JEE Main Study Notes for p-Block Elements: Boron, Nitrogen, Oxygen, Carbon Family, Halogens and Nobel Gases

Inorganic Chemistry has equal weightage compared to Organic and Physical Chemistry, but it is relatively simple, scoring and requires less time. Moreover, most of the questions come directly from NCERT. Study in detail to be clear with the concepts of bonding and coordination chemistry. NCERT will teach you all about block chemistry and its reactions. p-Block elements is one of the most important topics asked from Inorganic Chemistry. 

• The chapter on p-Block elements carries a weightage of around 6% in the Chemistry section in JEE Main. Therefore, 1-2 questions are sure to be asked from this chapter. Check JEE Main Chemistry Syllabus
• Questions can be asked from important sub topics such as general trends in physical and chemical properties of elements and compounds like Boron, Silicon, Nitrogen, Phosphorus, Sulphur, and Halogens.
• On a general note, elements that have a position within group 13 (group IIIA) to group 17 (group VIIA) of the periodic table along with group 18 which is the zero group elements form the p-Block of the periodic table together.

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In this article, you will find the necessary notes to ace p-Block elements.

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Positions of p-Block Elements in the Periodic Table

The last electron enters the furthest p orbital in the p-block elements. In the periphery shell, they have 3 to 8 electrons. As we know that the quantity of p orbitals is three, the most drastic number of electrons that can be compelled in a p orbital structure is six. Resultantly, the periodic table numbered from 13 to 18 shows six groups of p-block elements.

• First group: group IIIA known as Boron group

• Second group: group IVA known as Carbon group.

• Third group: group VA known as Nitrogen group.

• Fourth group: group VIA known as Chalcogens.

• Fifth group: group VIIA known as Halogens.

• Sixth group: zero group or group 18 known as Inert or Noble gasses group.

All three kinds of elements - the Metals, Non-Metals, and Metalloids, are available in the p-block. The crisscross path in the p-block isolates each of the elements that are metals from nonmetals. Metals are to be found on the left of the line, and those on the right are non-metals. We uncover the metalloids along the line. Due to the proximity of a wide range of elements, the p-block shows a lot of variety in properties.

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Elements in the p-block of the Periodic Table

Metals

1. Aluminium

2. Gallium

3. Indium

4. Thallium

5. Tin

6. Lead

7. Bismuth

Non-Metals

1. Helium

2. Carbon

3. Nitrogen

4. Oxygen

5. Fluorine

6. Neon

7. Phosphorus

8. Sulphur

9. Chlorine

10. Argon

11. Selenium

12. Bromine

13. Krypton

14. Iodine

15. Xenon

16. Radon

Metalloids

1. Boron

2. Silicon

3. Germanium

4. Arsenic

5. Antimony

6. Tellurium

7. Polonium

8. Astatine

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Boron Family (Group 13 Elements )

Members: B, Al, Ga, In & Tl

• Melting Point: Decreases from B to Ga and then increases up to Tl.

• Ionization Energies: 1st <<< 2nd < 3rd

Metallic Character: Increases from B to Tl. B is non-metal

Boron

Preparation of Boron

From Boric Acid: B2O3(s) + 3Mg(s) → 2B(s) +3 MgO(s)

• From Boron Trichloride

(at 1270 k): 2BCl3+ 3H2 (g) → 2B(s) + 6HCl (g)

(at 900 0C): 2BCl3(g) + 3Zn (s) → 2B(s) + 3 ZnCl2 (s)

• By electrolysis of fused mixture of boric anhydride (B2O3) and magnesium oxide (MgO) & Magnesium fluoride at 1100 0C

2 MgO- → 2Mg + O2(g)

B2O3 + 3Mg → 2B + 3MgO

• By thermal decomposition of Boron hydrides & halides:

B2H6 (g) + Δ → 2B(s) + 3H2 (g)

Compounds of Boron

1. Orthoboric acid (H3BO3)

Preparation of Orthoboric acid

• From borax : Na2B4O7 + H2SO4 + 5H2O → Na2SO4 + 4H3BO3

• From colemanite : Ca2B6O11 + 2SO2 + 11H2O → 2Ca(HSO3)2 + 6H3BO3

Properties of Orthoboric acid

• Action of Heat

• Weak monobasic acidic behavior

B(OH)3 ↔ H3BO3 ↔ H+ + H2O +

Thus, in titration with NaOH, it gives sodium metaborate salt.

