p-Block Formulas

Learning and memorizing p-Block Formulas helps you to kickstart your exam preparation. So, get p-Block Formulae Sheet & Tables and grasp the concept in a better way. Our P-block formula tables aid you to make your problems easy and quick during your homework, assignments, and exam preparation. Also, you will get a great hold on all the topics included in this p-block formulas list. 

BORON FAMILY
Introduction:
The boron family consists of five elements: – Boron, aluminium, gallium, indium and thallium.

Oxidation state:
These elements show +1 and +3 oxidation states ns², np¹.

Tendency to form ionic compounds:
It increases from B to Tl.

Chemical properties:
1. Reducing properties:
Due to high charge density (M⁺³) the standard oxidation potential of these elements are quite high.

2. Complex formation:
Due to higher charge density (M⁺³) these elements have much greater tendency to form complexes than the s-block elements.

Chemistry Of Boron
Occurrence
Boron is found in combined states. Its important minerals are-
Borax (Tincal) → Na₂B₄O₇. 10H₂O
Kernite (Rasorite) → Na₂B₄O₇. 4H₂O
Boric acid → H₃BO₃

Extraction of Boron:
1. Chemical method :- Boron is extracted from borax minerals.

2. Extraction of Boron : –

(b) By Van – Arkel method : –

COMPOUNDS OF BORON
Boron forms numerous compounds. Some of the compounds have special importance in chemistry. The main compounds are. Boron trioxide (B₂O₃), Borax (Na₂B₄.10H₂O) or Suhaga or Titical. Orthoboric acid, Boron halides, Boron, hydrides, Boron nitride, etc.

a. Structure of borax

b. Structure of borazole

c. Structure of diborane

d. Structure of boron nitride

Boric acid:-
Ortho boric acid does not donate proton, like most of the acids but rather it accepts OH^–, therefore it is a lewis acid and is represented by B(OH)₃.
B(OH)₃ + 2H₂O ⇌ H₃O⁺ + [B(OH)₄]^–pk = 9.25

(b) At higher concentration polymeric metaborate species are formed.
3B(OH)₃ ⇌ H₃O⁺ + [B₃O₃(OH)₄]^– + H₂O pk = 6.84

(c) B(OH)₃ partially reacts with water to form H₃O⁺ and [B(OH)₄]^– and behaves as a weak acid. Thus B(OH)₃ can not be titrated satisfactorily with NaOH because a sharp end point is not obtained. If polyhydroxy compound like glycerol, mannitol or sugar are added to titration mixture than is can be titrated with NaOH.
B (OH)₃ + NaOH ⇌ Na[B(OH)₄]

(d) At 100° CH₃BO₃ losses water and convert into metaboric acid.

metaboric acid form tetraboric acid on heating at 160°C.

On strong heating B₂O₃ is produced
H₂B₄O₇ → 2B₂O₃ + H₂O

(e) Boric acid react with ethyl alcohol to form ethyl borate which bums with green edged flame.
H₃BO₃ + 3C₂H₅OH → B(OC₂H₅)₃ + 3H₂O

Boron trihalides (BX₃, (X = F, Cl, Br, I)
Properties:
The relative lewis acid character of boron trihalides is found to follow the following order.
BI₃ > BBr₃ > BCl₃ > BF₃

Boron hydrides
(a) Boron forms a number of hydrides. These are called boranes by analogy with alkanes.

(b) These belong to one of the two B_nH_n+4 and B_nH_n+6.
series, viz,

• B_nH_n+4 (called nido-boranes) B₂H₆, B₅H₉, B₆H₁₀, B₁₀H₁₄.
• B_nH_n+6 (called arachno-boranes) B₄H₁₀, B₅H₁₁, B₆H₁₂, B₉H₁₅.

Borax (Na₂B₄O₇.H₂O)
(a) Aqueous solution of borax is basic in nature

(b) Reaction with strong acid: –
Na₂B₄O₇ + H₂SO₄ + 5H₂O → Na₂SO₄ + 4H₃BO₃

(c) Reaction with NaOH: –
Na₂B₄O₇ + 2NaOH → 4Na₂BO₂ + H₂O

5. CHEMISTRY OF ALUMINIUM:
5.1 Occurrence:
The important minerals of Al are
(i) Oxides – Corundum, ruby, sapphire, emerald, Al₂O₃,
– diaspore Al₂O₃.H₂O
– gibbsite Al₂O₃.3H₂O
– Bauxite Al₂O₃.2H₂O
(ii) Fluoride – Cryolite Na₃AlF₆
(iii) Basic sulphate – Alunite or alum stone,
K₂SO₄.Al₂(SO₄)₃.4Al(OH)₃
(iv) Basic phosphate – Turquoise AlPO₄Al(OH)₃.H₂O
(v) Silicates – Felspar, KAlSi₃O₈, Kaoline, porcelain,
mica, china clay, slate etc,
Al₂O₃.2SiO₂.2H₂O

Extraction:
The extraction of aluminium from bauxite ore involves the following three steps.

• Concentration of the ore i.e., removal of ferric oxide (By Baeyer’s process) and silica (By Serpeck’s process).
• Electrolytic reduction of Al₂O₃.
• Electrolytic purification of aluminium.

CARBON FAMILY
Introduction: Carbon family is basically the group IV A or Group 14 of the periodic table.

GENERAL CHARACTERISTICS OF CARBON FAMILY
1. Electronic configuration:
Each element of this group possesses ns2np2 electronic configuration

2. Metallic character:

• Metallic character increases from top to bottom.
• This change is due to less effective nuclear charge and increased number of available orbitals with increase in the size of the atom.

3. Oxidation state:
The configuration of the valency shell ns²p² of these elements indicates that these elements can loose or gain 4 electrons to attain inert gas configuration. That implies that they can show – 4 to + 4 oxidation state.

4. Melting and boiling points:
Low, the reason of lower melting points is weaker M-M bonds and that they do not use all the four electrons for metallic bonding.

