-filled p-orbitals. 
• Allotropy: All components of Group 15 with the exception of nitrogen show allotropy. 

Catenation: 

• Nitrogen shows catenation somewhat because of triple bond however phosphorus shows catenation to most extreme degree. 
• The inclination to show catenation diminishes down the gathering. 

Oxidation states: 

• The basic oxidation states are +3, +5 and – 3. 
• The inclination to show – 3 oxidation state diminishes down the gathering on account of decline in electronegativity by the expansion in nuclear size. 
• The soundness of +5 oxidation state diminishes while solidness of +3 oxidation state increments because of inactive pair impact. 
• Nitrogen shows oxidation states from – 3 to +5. 
• Nitrogen and phosphorus with oxidation states from +1 to +4 experience oxidation just as decrease in acidic medium. This procedure is called disproportionation. 

Reactivity towards hydrogen: 

• All gathering 15 components from trihydrides, . 
• It has a place with hybridisation. 
• The solidness of hydrides diminishes down the gathering because of reduction in bond separation vitality down the gathering. 
• Breaking point: 
• Breaking point increments with increment in size because of increment in van der Waals powers. 
• Breaking point of NH3 is more a result of hydrogen holding. 

Bond point: 

• Electronegativity of N is most noteworthy. Consequently, the solitary sets will be towards nitrogen and henceforth more aversion between bond sets. Along these lines bond edge is the most elevated. After nitrogen, the electronegativity diminishes down the gathering. 
• Basicity diminishes as NH3> PH3> AsH3> SbH3< BiH3. This is on the grounds that the solitary pair of electrons are focused more on nitrogen and thus the basicity will be most extreme on account of NH3. It will diminish down the gathering as the electronegativity diminishes down the gathering. The lessening intensity of hydrides increments down the gathering because of reduction in bond separation vitality down the gathering. 
• Reactivity towards oxygen: 
• All gathering 15 components from trioxides () and pentoxides (). 
• Acidic character of oxides diminishes and basicity increments down the gathering. This is on the grounds that the size of nitrogen is extremely little. 
• It has a solid positive field in an extremely little territory. In this manner, it draws in the electrons of water O-H cling to itself and delivery H+ particles without any problem. 
• As we descend the gathering, the nuclear size increments thus, the acidic character of oxide diminishes and basicity increments down the gathering. 

Reactivity towards halogen: 

• Gathering 15 components structure trihalides and pentahalides. 
• Trihalides: These are covalent mixes and get ionic down the gathering with hybridisation, pyramidal shape. 
• Pentahalidesa). They are lewis acids on account of the nearness of empty d – orbitals.b). They have hybridisation and henceforth have trigonalbirpyamidal shape. 
• PCl5 is ionic in strong state and exist as 
• In PCl5, there are three tropical bonds and two pivotal bonds. The pivotal bonds are longer than tropical bonds on account of more noteworthy aversion from central bonds. 
• Nitrogen doesn’t frame pentahalides because of nonattendance of d–orbitals. 
• Reactivity towards metals: All components respond with metals to frame twofold mixes in – 3 oxidation state. 
• Abnormal conduct of nitrogen: The conduct of nitrogen contrasts from rest of the components.

(Chapter 7:The p-block elements)

Reasons:

1. It has a small size.
2. It does not have d – orbitals

iii. It has high electronegativity

1. It has high ionization enthalpy

• Dinitrogen:

1. a) Preparation:
2. b) Physical Properties:
3. i) It is a colourless, odourless, tasteless and non – toxic gas.
4. ii) It is chemically un-reactive at ordinary temperature due to triple bond in N ≡ N which has high bond dissociation energy.

• Ammonia:

1. Ammonia molecule is trigonal pyramidal with nitrogen atom at the apex.
2. It has 3 bond pairs and 1 lone pair.
3. N is hybridised.
4. Preparation:

Haber’s process:

Pressure 20010 Pa Temperature 773 K Catalyst is FeO with small amounts of and

• Nitric Acid:

Ostwald Process: The NO thus formed is recycled and the aqueous can be concentrated by distillation upto ~ 68% by mass. Further concentration to 98% can be achieved by dehydration with concentrated . Nitric acid is strong oxidizing agent in the concentrated as well as in the dilute state.