H3BO3 + NaOH ↔ NaBO2 + 2H2O

Reaction with Metal Oxide

Reaction with Ammonium boron fluoride

2. Borax (sodium tetraborate) Na2B4O7. 10H2O

Preparation from Boric Acid

4H3BO3 + Na2CO3 --> Na2B4O7 + 6H2O + CO2

Properties of Borax

• Basic Nature

The aqueous solution of borax is alkaline in nature because of its hydrolysis

Na2B4O7 + 3H2O → NaBO2 + 3H3BO3

NaBO2 + 2H2O → NaOH + H3BO3

3. Diborane ( B2H6)

Preparation of Diborane

Reduction of Boron Trifluoride

BF3 + 3LiAlH4 → 2B2H6 + 3 LiAl F4

From NaBH4:

2NaBH4 + H2SO4 → B2H6 + 2H2 + Na2SO4

2NaBH4 + H3PO4 → B2H6 + 2H2 + NaH2PO4

Properties of Diborane

Reaction with water: B2H6 + H2O -->2H3BO3 + 6H2

Combustion: B2H6 +2O2 --? B2O3 + 3H2O ΔH = -2615 kJ/mol

Compounds of Aluminium

1. Aluminium Oxide or Alumina (Al2O3)

2Al(OH)3 +Heat → Al2O3 + 2H2O

2Al(SO4)3 +Heat → Al2O3 + 2SO3

(NH4)2Al2(SO4)3·24H2O --> 2NH3 +Al2O3 + 4SO3 + 25 H2O

2. Aluminum Chloride AlCl3

Properties of Aluminium Chloride

White, hygroscopic solid

• Sublimes at 183 0C

• Forms addition compounds with NH3, PH3, COCl2 etc.

• Hydrolysis: AlCl3 + 3H2O --> Al(OH)3 + 3HCl + 3H2O

Action of Heat: 2AlCl3 .6H2O --> 2Al(OH)3 à Al2O3+ 6HCl + 3H2O

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Carbon Family (Group 14 Elements)

Elements: C, Si, Ge, Sn, & Pb

• Ionization Energies: Decreases from C to Sn and then increases up to Pb.

• Metallic Character: C and Si are nonmetals, Ge is a metalloid, and Sn and Pb are metals.

Catenation: C and Si display a tendency to bond with its own atoms to form long chain polymers.

Compounds of Carbon

1. Carbon Monoxide

Preparation of Carbon Monoxide

By heating carbon in limited supply of oxygen: C + 1/2O2 --> CO.

• By heating oxides of heavy metals e.g. iron, zinc etc with carbon.

  • Fe2O3 + 3C → 2Fe + 3CO

  • ZnO + C → Zn + CO

• By passing steam over hot coke: C + H2O → CO + H2 (water gas)

By passing air over hot coke: 2C + O2 + 4N2 → 2CO + 4N2 (Producer gas)

Properties of Carbon Monoxide

A powerful reducing agent : Fe2O3 + 3CO → 2Fe + 3CO2

CuO + CO → Cu + CO2

Burns in air to give heat and carbon dioxide: CO + 1/2O2 → CO2 + heat.

Tests For Carbon Monoxide

• Burns with a blue flame

• Changes the filter paper that is soaked in platinum or palladium chloride to pink or green.

2. Carbon di-oxide

Preparation of Carbon di-oxide

• By action of acids on carbonates: CaCO3 + 2HCl → CaCl2 + H2O + CO2

• By combustion of carbon: C + O2 → CO2

Properties of Carbon di-oxide

Turns lime water milky Ca(OH)2 + CO2 → CaCO3 ¯ + H2O,

• Milkiness disappears when CO2 is passed in excess
  CaCO3 + H2O + CO2 → Ca(HCO3)2

• Solid carbon dioxide or dry ice is availed by cooling CO2 under pressure. It changes from the solid state straight to gaseous state without liquefying (hence dry ice).