5. Complex forming tendency:

• Except carbon all these elements have complex forming tendency.
• This is due to small size, high charge and availability of vacant orbitals of appropriate energy.

6. Allotropy: Various allotropes are (except lead) –

• C – diamond, graphite, coal, charcoal and lamp black
• Si – crystalline and amorphous
• Ge – two crystalline forms
• Sn – grey Sn, white Sn & rhombic Sn

7. Chain forming tendency ( Catenation or self linkage):

• All elements of this group have catenation tendency
• This property decreases from carbon to lead.

8. Reactivity: The elements of this group are relatively less reactive but reactivity increases down the group.

9. Halides:
(i) The thermal stability of tetrahalides of the same element decreases with increases in molecular mass of the tetrahalide :
CF₄ > CCl₄ > CBr₄ > CI₄

(ii) The thermal stability of tetrahalides of the elements with a common halogen decreases with increasing atomic number:
CX₄ > SiX₄ > GeX₄ > SnX₄ > PbX₄

(iii) Except carbon halides, other halides are readily hydrolysed by water. The hydrolysis is due to utilisation of d-orbitals to which water molecules can get attached

(iv) Stability of dihalides increases down the group due to inert pair effect.

COMPOUNDS OF CARBON:
1. Carbon Dioxide (CO₂):
Chemical properties

Recovery of CO₂:

2. Carbon monoxide (CO):
Chemical reactions

Carbides:
(a) Compounds of carbon with elements of nearly the same or lower electronegativity than itself (excluding hydrogen) are called carbides. Therefore O, S, P, N & halogen do not form carbide.

(b) Carbides are of three types according to the type of bonding

• Salt like carbides (ionic bonding)
  Covalent carbides (giant molecule structure)
• Alloy type or interstitial type carbide (metallic bonding)

CHEMISTRY OF SILICON:
It exist in two allotropic form –
(a) Amorphous, which is obtained by heating dry powdered silica with magnesium
SiO₂ + 2Mg → Si + 2MgO

(b) Crystalline, which is obtained by heating a finely powdered quartz with carbon in an electric furnace:
SiO₂ + 2C → Si + 2CO
A small amount of iron is added to present the formation of carborandum (SiC).

COMPOUNDS OF SILICON

• Carborundum (Silicon carbide), SiC
• Sodium silicates, Na₂SiO₃, (Water glass)
• Silicones: These are organo silicon polymers containing

Si – O – Si linkages.
Cyclic (ring) silicones are formed when water is eliminated form the terminal – OH groups of linear silicones.
Cyclic silicon:

Silicates: These are metal derivatives of silicic acid, H₄SiO₄ or Si(OH)₄

Classification of silicates:
Silicates have basic unit of, each silicon atom is bonded with four oxide ions tetrahedrally. The way in which the (SiO₄)^4- tetrahedral, units are linked provides a convenient classification of the many silicate minerals.

(a) Orthosilicates (neso- silicates):
These silicates contain single discrete (SiO₄)^4- tetrahedra, that is they share no corners.
Structure of orthosilicates:

(b) Pyrosilicates (soro-silicates, disilicates, island silicate):
In these silicates two tetrahedral units are joined by sharing the O at one comer, thus giving the unit (Si₂O₇)^6-
Structure of pyrosilicates (Si₂O₇)^6- (After T. Moeller)

(c) Cyclic silicates:

• Cyclic or ring silicates have general formula (SiO₃^2-)_n or (SiO₃)_n^2n-
• Rings containing three, four, six and eight tetrahedral units are known, but those with three and six are most common.

Structure of Cyclic Silicates Si₃O₉^6-:

(d) Chain silicates:

• These are formed by sharing two oxygen atoms by each tetrahedra.
• Anions of chain silicate have two general formula.

(1) (SiO₃)_n^2n-
(2) (SiO₁₁)_n^6n-

e.g. Synthetic silicates Li₂SiO₃Na₂SiO₃
Spodumene LiAl(SiO₃)₂

(e) Two dimensional sheet silicates:
(i) In such silicates, three oxygen atoms of each tetrahedraare shared with adjacent SiO₄^4- tetrahedra.

(ii) Such sharing forms two dimensional sheet structure with general formulae (Si₂O₅)₂^2n- as shown below –

(iii) Some examples are – Talc Mg (Si₂O₅)₂ Mg(OH)₂, Kaolin Al₂(OH)₄ (Si₂O₅)

(f) Three dimensional sheet silicates:
These silicates involve all four oxygen atoms in sharing with adjacent SiO₄^4- tetrahedra. e.g., Quartz, Tridymide, Crystobalite, Feldspars, Zeolites and ultramarines.

CHEMISTRY OF TIN
(a) It is extracted from cassiterite ore. The extraction involves following steps –

• Concentration
• Smelting:

(b) Purification:

• Liquation
• Poling
• Electrolytic refining

Physical properties:
It exists in three allotropic froms. White tin in the ordinary and most stable form.

Chemical properties:
(a) Action of air:

(b) Action of acid:

• Sn + H₂SO₄ (dil.) → SnSO₄ + H₂
• Sn + 2H₂SO₄ (cone.) → SnSO₄ + SO₂ + 2H₂O
• Sn + 4H₂SO₄ (cone.) → Sn(SO₄)₂ + 2SO₂ + 4H₂O

(c) Action of alkalies:
Sn + 2NaOH + H₂O → Na₂SnO₃ + 2H₂(sod. stannate)
or Sn + 2NaOH + 4H₂O → Na₂Sn(OH)₆ + 2H₂

(d) Action of chlorine and sulphur:
Sn + 2Cl₂ → SnCl₄
Sn + 2S → SnS₂

COMPOUNDS OF TIN:

• Stannous oxide (SnO): In stannous oxide, tin is in +2 oxidation state.
• Stannic Oxide, (SnO₂): In stannic oxide (SnO₂), tin is +4 oxidation state.