Phosphorus:

1. a) It shows the property of catenation to maximum extent due to most stable P – P bond.
2. b) It has many allotropes, the important ones are:
3. i) White phosphorus
4. ii) Red phosphorus

iii) Black phosphorus

• White phosphorus:

1. Discrete tetrahedral P4 molecules
2. Very reactive
3. Glows in dark
4. Translucent waxy solid
5. Soluble in but insoluble in water
6. It has low ignition temperature, therefore, kept under water

• Red phosphorus

1. Polymeric structure consisting of chains of P4 units linked together
2. Less reactive than white phosphorus
3. Does not glow in dark
4. Has an iron grey lustre
5. Insoluble in water as well as

• Black phosphorus

1. Exists in two forms – black phosphorus and black phosphorus
2. Very less reactive
3. Has an opaque monoclinic or rhombohedral crystals

• Phosphine

1. It is highly poisonous, colourless gas and has a smell of rotten fish.
2. Preparation
   
   
   

• Chlorides of Phosphorous:

1. a) Phosphorus Trichloride
2. i) It is a colourless oily liquid.
3. ii) Preparation

iii) With water,

It gets hydrolysed in the presence of moisture.

1. iv) Pyramidal shape, sp3 hybridisation
2. v) With acetic acid

vi). With alcohol

1. b) Phosphorus pentachloride

1. Yellowish white powder.
2. Trigonalbipyramidal shape, sp3dhybridisation .
3. Preparation
4.  
5. With water
6.  
7. With acetic acid
8. With alcohol
9. With metals

(Chapter 7:The p-block elements)

GROUP 16 ELEMENTS:

• Oxidation states:

1. They show -2, +2, +4, +6 oxidation states.
2. Oxygen does not show +6 oxidation state due to absence of d – orbitals.
3. Po does not show +6 oxidation state due to inert pair effect.
4. The stability of -2 oxidation state decreases down the group due to increase in atomic size and decrease in electronegativity.
5. Oxygen shows -2 oxidation state in general except in and
6. Thus, the stability of +6 oxidation state decreases and +4 oxidation state increases due to inert pair effect.

• Ionisation enthalpy:

1. Ionisation enthalpy of elements of group 16 is lower than group 15 due to half-filled p-orbitals in group 15 which is more stable.
2. However, ionization enthalpy decreases down the group.

• Electron gain enthalpy:

1. Oxygen has less negative electron gain enthalpy than S because of small size of O.
2. From S to Po electron gain enthalpy becomes less negative to Po because of increase in atomic size.

• Melting and boiling point:

1. It increases with increase in atomic number.
2. Oxygen has much lower melting and boiling points than sulphur because oxygen is diatomic ( ) and sulphur is octatomic ().

• Reactivity with hydrogen:

1. All group 16 elements form hydrides.
2. They possess bent shape.
3. Bond angle:

• Acidic nature:
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.

• Thermal stability:
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.

• Reducing character:
  This is because the H-E bond length increases down the group. Therefore, the bond dissociation enthalpy decreases down the group.

• Reactivity with oxygen:  and

1. Reducing character of dioxides decreases down the group because oxygen has a strong positive field which attracts the hydroxyl group and removal of becomes easy.
2. Acidity also decreases down the group.
3. is a gas whereas SeO2 is solid. This is because has a chain polymeric structure whereas SO2 forms discrete units.

• Reactivity with halogens: EX2, EX4 and EX6

1. The stability of halides decreases in the order
2. This is because E-X bond length increases with increase in size.
3. Among hexa halides, fluorides are the most stable because of steric reasons.
4. Dihalides are hybridised and so, are tetrahedral in shape.
5. Hexafluorides are only stable halides which are gaseous and have hybridisation and octahedral structure.
6. is a liquid while H2S is a gas. This is because strong hydrogen bonding is present in water. This is due to small size and high electronegativity of O.

Oxygen:

The compounds of oxygen and other elements are called oxides.

• Oxides: The compounds of oxygen and other elements are called oxides.

• Types of oxides:

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1. Acidic oxides: Non- metallic oxides are usually acidic in nature.
2. Basic oxides: Metallic oxides are mostly basic in nature. Basic oxides dissolve in water forming bases e.g.,
3. Amphoteric oxides: They show characteristics of both acidic as well as basic oxides.
   
4. Neutral oxides: These oxides are neither acidic nor basic. Example: Co, NO and N2O

• Ozone:

1. Preparation: It is prepared by passing silent electric discharge through pure and dry oxygen 10 – 15 % oxygen is converted to ozone.
   