3. Carbides

Salt like Carbides : Ionic salts containing either C22- (acetylide ion) or C4- (methanide ion)e.g. CaC2, Al4C3, Be2C.

• Covalent Carbides : Carbides of non-metals such as silicon and boron. The atoms of two elements are bonded to each other through covalent bonds. SiC also known as Carborundum.

• Interstitial Carbides : Formed by transition elements and consist metallic lattices with carbon atoms in the interstices. e.g. tungsten carbide WC, vanadium carbide VC.

Compounds of Silicon

1. Sodium Silicate (Na2SiO3)

Prepared by fusing soda ash with pure sand at high temperature:

Na2CO3+ SiO3 → Na2SiO3 +CO2

2. Silicones

Silicon polymers containing Si – O – Si linkages are formed by the hydrolysis of alkyl or aryl substituted chlorosilanes and their subsequent polymerisation.

3. Silicates

Salts of silicic acid, H4SiO4 comprised of SiO44- units, having tetrahedral structure formed as a result of sp3hybridization.

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Nitrogen Family (Group 15 Elements)

Elements: N, P, As, Sb & Bi

• Atomic Radii: Increases down the group. Only a small increase is seen from As to Bi.

• Oxidation state: -3 to +5. Stability of +3 oxidation state increases down the group.

Ionization energy: Decreases from N to Bi.

Nitrogen

Preparation of Nitrogen

3CuO + 2NH3 + Heat --> N2 + Cu + 3H2O

• CaOCl2 + 2NH3 + Heat --> CaCl2+ 3H2O + N2

• NH4NO2 +Heat --> 3H2O + N2 +Cr2O3

Properties of Dinitrogen

Formation of Nitrides (with Li, Mg, Ca & Al): Ca + N2 +Heat → Ca3N2

• Oxidation: N2 + O2 → 2NO

• Reaction with carbide (at 1273 K): CaC2 + N2 → CaCN2 + C

Oxides of Nitrogen

Oxy -Acids of Nitrogen

Ammonia (NH3)

Preparation of Ammonia

By heating an ammonium salt with a strong alkali ;NH4Cl + NaOH --> NH3­ + NaCl + H2O

• By the hydrolysis of magnesium nitride: Mg3N2 + 6H2O --> 3Mg(OH)2 + 2NH3.
• Haber's process : N2(g) + 3H2(g) --> 2NH3(g).

Properties of Ammonia

Basic nature : Its aqueous solution is basic in nature and turns red litmus blue.

NH3 + H2O ↔ NH4+ + OH-

• Reaction with halogens:
  • 8NH3 + 3Cl2 --> 6NH4Cl + N2
  • NH3 + 3Cl2 (in excess) → NCl3 + 3HCl
  • 8NH3 + 3Br2 → 6NH4Br + N2
  • NH3 + 3Br2 (in excess) → NBr3 + 3HBr
  • 2NH3 + 3I2 → NH3.NI3 + 3HI
  • 8NH3.NI3 → 6NH4I + 9I2 + 6N2
• Complex formation:
  • Ag+ + NH3 → [Ag(NH3)2]+
  • Cu2+ + 4NH3 → [Cu(NH3)4]2+
  •  Cd2+ + 4NH3 → [Cd(NH3)4]2+

Precipitation of heavy metal ions from the aq. solution of their salts:

FeCl3 + 3NH4OH → Fe(OH)3 + 3NH4Cl

Brown ppt.

• AlCl3 + 3NH4OH → Al(OH)3 + 3NH4Cl
  White ppt.

CrCl3 + 3NH4OH → Cr(OH)3 + 3NH4Cl

Phosphorus

Allotropy of Phosphorus

a) White phosphorus

Translucent white waxy solid

• Highly reactive
• Poisonous and insoluble in water

b) Red Phosphorus

Formed by heating white phosphorus in the absence of air.

Does not burn spontaneously at room temperature.

c) Black Phosphorus: Formed by extra heating of red phosphorus.