Preparation

Properties:
(a) SnO₂ + 2H₂SO₄→ Sn(SO₄)₂ + 2H₂O
On fusion with NaOH it gives sodium stannate.
SnO₂ + 2NaOH → Na₂SnO₃ + 2H₂O
Tin (II) Chloride,stannous chloride (SnCl₂):
Tin (II) chloride is commonly called as tin salt.

Properties of SnCl₂:
(a) Hydrolysis:
SnCl₂(milkiness) + H₂O → Sn(OH)Cl + HC1

(b) Reducing character:
Tin (II) chloride (SnCl₂) in acidic medium acts as a strong reducing agent.

(c) With H₂S: SnCl₂ + H₂S → SnS + 2HCl

(c) With sodium hydroxide:
SnCl₂ + 2NaOH → Sn(OH)₂ + 2NaCl
Sn(OH)₂ + NaOH → Na₂SnO₂(Sodium stannite) + 2H₂O

Sodium stannite absorbs atmospheric oxygen to give sodium stannate, Na₂SnO₃
Na₂SnO₂ + O₂ → Na₂SnO₃

Tin (IV) Chloride,stannic chloride (SnCl₄)
Preparation

• Sn(fused) + 2Cl₂ → SnCl₄
• SnCl₂ + Cl₂ (g) → SnCl₄

Properties:
In dilute solutions, in water it undergoes hydrolysis.:

COMPOUNDS OF LEAD:
1. Lead dioxide, PbO₂
Properties

(b) Action of acids
PbO₂ + 4HCl → PbCl₂ + Cl₂ + H₂O
PbO₂ + 4HCl (cold cone.) → PbCl₄ + H₂O

(c) Action of Alkalies

(d) Action of sulphur dioxide to form lead sulphate.

2. Red lead (Minium or sindhur), Pb₃O₄
Preparation:

Properties:

(b) Pb₃O₄ → 4HNO₃ → 2Pb(NO₃)₂ + PbO₂ + 2H₂O

(c) It acts as an oxidizing agent
Pb₃O₄ + 8HCl → 3PbCl₂ + 4H₂O + Cl₂

3. PbCl₂ or Lead chloride or Plumbous chloride
Preparation:
Pb(NO₃)₂ + 2HCl → PbCl₂↓+ 2HNO₃

Properties:
It is fairly soluble in concentrated hydrochloric acid forming chloroplumbous acid.
PbCl₂ + 2HCl → H₂PbCl₄

4. Lead tetrachloride or Plumbic chloride, PbCl4 Preparation :
PbO₂ + 4HCl → PbCl₄ + 2H₂O

Properties:
(a) It is only slightly stable :
PbCl₄ → PbCl₂ + Cl₂

(b) It forms a double chloride with ammonium chloride known as ammonium plumbichloride or ammonium chloro plumbate, (NH₄)₂PbCl₆. Although the later compound is fairly stable, when treated with sulphuric acid it decomposes back to lead tetrachloride.
[(NH₄)₂PbCl₆] + H₂SO₄ → (NH₄)₂SO₄ + PbCl₄ + 2HCl

NITROGEN FAMILY

1. INTRODUCTION:
Group 15 of long form of periodic table (previously reported as group VA according to Mendeleef s periodic table) includes nitrogen (N) phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi), These are known as: “pnicogens”.

2. GENERAL PHYSICAL PROPERTIES OF NITROGEN FAMILY:
Electronic structure and oxidation states – The elements of this group all have five electrons in their outer shell. They show a maximum oxidation state of five towards oxygen by using all five outer electrons in forming bonds and also (+ 3) due to inert pair effect.

Non-metallic nature- Like group 13 and group 14, there is a transition from non- metallic to metallic nature on moving down the group.

Density – The density of these elements also increases down the group

Melting point and boiling point – The m.pt. of pnicogens increases from nitrogen to arsenic and then decreases from arsenic to bismuth

Atomic radii – Atomic radii increase with the increase of atomic number.

Ionisation energy – The first ionisation energy decreases regularly on moving down the group.

Electronegativity – Electronegativity decreases gradually on moving down the group from N to Bi.

Oxidation state and valency – Except N which shows + 5 oxidation state by co-ordinate bonding). All elements of this group exhibit -3 oxidation state (covalency) in hydrides. The tendency to show -3 ionic nature is not possible due to high energy considerations.

However nitrogen being smallest and most electronegative and thus forms N^3- ions in nitrides such as Mg₃N₂, Ca₃N₂ etc. Rest all show -3 oxidation state in covalent compounds. The +3 oxidation states are covalent and ionic both. The tendency to show M³⁺ ionic state however increases down the group due to inert pair effect.

Electrical and thermal conductivities – Both these properties increase on , moving down the group. Thus nitrogen and phosphorus are non-conductors while bismuth is an excellent conductor of electricity.

Catenation – The lighter elements N, P and As exhibit the property of catenation but this property is much less than group 14 elements.

3. GENERAL CHEMICAL PROPERTIES OF NITROGEN FAMILY:
Oxides – X₂O₃, X₂O₄ and X₂O₅.

• Greater is the electronegativity more is the acidic character of its oxides.
• Stability of oxides of higher oxidation states decreases with increasing atomic number. Thus N₂O₅, is quite stable while Bi₂O₅, is not known.

Hydrides – All elements of this group form volatile hydrides of the type MH₃.

• NH₃ is highly soluble in water due to hydrogen bonding but other hydrides are less soluble. The basic character decreases form NH₃ to BiH₃.
• Thermal stability decreases gradually from NH₃ to BiH₃.
• The reducing nature increases.

Halides of Nitrogen family – Elements of Nitrogen family form two types of halides viz. trihalides (MX₃) and pentahalides (MX₅).

