A lone pair is the  valence electron pair without bonding or withut sharing with other atoms. They are found in the outermost electron shell of an atom, so the lone pairs are a subset of a molecule's valence electrons. They can be identified by examining the outermost energy level of an atom—lone electron pairs consist of paired electrons as opposed to single electrons, which may appear if the atomic orbital is not full. Electron pairs are therefore considered lone pairs if two electrons are paired but are not used in chemical bonding. Thus, the number of lone electrons plus the number of bonding electrons equal the total number of valence electrons in a compound.

The strength of the chemical bond between any two species can be estimated with the help of bond dissociation enthalpy. Although it is generally measured as the enthalpy change at standard conditions (298K), the bond dissociation energy of a chemical bond is often defined as the enthalpy change of the homolytic fission of the bond at absolute zero (0K).

Some important features of the concept of bond dissociation enthalpy include:

• It is the amount of energy which needs to be supplied in order to break a chemical bond between two species.
• It is a means of calculating the strength of a chemical bond.
• In diatomic molecules specifically, it is equal to the value of bond energy
• The bond between silicon and fluorine is said to have the strongest bond dissociation enthalpy.
• Covalent bonds between atoms or molecules are said to have weak bond dissociation energies.

Halogens are coloured because of their low ionisation energy.

As a result, the electrons get excited in the visible region of light.

The remaining light falls in the visible region and halogens appear to be coloured.

For example: Fluorine absorbs wavelengths corresponding to violet light and the remaining light appears to be yellow. Thus, fluorine has a yellow colour. Similarly, chlorine is greenish yellow, bromine reddish and iodine is a violet solid.

Halogens are coloured because of their low ionisation energy.

As a result, the electrons get excited in the visible region of light.

The remaining light falls in the visible region and halogens appear to be coloured.

For example: Fluorine absorbs wavelengths corresponding to violet light and the remaining light appears to be yellow. Thus, fluorine has a yellow colour. Similarly, chlorine is greenish yellow, bromine reddish and iodine is a violet solid.

Halogens are coloured because of their low ionisation energy.

As a result, the electrons get excited in the visible region of light.

The remaining light falls in the visible region and halogens appear to be coloured.

For example: Fluorine absorbs wavelengths corresponding to violet light and the remaining light appears to be yellow. Thus, fluorine has a yellow colour. Similarly, chlorine is greenish yellow, bromine reddish and iodine is a violet solid.

The molecules of each substance attract each other through dispersion (London) intermolecular forces. Whether a substance is a solid, liquid, or gas depends on the balance between the kinetic energies of the molecules and their intermolecular attractions
In fluorine, the electrons are tightly held to the nuclei. The electrons have little chance to wander to one side of the molecule, so the London dispersion forces are relatively weak
As we move from fluorine to iodine, the electrons are further from the nuclei so the electron clouds can more easily distort. The London dispersion forces become progressively stronger.
At a low enough temperature, the molecules will all be solids. At a high enough temperature, they will all be gases.
It is only at a temperature  between −7oC and 59o that fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.

The molecules of each substance attract each other through dispersion (London) intermolecular forces. Whether a substance is a solid, liquid, or gas depends on the balance between the kinetic energies of the molecules and their intermolecular attractions
In fluorine, the electrons are tightly held to the nuclei. The electrons have little chance to wander to one side of the molecule, so the London dispersion forces are relatively weak
As we move from fluorine to iodine, the electrons are further from the nuclei so the electron clouds can more easily distort. The London dispersion forces become progressively stronger.
At a low enough temperature, the molecules will all be solids. At a high enough temperature, they will all be gases.
It is only at a temperature  between −7^oC and 59^o C that fluorine and chlorine are gases, bromine is a liquid and iodine is a solid.

1. Oxidation states: They show variable oxidation states like:

Fluorine: -1

Chlorine: -1, +1, +3, +7

Bromine : -1, +1, +3, +5, +7

Iodine:  -1, +1, +3, +5, +7

• Higher oxidation states of halogens are used when they are combining with small size highly electronegative ions.
• All halogens are very reactive and reactivity decreases down the group.
• All act as Lewis acids as they accept electron.
• Fluorine is the strong oxidizing agent among all.

Preparation of Sulphur Dioxide

1. In the laboratory, sulphur dioxide is prepared by the reaction of metallic sulphite or a metallic bisulphite with dilute acid. For example, a reaction between the dilute sulphuric acid and sodium sulphite will result in the formation of SO₂.

Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂

2. Commercially it is obtained as a by-product released from the roasting of sulphide ores. The gas obtained is dried, liquefied and then stored in steel cylinders.

4FeS₂ (s) + 11 O₂ (g) →2Fe₂O₃ (s) + 8SO₂ (g)

Electron gain enthalpy:

• Along period it is more negative because of attraction towards electron as of increased nuclear charge.
• Down the group electron gain enthalpy keeps on becoming less negative because nuclear charge decreases and size.
• So, the group 17 has highest electron gain enthalpy due to smallest size in periodic table.
• Please note that: Fluorine has less electron gain enthalpy than chlorine because due to small size of oxygen the incoming electron suffers repulsion therefore, electron gain enthalpy is less negative as compared to chlorine.