Compounds of Phosphorus

a) Phosphine, PH3

Preparation

• Ca3P2 + 6H2O → 2 PH3 + 3 Ca(OH)2
• 4H3PO3 +Heat → PH3+ 3 H3PO4
• PH4I +KOH → PH3+KI + H2O
• P4 + 3KOH + 3H2O → PH3 +3KH3PO2

Properties

Formation of Phosphonium Iodide: PH3 + HI à PH4I

Combustion: PH3 + 2O2 à H3PO4

b) Phosphorous Halides

Preparation

P4+ 6Cl2 → 4PCl3

• P4+ 10Cl2 → 4PCl5
• P4+ 8SOCl2 → 4PCl3 + 4SO2+ 2S2Cl2
• P4+ 10SOCl2 → 4PCl5 + 10SO2

Properties

PCl3 + 3H2O → H3PO3 + 3HCl

• PCl5 + 4H2O → POCl3 à H3PO4 +5HCl

• PCl3 + 3CH3COOH → 3 CH3COCl +H3PO3

• PCl5 + CH3COOH → CH3COCl + POCl3+ HCl

• 2Ag + PCl5 → 2AgCl + PCl3

• 2Sn + PCl5 → SnCl4 + 2PCl3

• PCl5 + Heat → PCl3 + Cl2

c) Oxides of Phosphorus

d) Oxy – Acids of Phosphorus

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Oxygen Family (Group 16 Elements)

Chemical Properties

Formation of volatile Hydrides

Formation of Halides

Formation of Oxides

a) All the elements (except Se) form monoxide.

b) All the elements form dioxide with formula MO2. SO2 is a gas, SeO2 is volatile solid, while TeO2 and PoO2 are non – volatile crystalline solids.

c) Ozone: It is unstable and easily reduces into oxygen. It behaves like a strong oxidising agent due to the ease with which it can liberate nascent oxygen.

Oxyacids

Allotropes of Sulphur

1. Rhombic sulphur

• It has a bright yellow colour.

• It is insoluble in water and carbon disulphide. Its density is 2.07 gm cm-3 and exists as S8 molecules. The 8 sulphur atoms in the S8­ molecule form a puckered ring.

2. Monoclinic Sulphur

• Stable only above 369 K. It is a dull yellow coloured solid, also known as b - sulphur. It is soluble in CS2 but insoluble in H2O.

• It gradually changes into rhombic sulphur. It also exists as S8 molecules which possess a puckered ring structure. It differs from the rhombic sulphur in the symmetry of the crystals.

3. Plastic Sulphur

• It is obtained by pouring molten sulphur to cold water.

• It is an amorphous form of sulphur.

• It is insoluble in water as well as CS2.

Sulphuric Acid

Due to its strong affinity for water, H2SO4 acts as a powerful dehydrating agent.

• Concentrated H2SO4 reacts with sugar, wood, and paper to form black mass of carbon. This phenomenon is called charring.

• It is a moderately strong oxidizing agent.

• Decomposes carbonates, bicarbonates, sulphides, sulphites, thiosulphates and nitrites at room temperatures.

• Salts like chlorides, fluorides, nitrates, acetates, oxalates are decomposed by hot conc. H2SO4 liberating their corresponding acids.

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Halogen Family ( Group 17 Elements)

Interhalogen compounds

Hydrogen Halides

Properties of Hydrogen Halides

· All the three acids are reducing agents. HCl is not attacked by H2SO4.

• • 2HBr + H2SO4 → 2H2O + SO2 + Br2 ­

  • 2HI + H2SO4 → 2H2O + SO2 + I2

• All the three react with KMnO4 and K2Cr2O7

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

  • K2Cr2O7 + 14HBr → 2KBr + 2CrBr3 + 7H2O + 3Br2­

• Other reactions are similar.

  • Dipole moment : HI < HBr < HCl < HF

  • Bond length: HF < HCl < HBr < HI

  • Bond strength: HI < HBr < HCl < HF

  • Thermal stability: HI < HBr < HCl < HF

  • Acid strength: HF < HCl < HBr < BI

o Reducing power: HF < HCl < HBr < HI

Pseudohalide ions and pseudohalogens

Ions that have two or more atoms of which at least one is nitrogen and have properties like those of halide ions are known as pseudohalide ions. A few of these pseudohalide ions could get oxidised to form covalent dimers comparable to halogens (X­2). Such covalent dimers of pseudohalide ions are known as pseudohalogens.