Trihalides – All the elements of nitrogen family form trihalides (except NI₃)

Pyramidal structure of phosphorus trihalide, (PX₃)
The tendency towards hydrolysis follows the order
N > P > As > Sb > Bi and I > Br > Cl > F

Pentahalides – Nitrogen does not form pentahalides.

Oxyacids of Nitrogen family – The acids which contain oxyanions are called oxyacids. Except bismuth, all other elements of Group 15 form oxyacids.

Strength of Oxyacids of Group 15 elements – The strength and solubility of oxyacids having central atom in the same oxidation state follow the order. N > P > As > Sb

Sulphides of Nitrogen family –

• With the exception of nitrogen all other elements of this group form sulphides.
• The stability of sulphides increases with increase in atomic number of the elements.

4. CHEMISTRY OF NITROGEN (N2)Preparation:
(i) NH₄Cl + NaNO₂ → NH₄NO₂(Ammonium nitrite) + NaCl
NH₄NO₂ → N₂ + 2H₂O

(ii) 2NH₃ + 3CuO → N₂ + 3Cu + 3H₂O

(iii) NH₂CONH₂(urea) + 2HNO₂ → 2N₂ + CO₂ + 3H₂O

(iv) 5Cu + 2HNO₃ → 5CuO + N₂ + H₂O

(v) Ba(N₃)₂(Barium azide) → 3N₂ + Ba

Chemical Properties:

(iv) 2B + N₂ → 2BN (boron nitride)
(v) CaC₂ + N₂ → (CaCN₂ + C) (Nitrolim)
(vi) CaCN₂ + 3H₂O → Ca(OH)₂ + NH₂CONH₂

5. COMPOUNDS OF NITROGEN
5.1 Ammonia (NH₃) Preparation of ammonia
(a) NH₄Cl + NaOH → NH₃ + NaCl + H₂O
(b) NH₄Cl + PbO → 2NH₃ + PbCl₂ + H₂O
(c) AIN + 3H₂O → Al(OH)₃ + NH₃
(d) Mg₃N₂ + 6H₂O → 3Mg(OH)₂ + 2NH₃

(f) By Haber’s process (a commerical method)
N₂ + 3H₂ ⇌ 2NH₃; H = -2.40 k cal.

Chemical properties of ammonia
(a) Stability:

(b) Combustion:
4NH₃ + 3O₂ → 2N₂ + 6H₂O

(c) Basic nature:
Ammonia acts as Lewis base, accepting proton to form ammonium ion because of the tendency to donate an electron pair.

(d) Oxidation:
NH₃ is oxidised by CuO or PbO,
3CuO + 2NH₃ → 3Cu + N₂ + 3H₂O
3PbO + 2NH₃ → 3Pb + N₂ + 3H₂O

(e) Formation of amides:
2Na + 2NH₃ → 2NaNH₂(Sodamide) + H₂

(f) Reaction with Nessler’s reagent:
2K₂HgI₄ + NH₃ + 3KOH→ H₂NHgOHgI(Brown ppt.) + 7KI + 2H₂O

6. Oxides of nitrogen: Nitrogen forms a number of oxides. The well known oxides of nitrogen are –

• Nitrous oxide, N₂O,
• Nitric oxide NO
• Nitrogen trioxide, N₂O₃
• Nitrogen dioxide or Dinitrogen tetra oxide NO₂ or N₂O₄
• Nitrogen pentoxide, N₂O₅

6.1 Nitrous oxide, N₂O or laughing gas
Preparation
NH₄NO₃ → N₂O + 2H₂O

Chemical properties

Structure
1.13Å 1.19Å ; dipole moment (0.116 D)
N——-N——-O

6.2 Nitric oxide (NO)
Structure –
Structure of Nitric oxide dimer:

Preparation:
3Cu + 8HNO₃ → 3Cu(NO₃)₂ + 2 NO + 4H₂O
Chemical Properties

6.3 Dinitrogen trioxide, N₂O₃

N₂O₃ + H₂O → 2HNO₂
Structure:

6.4 Nitrogen dioxide, NO₂ or Dinitrogen tetraoxide, N₂O₄
Preparation:

• 2NO₂(Brown gas) ⇌ N₂O₄ (Colourless solid)
• 2Pb(NO₃)₂ → 2PbO + 4 NO₂ + O₂

Properties:

Structure:
NO₂ molecule possess V-shaped structure with O-N-O bond angle 132° and N-O bond length of about 1.19A which is intermediate between a single (N-O) bond and a double (N = O) bond. Hence, NO₂ is supposed to be a resonance hybrid of the following two structures.

6.5 Nitrogen pentoxide, nitric anhydride (N₂O₅)
Preparation:

Properties:

• N₂O₅ + H₂O → 2HNO₃
• I₂ + 5N₂O₅ → I₂O₅ + 10 NO₂

Structure:

7. Oxy Acid of Nitrogen:

7.1 Nitrous acid (HNO₂)
Preparation:
N₂O₃ + H₂O → 2HNO₂
NH₃ + 3H₂O₂ → HNO₂ + 4H₂O

Properties of nitrous acid:
(i) 2HNO₂ ⇌ NO + NO₂ + H₂O
(ii) It shows both oxidising and reducing properties
2HNO₂ → 2NO + H₂O + [O]
H₂S + [O] → H₂O + S ↓

Structure:

The bond angle ONO is 127°.