Fluorine is the most electronegative element and cannot exhibit any positive oxidation state. Other halogens have d orbitals and therefore, can expand their octets and show + 1, + 3, + 5 and + 7 oxidation states also.

Fluorine is the most electronegative element and cannot exhibit any positive oxidation state. Other halogens have d orbitals and therefore, can expand their octets and show + 1, + 3, + 5 and + 7 oxidation states also.

Oxidation state indicates the degree of oxidation for an atom in a chemical compound; it is the hypothetical charge that an atom would have if all bonds to atoms of different elements were completely ionic. Oxidation states are typically represented by integers, which can be positive, negative, or zero.

1. The oxidation number of a free element is always 0.
2. The oxidation number of a monatomic ion equals the charge of the ion.
3. The oxidation number of H is +1, but it is -1 in when combined with less electronegative elements.
4. The oxidation number of O in compounds is usually -2, but it is -1 in peroxides.

The nuclear and atomic radii of these elements keep on increasing as we move down the group. This happens because of the addition of an extra energy level. They have the minimal atomic radii compared to the other elements in the related periods. This can be attributed to the fact that their atomic charge is quite powerful.

• Physical state: The group 17 elements are found in diverse physical states. For example, Fluorine and Chlorine are gases. On the other hand, Bromine is a liquid and Iodine is solid.
• Colour: These elements have a variety of colours. For example, while Fluorine is pale yellow in colour, Iodine is dark violet in colour.
• Solubility: Florine and Chlorine are soluble in water. On the other hand, Bromine and Iodine are very less soluble in water.
• Melting and boiling points: Melting and boiling points of these elements increase as we move down the group from Fluorine to Iodine. Thus, Fluorine has the lowest boiling and melting points.

1. Electronic configuration-The general electronic configuration for this group is ns²np⁵

• Fluorine (F) [He]2s²,2p⁵
• Chlorine (Cl) [Ne]3s²,3p⁵
• Bromine(Br)[Ar]3d¹⁰4s²4p⁵
• Iodine (I)[Kr]4d^10,5s²,5p⁴
• Astatine (At) [Xe]4f¹⁴,5d¹⁰,6s²,6p⁴

The members of group 17 are:

• Fluorine(F) - 9
• Chlorine(Cl) - 17 
• Bromine(Br) - 35
• Iodine (I) - 53
• Astatine (At) - 85

The members of group 17 are:

• Fluorine(F)
• Chlorine(Cl)
• Bromine(Br)
• Iodine (I)
• Astatine (At)

The halogens are the elements that form group 17 of the periodic table. They are reactive nonmetals and include fluorine, chlorine, bromine, and iodine.

Nascent hydrogen is most powerful reducing agent than ordinary hydrogen because nascent hydrogen is in atomic state and atoms are more active than molecules, nascent hydrogen is evolved in small bubbles containing the gas under great pressure, nascent hydrogen is activated by the energy liberated in the reaction in which it is formed, so the nascent hydrogen becomes more energized and active.

1. In the laboratory, sulphur dioxide is prepared by the reaction of metallic sulphite or a metallic bisulphite with dilute acid. For example, a reaction between the dilute sulphuric acid and sodium sulphite will result in the formation of SO₂.

Na₂SO₃ + H₂SO₄ → Na₂SO₄ + H₂O + SO₂

2. Commercially it is obtained as a by-product released from the roasting of sulphide ores. The gas obtained is dried, liquefied and then stored in steel cylinders.

4FeS₂ (s) + 11 O₂ (g) →2Fe₂O₃ (s) + 8SO₂ (g)

The space in a molecule in which the probability of finding an electron is maximum can be calculated using the molecular orbital function. Molecular orbitals are basically mathematical functions that describe the wave nature of electrons in a given molecule.

These orbitals can be constructed via the combination of hybridized orbitals or atomic orbitals from each atom belonging to the specific molecule. Molecular orbitals provide a great model via the molecular orbital theory to demonstrate the bonding of molecules.

The temperature at which electrical resistivity of the material suddenly drops of zero and the material changes from normal conductor to a super conductor is called the Transition temperature or critical temperature T C ​ .

Sulphur forms numerous allotropes, but let us study the two most important allotropes of sulphur-

yellow rhombic sulphur (α-sulphur) and the monoclinic (β-sulphur). The most interesting feature is their thermal stability, the allotropes of sulphur are inter-convertible i.e. rhombic sulphur when heated above 369K gives monoclinic sulphur. 

Rhombic sulphur (α-sulphur)

Rhombic sulphur is crystalline in nature and has octahedral shape. On heating the solution of roll sulphur in CS2 we get rhombic sulphur. It is yellow with a melting point of 385.8K and specific gravity 2.06. Rhombic sulphur cannot be dissolved in water but can be dissolved in benzene, ether, alcohol etc.

Monoclinic sulphur (β-sulphur)

When we take a dish and melt rhombic sulphur in that dish we obtain monoclinic sulphur after cooling it. In this process we make two holes in the crust and pour out the remaining liquid. After this we get colourless needle-shaped crystals of β-sulphur when the crust is removed.