The best known pseudohalide ion is CN–

Pseudohalogen

· (CN)2 cyanogen

· (SCM)2 thiocyanogen

Important stable compounds of Xenon

XeO3 Pyramidal

• XeO4 Tetrahedral

• XeOF4 Square pyramidal

XeO2F2 Distorted octahedral

Did you know?

The first rare gas compound was discovered by Bartlett and was known as Xe+ (PtF6]–.

Oxyacids of Chlorine

Acidic Character: Acidic character of halogens increases with the increase in oxidation number of the halogen: HClO4 > HClO3> HClO2 > HOCl

Preparation

HOCl :

· Ca(OCl)2 + 2HNO3 → Ca(NO3)2 + 2HOCl

HClO2 :

· BaO2 + 2ClO2 → Ba(ClO2)2 (liquid) + O2

· Ba(ClO2)2 + H2SO4(dil.) → BaSO4 ¯ + 2HClO2

HClO3 :

· 6Ba(OH)2 + 6Cl2 → 5BaCl2 + Ba(ClO3)2 + 6H2O

· Ba(ClO3)2 + H2SO4(dil.) → BaSO4 ¯ + 2HClO3

HClO4 :

· KClO4 + H2SO4 → KHSO4 + HClO4

· 3HClO3 → HClO4 + 2ClO2 + H2O

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The Noble Gases (Group 18 Elements)

The noble gases are inert. They do not take part in the reactions easily because they have -

• stable electronic configuration (a complete octet).

• high ionization energies.

• low electron affinity.

Practice with JEE Main Question Paper

Compounds of Xenon

Uses of Nobles Gasses

A) Helium

Filling airships and observation balloons.

• Oxygen mixture of deep sea divers.

• Treatment of asthma.

• Inflating aeroplane tyres.

Providing inert atmosphere in melting and welding of easily oxidizable metals.

B) Neon

Filling discharge tubes, that have different characteristic colours and are used for advertising purposes.

Beacon lights for safety of air navigators as the light contains fog and steam perpetrating power.

C) Argon

In addition to nitrogen, it is used in gas–filled electric lamps as it is more inert than nitrogen.

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Properties of Elements in p-Block of Periodic Table

1. Electronic Configuration

The usual electronic design of p-block components for the valence shell is ns 2 np1-6 (except He). The internal core of the electronic structure could however contrast.

The general electronic structure displayed by p-block group 13 to 18 elements is is displayed below -

• Group 13 (Boron family) - ns2 np1
• Group 14 (Carbon family) - ns2 np2
• Group 15 (Nitrogen family) - ns2 np3
• Group 16 (Oxygen family) - ns2 np4
• Group 17 (Halogen family) - ns 2 np5
• Group 18 (Noble gases) - ns2 np6 (except Helium)

Helium 's standard electronic configuration is 1s2. Owing to their unique electronic structure, p-block elements show a great range of properties.

2. Metallic Character

As previously mentioned, p-block includes a broad variety of elements including metals, non-metals, and metalloids. The p-block is the largest locale of the periodic table comprising metalloids. The non-metallic character falls down the group while in the p-block there is a steady change in non-metallic character from left to right. The metallic character continues to increase down each group as it reduces as we pass over a period from left to right. Evidently, the heaviest element in a p-Block group is the most metallic in nature.

3. Atomic Density

The Atomic Density of elements in p-block increases down the group, which is due to the change in the size of the atom down the chain. Although it decreases as we pass over the period from left to right, this is due to the reduction in the nuclear size of all elements in the p-block over the period. Aluminum has a low density of the significant number of elements, and is commonly used as a building material.

4. Melting and Boiling Points

The melting and boiling points increase in the group gradually given that the atomic mass increases down the group and hence the intermolecular forces rise as well. On the other hand, due to greater intermolecular forces (van der Waals interaction), the melting point of group 17 and 18 rises down the group.