7.2 Nitric Acid (HNO₃)
Preparation of nitric acid –
(a) laboratory method:
2NaNO₃ + H₂SO₄ → Na₂SO₄ + Na₂SO₄ + 2HNO₃

(b) Birkland – Eyde Process:
I step N₂ + O₂ → 2NO
II step 2NO + O₂ → 2NO₂
III step 3NO₂ + H₂O → 2HNO₃ + NO

(c) Ostwald Process:
I step

II step 2NO + O₂ → 2NO₂
III step 3NO₂ + H₂O → 2HNO₃ + NO^–

Properties:

(ii) C₁₂H₂₂O₁₁ + 18[O] → 6H₂C₂O₄ + 5H₂O

(iii) Mg + 2HNO₃ (dil.) → Mg(NO₃)₂ + H₂ ↑

(iv) 3Mg + 8HNO₃ → 3Mg(NO₃)₂ + 4H₂O + 2NO ↑

(v) with copper: with cold, dil HNO₃
4Cu + 10HNO₃ → 4Cu(NO₃)₂ + 5H₂O + N₂O ↑

(vi) with mercury with cold dilute HNO₃
6Hg + 8HNO₃ → 3Hg₂(NO₃)₂ + 4H₂O + 2NO

(vii) Iron and tin reacts with nitric acid in the same manner e.g.,
4Fe + 10HNO₃(very dil.) → 4Fe(NO₃)₂ + NH₄NO₃ + 3H₂O
4Sn + 10HNO₃ → 4Sn(NO₃)₂ + NH₄NO₃ + 3H₂O

Structure:

Action of heat on metal nitrates:
(a) Alkaline earth and other heavy metal nitrates :
2Ba(NO₃)₂ → 2BaO + 4NO₂↑ + O₂↑

(b) Silver and mercury nitrates
2AgNO₃ → 2Ag + 2NO₂ ↑+ O₂↑

(c) Ammonium nitrate
NH₄NO₃ → N₂O ↑+ 2H₂O

8. CHEMISTRY OF PHOSPHORUS:
Occurence:

• Ca₃(PO₄)₂ Phosphorite ;
• Ca₃(PO₄)₂ CaF₂ Flourapatite
• Ca₃(PO₄)₂. CaCl₂ Chlorapatite

Extraction of phosphorus
(A) Retort process or old process
(B) Electrothermal process or modern process

White phosphorus
Persons working with phosphorus develop a disease known as phossy jaw in which jaw bones decay.

Structure:

Red Phosphorus:

It changes to white phosphorus when it is vaporised and the vapours are condensed.

Structure:
The exact structure of red phosphorus is not yet known. It is, however, believed as a polymer consisting of canines of P4 tetrahedra linked together.

9. COMPOUNDS OF PHOSPHORUS:

• Phosphorus trioxide P₂O₃ or P₄O₆
• Phosphorus pentoxide P₂O₅ or P₄O₁₀
• Phosphorus tetroxides P₂O₄ or P₄O₈

(a) Phosphorus trioxide P₂O₃ or P₄O₆:
P₄ + 3O₂ (limited) → P₄O₆

Heating in air:
P₄O₆ + 2O₂(Phosphorus (V) oxide) → P₄O₁₀

Action of water:
P₄O₆ + 6H₂O (cold) → 4H₃PO₃(Phosphorus acid)

Action with chlorine:

Structure:
Structure of P₄O₆:

(b) Phosphorus (V) oxide (P₄O₁₀)
Preparation :
P₄ + 5O₂ (excess) → P₄O₁₀

Properties:
(i) Action with water:
P₄O₁₀ + 2H₄O(cold) → 4HPO₃(Metaphosphoric acid)
With hot water it gives phosphoric acid.
P₄O₁₀ + 6H₂O(cold) → 4H₃PO₄(Phosphoric acid)

Structure:
Structure of P₄O₁₀

10. Oxvacids of phosphorus:

The structures of various oxyacids of phosphorus are given below –



H₄P₂O₇

10.1 Phosphorus acid, H₃PO₃
Preparation – P₄O₆ + 6H₂O → 4H₃PO₃

Properties

(b) It acts as a strong reducing agent,
H₃PO₃ + H₂O → H₃PO₄ + 2H

10.2 Orthophosphoric acid, H₃PO₄
Properties
(a) Action of heat:

(b) Action of heat on sodium salt –

(c) Acidic nature:

10.3 Pyrophosphoric acid, H₄P₂O₇
Preparation:

Properties

10.4 Metaphosphoric acid, HPO3 Preparation
P₄O₁₀ + 2H₂O → 4HPO₃
Properties

10.5 Phosphine
Preparation of phosphine
P₄ + 3NaOH + 3H₂O → 3NaH₂PO₂(Sod. hypophosphite) + PH₃

Chemical properties
(a) Decomposition:

(b) Combustibility:

(c) Action of chlorine:
PH₃ + 4Cl₂ → PCl₅ + 3HCl

(d) Basic nature:
with HCl, HBr or HI forms phosphonium compounds.
PH₃ + HCl → PH₄Cl (Phosphonium chloride)

(e) Action of nitric acid:
2PH₃ + 16HNO₃ → P₂O₅ + 16NO₂ + 11H₂O

(f) Addition compounds :
AlCl₃ + 2PH₃ → AlCl₃. 2PH₃

(g) Formation of phosphides:
3CUSO₄ + 2PH₃ → Cu₃P₂ + 3H₃SO₄

OXYGEN FAMILY

INTRODUCTION:
Group 16 or VI A of the periodic table consists of oxygen (O), sulphur (S), selenium (Se), tellurium (Te) and polonium (Po). Except polonium, these are called by Chalcogens

2. GENERAL PROPERTIES OF THE GROUP
Metallic nature:
The metallic character increases with increase in atomic number.

Atomic radius, atomic volume and density- increase gradually with increase in atomic number.

Ionisation potentials- The ionisation potentials are high and The values decrease as the atomic number increases from O to Po.

Electro negativity – Electronegativity decreases down the group

Melting and boiling points – The m.p and b.p increase gradually with increase in atomic number.

Allotropy – All the elements show allotropy oxygen as ordinary oxygen and ozone ; Sulphur as Rhombic, monoclinic, plastic and amorphous ; Selenium as Red form (non-metallic), and grey form (metallic form); Tellurium as crystalline and amorphous ; Polonium as a and b forms (Both are metallic forms)

Catenation – Oxygen and sulphur show the property of catenation.