2. Structure of Ozone: Ozone has angular structure. Both O = O bonds are of equal bond length due to resonance.

• Sulphur:

1. Sulphur exhibits allotropy:
   1. Yellow Rhombic ( – sulphur)
   2. Monoclinic (- sulphur)
2. At 369 K both forms are stable. It is called transition temperature.
3. Both of them have S8 molecules.
4. The ring is puckered and has a crown shape.
5. Another allotrope of sulphur – cyclo ring adopts a chair form.
6. S2is formed at high temperature (1000 K).
7. It is paramagnetic because of 2 unpaired electrons present in anti bonding* orbitals like O2.

Sulphuric acid:

By contact process

1. Preparation:
2. Exothermic reaction and therefore low temperature and high pressure are favourable.
3. It is dibasic acid or diprotic acid.
4. It is a strong dehydrating agent.
5. It is a moderately strong oxidizing agent.

(Chapter 7:The p-block elements)

GROUP 17 ELEMENTS:

• Atomic and ionic radii: Halogens have the smallest atomic radii in their respective periods because of maximum effective nuclear charge.

• Ionisation enthalpy: They have very high ionization enthalpy because of small size as compared to other groups.

• Electron gain enthalpy:

1. Halogens have maximum negative electron gain enthalpy because these elements have only one electron less than stable noble gas configuration.
2. Electron gain enthalpy becomes less negative down the group because atomic size increases down the group.

• Electronegativity:

1. These elements are highly electronegative and electronegativity decreases down the group.
2. They have high effective nuclear charge.

• Bond dissociation enthalpy:

1. Bond dissociation enthalpy follows the order:
2. This is because as the size increases bond length increases.
3. Bond dissociation enthalpy of Cl2 is more than F2 because there are large electronic repulsions of lone pairs present in F2.

• Colour: All halogens are coloured because of absorption of radiations in visible region which results in the excitation of outer electrons to higher energy levels.

• Oxidising power:

1. All halogens are strong oxidisingagents because they have a strong tendency to accept electrons.
2. Order of oxidizing power is:

• Reactivity with Hydrogen:

1. Acidic strength: HF <HCl<HBr< HI
2. Stability: HF >HCl>HBr> HI. This is because of decrease in bond dissociation enthalpy.
3. Boiling point: HCl<HBr< HI < HF. HF has strong intermolecular H bonding. As the size increases van der Waals forces increases and hence boiling point increases.
4. % Ionic character: HF >HCl>HBr> HI Dipole moment: HF >HCl>HBr> HI. Electronegativity decreases down the group.
5. Reducing power: HF <HCl<HBr< HI

• Reactivity with metals:

1. Halogens react with metals to form halides.
2. Ionic character: MF >MCl>MBr> MI. The halides in higher oxidation state will be more covalent than the one in the lower oxidation state.

• Interhalogen compounds:

Reactivity of halogens towards other halogens:

1. Binary compounds of two different halogen atoms of general formula X are called interhalogen compounds where n = 1, 3, 5, or 7. All these are covalent compounds.
2. Interhalogen compounds are more reactive than halogens because X-X is a more polar bond than X-X bond.
3. All are diamagnetic.
4. Their melting point is little higher than halogens.
5. XX’ (CIF, BrF, BrCl, ICl, IBr, IF) (Linear shape) XX’3 () (Bent T- shape) XX’5 –, (square pyramidal shape) XX’7 – (Pentagonal bipyramidal shape)

• Oxoacids of halogens:

1. Fluorine forms only one oxoacid HOF (Fluoric (I) acid or hypofluorous acid) due to high electronegativity.
2. Acid strength:
3. Reason:
4. Acid strength: HOF >HOCl>HOBr> HOI. This is because Fluorine is most electronegative.

(Chapter 7:The p-block elements)

GROUP 18 ELEMENTS:

• Ionisation enthalpy:

1. They have very high ionization enthalpy because of completely filled orbitals.
2. Ionisation enthalpy decreases down the group because of increase in size.

• Atomic radii: Increases down the group because the number of shells increases down the group.

• Electron gain enthalpy: They have large electron gain enthalpy because of stable electronic configuration.
• Melting and boiling point: It has low melting and boiling point due to the presence of only weak dispersion forces.

• Shapes:

is linear, is square planar and is distorted octahedral. is known but no true compound of He Ne and Arare known.

• Compounds of Xe and F:

and are powerful fluorinating agents.

• Compounds of Xe and O:

(Chapter 7:The p-block elements)

Questions:

Q: Why do noble gases not participate in chemical reactions?

Ans: The noble gases are inert in nature. They do not participate in the reactions easily because of their stable electronic configuration, high ionization energies and low electron affinity.