5. Oxidation State

The p-block elements exhibit a variable state of oxidation. Within the periodic table, the oxidation shows increments as we pass from left to right. The greatest state of oxidation that a p-block element has is equivalent to the aggregate number of valence electrons. As shown, the oxidation states found in different classes are as follows -

• Boron family (Group 13): - + 3
• Carbon family (Group 14): - + 4
• Nitrogen family (Group 15): - + 5
• Oxygen family (Group 16): - + 6
• Halogen family (Group 17): - + 7
• Noble gases (Group 18): - + 8

In any case , given such p-block elements, there might also be other oxidation states that usually vary by a unit of two from an aggregate number of valence electrons. The other two units of oxidation state, not exactly the group oxidation state that appeared by different groups, are as follows -

• Boron family (Group 13): - + 1
• Carbon family (Group 14):- + 2, -4
• Nitrogen family (Group 15):- + 3, -3
• Oxygen family (Group 16): - + 4, + 2, -2
• Halogen family (Group 17): - + 5, + 3, + 1, -1
• Noble gases (Group 18): - + 6, + 4, + 2

The relative inertia of these two oxidation states , the group oxidation state and the other oxidation state, two units less than the group oxidation state, may change from group to group in any case.

6. Atomic and Ionic Radii

As we pass down the p-block group, one additional shell is included in the following element compared to the previous one. This eventually increases the nuclear and ionic radius of every next element down the group, showing that the nuclear and ionic radii are increasing down the group. Over the period, the trend is not the same. When we pass in a period to the right, the Atomic radii and the Ionic radii of p-block elements are decreasing. Atomic radius from Boron to Aluminium increases tremendously. This expansion is due to the more prominent effect of filtering produced by the eight electrons shown in the penultimate shell.

7. Electrode Potential

The p-block elements have a positive anode potential, by and large. It reduces the down the groups for the most part.

For example, consider the following halogen-group anode possibilities -

• Fluorine = 2.87 V
• Chlorine = 1.36 V
• Bromine = 1.09 V
• Iodine = 0.53 V

We can say from the scientific data above that the anode potential in the p-block is decreasing down the groups.

8. Ionization Energies

The p-block elements give strong ionization possibilities. Because of the efficient expansion of atomic charge, the ionization energies of p-block elements increase towards the right in a period. The ionization energy values decrease down the group as shown by the general trends and do not decrease smoothly. Ionizing Non-metal energies are greater than metals. It is most intense for a noble gas, as the structure of noble gasses is fully filled. Some elements at the base of the group such as Iron, Silver, Thallium, Bismuth, and so on behave more as a metal of weak ionizing energies.

9. Magnetic Properties

The p block elements Radon, Astatine, Iodine, and Polonium in nature are Non-Magnetic. Tin is Paramagnetic, and the other p-block elements are Diamagnetic in nature.

10. Complex Formation

The small size and the more noticeable charge of the elements of different p-block groups empower them to have a more influential propensity to form complexes compared with the s-block elements. This complex formation propensity diminishes down the group as the size of the atoms increases down the group.

11. Chemical Reactivity

The Chemical Reactivity of elements in p-block increases as we pass in a period to the right. The chemical reactivity of the elements decreases down the group as we step down within a group.

• All of the noble gasses' orbitals are fully packed with electrons, so it is extremely difficult to crack their stability in any manner, whether it be the elimination of electrons or the insertion of electrons. Therefore, the noble gasses exhibit poor compound reactivity. Considering their low reactivity, noble gasses are used routinely in welding, for example, where a non-reactive atmosphere is needed. There are two chemically important groups of non-metals going up before the noble gas family. These are halogens (Group 17) and chalcogens (Group 16). Such two groups of elements have high enthalpies of electron gain which can provide one or two electrons framing an anion promptly or obtain the stable noble gas structure suggesting great chemical reactivity in this manner.

Properties of Halogen

A) All halogens are typically present in a combined form.

B) Fluorine responds rapidly to every material that interacts with it.

C) Chlorine, bromine, and iodine are dynamically less reactive while at the same time, frame compounds with a variety of elements, particularly metals.

D) Halogens are always solid oxidizing agents. The halogens oxidize different compounds, and reduce themselves.

E) All halogens bind directly to frame sodium halides with sodium.

F) All halogens bond to form phosphorus halides with red phosphorus.

G) Halogens immediately react with salt-forming alkali metals.

H) The presence of chlorine, bromine, and iodine may be detected by acidified silver nitrate solution.