Oxidation states:

• Oxygen shows ‘-2’ state,
• Sulphur shows – 2, + 4, & + 6 states.
• The stability of + 4 oxidation state increases down the group due to inert pair effect

Chemical properties- Hydrides:
H₂M where M = O, S, Se, Te & Po

• “Thermal stability” decreases as the atomic mass increases.
• “Acidic nature” increases down the group.
• “Reducing nature” increases down the group

Oxides: Order of acidity SO₂ > SeO₂ > TeO₂

Halides: These elements forms binary halides and all elements form hexafluorides

Oxyhalides: Only S and Se form oxyhalides. They are called thionyl and selenyl halides.
e. g. SOF₂, SeOF₂, SOCl₂, SeOCl₂ etc

3. CHEMISTRY OF OXYGEN:
Preparation of Oxygen
Laboratory m ethod :

Heating the oxides:

Heating the salts/H₂O₂:

Electrolysis method:
(a) By the electrolysis of acidified water using Platinum electrodes
H₂SO₄ ⇌ 2H⁺ + SO₄^2-
At cathode: 2H⁺ + 2e^– → 2H → H₂
At anode: H₂O + SO₄^2- → 2H⁺ + SO₄^2- + O + 2e^–
O + O → O₂↑

(b) By the electrolysis of alkaline water using Ni electrodes :
Ba(OH)₂ ⇌ Ba²⁺ + 2OH^–
At cathode:
Ba²⁺ + 2H₂O + 2e^– → Ba²⁺ + 2OH^– + H₂ ↑
At anode: 2OH^– → H₂O + O + 2e^–
O + O → O₂↑

Chemical properties:
(a) Reaction with metals & non-metals
4Na + O₂ → 2Na₂O
S + O₂ → SO₂
3Fe + 2O₂ → Fe₃O₄
P₄ + 5O₂ → P₄O₁₀

(b) Reaction with compounds:

(c) Oxidation of organic compounds:
CH₄ + 2O₂ → CO₂ + 2H₂O

Structure of oxygen:

(b) “Pauling “suggested 3-electron bonds as like
(c) KK σ 2s² σ* 2s² σ 2p_z² π 2p_y² π* 2p_x² π* 2p_y¹ π* 2p_x¹
Bond order in this case is \(\frac{1}{2}\) (8 – 4) = 2 i.e. double bond is present between two oxygen atoms.

4. OXIDES – Classification of oxides on the basis of oxygen content
(i) Normal oxides -These contants only M—O bonds e.g. H₂O, MgO, Al₂O₃ etc.

(ii) Poly-oxides – Peroxides (BaO₂, Na₂O₂), super oxides (KO₂, RbO₂) and dioxides (MnO₂, PbO₂).

(iii) Sub-Oxides – N₂O, C₃O₂, Pb₂O etc.

(iv) Mixed (compound) oxides-Red lead, Pb₃O₄(2PbO + PbO₂), Ferrosoferric oxide, Fe₃O₄ (FeO + Fe₂O₃)

Clas sification of oxides on the basis of chemical behaviour
(i) Acidic oxides (Acid anhydrides) – These oxides are generally formed by non-metals.
e.g.: CO₂, SO₂, SO₃, B₂O₃, SiO₂, N₂O₅, Cl₂O₇, I₂O₅ etc.

(ii) Basic oxides- These oxides are formed from metals, e.g.: Na₂O, CaO, CuO, FeO, BaO, PbO etc.

(iii) Neutral oxides – The oxides which neither combine with acids nor with bases to form salts.
e.g.: CO, H₂O, N₂O, NO etc.

(iv) Amphoteric oxides – The oxides which react with acids as well as bases depends upon conditions.
e.g.: ZnO, Al₂O₃, BeO, Sb₂O₃, Cr₂O₃, PbO, PbO₂, SnO, SnO₂ etc.

5. OZONE, (O₃):
Preparation:
(i) 3O₂ ⇌ 2O₃; ΔH = + 68 K cal.

O₂ + O → O₃
3O₂ ⇌ 2O₃ – energy

Properties- Chemical properties:
(i) Decomposition:

(ii) Oxidising nature:

Reducing Nature:
O₃ + O → 2O₂
O₃ reduces peroxides to normal oxides like BaO₂ to BaO and H₂O₂ to H₂O

Bleaching properties:
Coloured substance + O → colourless

Structure of Ozone:

6. CHEMISTRY OF SULPHUR:
The name sulphur has been derived from Sanskrit word sulveri means killer of copper.
Occurrence:
Galena (PbS), Zinc blende (Zns), Gypsum (CaSO₄. 2H₂O), Epsom salt (MgSO₄. 7H₂O) etc.

Extraction
(i) From alkali wastes (which consists of calcium sulphide)
CaS + H₂O + CO₂ → CaCO₃ + H₂S

(ii) From spent oxides of coal gas (contains ferric sulphide):
2Fe₂S₃ + 3O₂ → 2Fe₂O₃ + 6S

(iii) From iron pyrites:

Allotropic forms
(i) Rhombic, octahedral or α-sulphur: (m.p 114.5°C, sp.gr 2.06)
(ii) Monoclinic or prismatic or β-sulphur: (m.p 199.2°C, sp gr 1.96)

(iii) Plastic or γ-sulphur (no sharp m.p, sp.gr.1.95)
(iv) Milk of sulphur:

Colloidal or d-Sulphur – Prepared by passing H₂S through a solution of an oxidising agent (eg., HNO₃, KMnO₄ etc) or water.
2HNO₃ + H₂S → 2NO₂ + 2H₂O + S↓

Chemical properties
(a) Oxidation: S + O₂ → SO₂↑
2SO₂ + O₂ → 2SO₃ ↑

(b) Reducing property
S + 2H₂SO₄ → 2H₂O + 3SO₂

(c) Reaction with alkalies
4S + 6 NaOH → 2Na₂S + Na₂S₂O₃ + 3H₂O
Na₂S + 4S (excess) → Na₂S₅ (sodium pentasulphide)

7. COMPOUNDS OF SULPHUR:

7.1 Hydrogen Sulphide (H₂S)
Preparation: H₂ + S → H₂S

Properties – Chemical properties
(a) Decomposition:

(b) Combustion : It bums with a blue flame.
2H₂S + 3O₂(excess) → 2H₂O + 2SO₂
2H₂S + O₂ (limited) → 2H₂O + 2S.