Elements in VIA

A) As we reach the right side of the periodic table, the parallels between the elements within the group become evidently more notable. It refers to category VIA, except for polonium which is not included due to its radioactivity.

B) The entities in the VIA form X2 – particles when reacting with extremely electropositive metals.

C) The tendency to decrease to the -2 oxidation is greatly decreased as we move down.

D) Oxygen is a gas at normal pressure and temperature. This occurs in all allotropic structures: O2, which accounts for 21 per cent of the world's air, or O3 (ozone), which slowly decreases to O2.

E) The ozone itself ingests long-wavelength ultraviolet radiation, preventing these toxic rays from touching the surface of the earth, which in some manner or another will increase the risk in malignancy of the human skin and may even contribute to other environmental issues.

F) Selenium and tellurium compounds are of limited commercial interest because they are toxic.

A) The organic reactivity of the metalloids depends on the reactive element. For example, Boron behaves as a non-metal when reacting with sodium, but when reacting with fluorine it goes on like a metal.

B) Hence, we can conclude from the above example that metalloids exhibit varying chemical properties.

C) They act as non-metals when metals react and they act like metals when non-metals react.

D) They are normally oxidized in reactions due to their low electronegativity. The metalloid oxides are usually amphoteric.

Group VA Elements

A) All group VA elements form trihydrates after reaction with hydrogen.

B) The reactivity weakens down the group.

C) The elements in group VA either frame trioxides or pentoxides when oxygen reacts.

D) Often, when reacting with halogens, they form trihalides or pentahalides.

E) All of the group VA elements form binary compounds by reacting with metals.

F) Nitrogen and phosphorus compounds are the most essential compounds of group VA elements.

G) Nitrogen and phosphorus are most widely used as manure.

Group IA and IIA Elements

A) Unlike groups IA and IIA, none of group IIIA elements directly form hydrides when reacting with hydrogen.

B) All group IIIA elements also correspond to form trihalides when reacting with halogens, rather than simply forming halides like group IA and group IIA.

A) Carbon is capable of forming strong bonds with other carbon molecules and frames a wide array of organic compounds along those lines.

B) In the +4 oxidation state, the lead functions as a solid oxidizing agent, increasing two electrons and reducing it to +2 oxidation state after taking up electrons.

C) Lead forms covalent compounds and bonds strongly to carbon in the +4 oxidation state.

D) In addition to the metals, certain tin and lead mixes are of market importance themselves. To repress dental concerns, for example, Tin (II) fluoride (stannous fluoride) is applied to a particular toothpaste.

E) Lead is also used in two business applications. One, the lead-acid storage batteries used to power autos, and the other is in the fuel for vehicles.

12. Conductivity

In p-block, the conductivity of elements increases down the group. The metals present in the p-block are good electricity and heat conductors whilst the non-metals are poor electricity and heat conduits. In the middle of the metals and nonmetals lies the conductivity of metalloids.

Colour of p Block Elements

Group IIIA elements

All of the elements are silvery solids except boron that is brown solid.

Group IVA elements

• Carbon: black in color
• Silicon and germanium: reddish brown or dull grey or black color
• Lead: bluish-white color
• Color of group VA elements:
• Nitrogen: colorless
• Phosphorus: white and red
• Arsenic: yellow and grey solid form
• Antimony: amorphous grey form
• Bismuth: silvery white

Color of group 16 elements:

• Oxygen: is a colorless gas
• Sulphur: pale yellow
• Tellurium: silvery-white

Halogens

• Fluorine: Pale yellow.
• Chlorine: Greenish yellow.
• Bromine: Reddish brown.
• Iodine: Violet black.
• Noble Gases
• Helium: Red
• Neon: Orange
• Krypton: Purple
• Xenon: White
• Radon: Colorless

Flame Coloration

• Boron imparts Bright green color
• Copper (I) impart Blue color
• Copper (II) (non-halide) impart Green color
• Copper (II) (halide) impart Blue- green color
• Indium and selenium impart Blue color
• Phosphorus impart Pale bluish green color
• Lead imparts Blue/White color
• Antimony and Tellurium impart Pale green color
• Thallium imparts pure green color