(c) Reducing agent: due to release of hydrogen by decomposition, it acts as reducing agent.

Tests of H₂S:

• Unpleasant odour resembling that of rotten eggs.
• It turns lead acetate paper black.
• It gives a violet colouration with a solution of sodium nitroprusside.

Structure of H₂S.
(a) Similar to str. of water molecule i.e. V-shaped structure with bond length (H-S) 1.35A0 and bond angle (H-S-H) is 92.5°.

7.2 Sulphur dioxide (SO₂)
Preparation:

Properties (Physical and chemical) –
(a) Decomposed on heating to give SO₃ & S (disproportionate reaction)

(b) Acidic nature:
SO₂ + H₂O → H₂SO₃

(c) Reducing property:

(d) Oxidising property:
SO₂ + 2H₂S → 2H₂O + 3S↓

(e) Bleaching property:
SO₂ + 2H₂O → H₂SO₂ + 2[H]

7.3 Sulphur Trioxide (SO₃):
Preparation:

8. Oxyacids of sulphur:
Oxyacids with S—S links are called thio acids.

8.1 Sulphurous acid, H2S03:
SO₂ + H₂O → H₂SO₃
It is a strong diprotic acid
H₂SO₃ ⇌ H⁺ + HSO₃^–
HSO₂^– ⇌ H⁺ + S₂^2-

(i) Structure:

8.2 Sulphuric acid, H₂SO₄ (Oil of vitriol):
Preparation: H₂SO₄ is prepared by two methods

(a) Lead chamber process:
Stage I: 2NO (Catalyst) + O₂ → 2NO₂
Stage II: 2SO₂ + 2NO₂ → 2SO₃ + 2NO
Stage III: SO₃ + H₂O → H₂SO₄

(b) Contact process:

Chemical properties
(a) As a strong acid:
H₂SO₄ + H₂O → HSO₂^– + H₃O⁺
HSO₄^– + H₂O ⇌ SO₄^2- + H₃O⁺

(b) As an oxidising agent:

(c) As a dehydrating agent:

(d) Displacing property: – Sulphuric acid displaces more volatile acids (like HC1, HN03, H2S) from their metal salts
KCl + H₂SO₄ → KHSO₄ + HCl

(e) As a sulphonating agent:
C₆H₆ + HO.SO₂.OH → C₆H₅. SO₂. OH (Benzene sulphonic acid)

(f) Reaction with PCl₅:
HO.SO₂.OH + PCl₅ → Cl.SO₂.OH + POCl₃ + HCl (Chloro sulphuric acid)

Structure
(a) SO₄2- is a resonance hybrid structure and sulphur atom is sp3
hybridised.

8.3 Sodium thiosulphate, Hypo (Na₂S₂O₃. 5H₂O)
Preparation

2Na₂S + 3SO₂ → 2Na₂S₂O₃ + S
Structure –

HALOGENS

1. INTRODUCTION- Group VII A or group 17 of the periodic table consists of five elements Fluorine (F), chlorine (Cl), bromine (Br), iodine (I) and artificially produced astatine (At)

2. GENERAL PHYSICAL PROPERTIES OF HALOGEN FAMILY:
Physical State – Halogens are diatomic covalent molecules are held together by weak van der Waals forces.

Ionization energy – On moving down this group ionization energy decreases.

Electronegativity – On moving down this group, electronegativity keeps on decreasing due to increase in atomic size.

Basic character – of halogens increases from fluorine to iodine.

Electron affinity – The decreasing order of electron affinity of halogens is : Cl > F > Br > I

Density – Density increases gradually as we move down the group,

Solubility – Solubility of halogens in water decrease from F₂ to I₂

Non-metallic character – On moving down in this group, non-metallic character goes on decreasing.

Bond energy – Bond energy decreases on moving down the group except for F – F bond which is exceptionally low.

Oxidizing character -Oxidizing character of halogens decreases down the group.

Colour – The colour deepens down the group.

3. GENERAL CHEMICAL PROPERTIES OF HALOGEN FAMILY:
Reaction with water:

Halides of metals:
The ionic nature of M – X bond M – F > M – Cl > M – Br > M – I.
Reducing power decreases in the order I^– > Br^– > Cl^– > F^–

Halides of non-metal
M – F > M – Cl > M – Br > M – I.
HI > HBr > HCl > HF

Reactivity: The reactivity of halogens goes on decreasing on moving down the group.

Oxidation State
Hydrides – H₂ + X₂ → 2HX (X = F, Cl, Br or I)
The dipole moments of these hydrogen halides in decreasing order are
HF > HCl > HBr > HI
Thermal stability of the hydrides decreases from HF to HI i.e.
HF > HCl > HBr > HI.
Reducing properties – HF < HCl < HBr < HI.
Acidic strength – The acidic strength of halogen decreases from HI > HBr > HCl > HF.
Conjugate base strengths of these acids increase in the order
I^– < Br^– < Cl^– < F^–.
Boiling point or volatility – Volatility decreases in the order:
HCl > HBr > HI > HF.

Reaction with Alkalies:

Oxides:
(i) Oxygen difluoride, OF₂:
2F₂ + 2NaOH → 2NaF + H₂O + F₂O
(ii) Oxygen mono fluoride, O₂F₂
(iii) Chlorine mono oxide, Cl₂O:
HgO + 2Cl₂ → Cl₂O + HgCl₂

(iv) Chlorine dioxide, ClO₂:

(v) Dichloro hexoxide, Cl₂O₆: –
(vi) Dichloro heptoxide, Cl₂O₇ –

(vii) Iodine pentoxide, I₂O₅:

4. CHEMISTRY OF FLUORINE:
Occurrence:

• Fluorspar (Fluorite)CaF₂
• Cryolite Na₃AlF₆
• Fluorapatite CaF₂, 3Ca₃ (PO₄)₂

Chemical properties of fluorine

5. CHEMISTRY OF CHLORINE:
Occurrence:

• Horn silver, AgCl;
• Camallite, KCl, MgCl₂, 6H₂O
• Chlorapatite, 3Ca₃(PO₄)₂.CaCl₂
• Sylvine (Potassium chloride), KCl

Manufacture –
(a) By the electrolysis of Brine in Nelson cell:
Most of the Cl₂ used for commercial purpose is obtained as the by-product in the manufacture of caustic soda by electrolysis of brine solution.

(b) Weldon’s process:
MnO₂ + 4 HCl → MnCl₂ + 2H₂O + Cl₂

(c) Deacon’s process:

(d) Nitrosyl chloride process :
3NaCl + 4HNO₃ → 3NaN₃ + NOCl + Cl₂ + 2H₂O

Chemical properties of chlorine

6. CHEMISTRY OF BROMINE:

Occurrence:

• Sea water contains bromides such as NaBr, MgBr₂, KBr, CaBr₂ etc.
• Mother liquor left after the crystallization of KCl from camallite, KCl. MgCl₂. 6H₂O contains some bromo camallite. KBr. MgBr₂. 6H₂O
• In mineral springs and salt lakes NaBr, MgBr₂ are found in small amounts
• Bromargyrite: – a mineral AgBr.

Chemical properties of Bromine
In chemical properties, bromine closely resembles chlorine

7. CHEMISTRY OF IODINE:
Preparation of iodine:
2KI + Cl₂ → 2KCl + I₂

Physical properties of iodine:

Chemical properties of iodine:
Iodine is chemically less reactive in comparison to chlorine and bromine

8. HYDROHALIC ACIDS:
Covalent character : HI > HCl > HBr > HF
Stability: The HX bond strength decreases from HF to HI.
Acidic strength: The relative strength increases from HF to HI
Reducing property: Increases from HF to HI as the stability decreases from HF to HI.
Action of ammonia:

Action of halogens:
2HCl + F₂ → HF + Cl₂
2HBr + Cl₂ → 2HCl + Br₂
2 HI + Br₂ → 2 HBr + I₂

Acidic properties:
On reaction with certain metals, their oxides, carbonates, and bicarbonates, they form salt

9. OXY ACIDS OF CHLORINE:
Hypo-chlorous acid, (HCIO)
Properties – It is unstable and decomposes as
2HClO → 2HCl + O₂
Chlorous acid, (HClO₂)

Preparation:
Ba(ClO₂)₂ + H₂SO₄ → BaSO₄ + 2HClO₂

Properties: It shows disproportionation reaction
2HClO₂ → HClO + HClO₃
Structure:

Chloric Acid, (HClO₃)

Preparation:
Ba(ClO₃)₂ + H₂SO₄ → BaSO₄ + 2HClO₃
Structure:

Per-chloric acid (HClO₄)
Preparation – KClO₄ + H₂SO₄ → KHSO₄ + HClO₄

Structure:

10. BLEACHING POWDER, (CaOCl₂)
Manufacture:
Ca(OH)₂ + Cl₂ → CaOCl₂ + H₂O

Physical properties of bleaching powder:
CaOCl₂ → Ca²⁺ + Cl^– + OCl^–

Chemical properties of bleaching powder
HClO → HCl + O
due to presence of nascent oxygen it has oxidising and bleaching properties.

NOBLE GASES
1. INTRODUCTION – The zero group or 18^th group of the periodic table consists of six gaseous elements Helium (He), Neon (Ne), Argon (Ar), Krypton (Kr), Xenon (Xe) and Radon (Rn).

2. OCCURRENCE OF NOBLE GASES: Air

3. EXTRACTION OR ISOLATION OF NOBLE GASES
Physical method – Fractional distillation of liquid air:
This method is based on difference in boiling points of all gases

Chemical method:
Dewar’s coconut charcoal adsorption method:
(a) This method is based on the principle that the adsorption capacity of these gases at low temp, increases with their atomic weights.

(b) Adsorption capacity depends upon temperature also i.e., the lower the atomic weight of the noble gas, the lower is the temperature needed to absorb it.

4. PROPERTIES OF NOBLE GASES
Physical state:
These gases are colourless, odourless and tasteless

Heating with air:
These gases do not bum nor support in burning

Atomic radii:
Their atomic radii increase with the increase in atomic number due to addition of a new shell

Ionization energy:
Their ionisation energy is too high due to stable electronic configuration

Atomicity:
Since their combining capacity is very low they exist as a single atom hence they are monoatomic. The specific heat ratio for all these gases is 1.66

Electron affinity:
Due to complete octet, noble gases have zero electron affinity

Ease of liquefication:
The vander waals forces between the atoms increases from He to Xe due to increase in polarizability.

Solubility in water:
Solubility in water increases with increasing atomic number

Adsorption by wood charcoal:
Except He, adsorbability of these gases increases considerably with the increase in their atomic weights

Characteristic spectra:
These gases have their characteristic spectra.

Electrical conductivity:
These gases have fairly high electrical conductivity.

Xenon hexafluoride (XeF₆)

Preparation:

Structure:

Xenon tetrafluoride (XeF₄)
Preparation:

Properties:
4NO + XeF₄ → 4NOF + Xe
4NO₂ + XeF₄ → 4NO₂F + Xe

Structure:

Xenon difluoride(XeF₂)
Preparation:

Structure:

Summary of Stable Compounds of Xe

Are you looking for all chemistry formulas in one place? Then, chemistrycalc.com is a reliable and trusted website. Make use of this site and learn all concepts of chemistry formulas using provided sheets, tables & lists.