In oxides of halogen, the bonds are mainly covalent due to small difference in electronegativity between the halogens and oxygen: the bond polarity, however, increases as we move from F to I. The stability of oxides of iodine is greater than those of chlorine while bromine oxides are the least stable. For fluorine F this is the only possible oxidation number, as fluorine is the most electronegative element and doesn't loose electrons in any chemical reactions. Other halogens can have oxidation numbers: -1, +1, +3, +5 and +7.

Preparation of Hydrogen Chloride

Muriatic acid is prepared by warming NaCl crystals with concentrated H2SO4 (Sulphuric acid).

NaCl+H₂SO₄→NaHSO₄+HCl

Usually, most of the hydrogen chloride/hydrochloric acid that is formed is a co-product of some other chemical reactions. HCl is also formed by the chlorination of hydrocarbons.

Properties of Hydrogen Chloride

• HCl is an uncoloured gas and has a pungent aroma.

• Hydrochloric acid is the aqueous solution of hydrogen chloride.
• HCl is soluble in water.
• It liquefies at 189K to form a colourless liquid and freezes at 159k to form a white solid.

Uses of Hydrogen Chloride

• HCl is used in the preparation of chlorine, aqua regia, and other chlorides.
• It is used as a solvent to dissolve noble gases.
• It acts as a reagent in laboratories.

Hydrochloric acid is an inorganic chemical. It is a strong corrosive acid with a chemical formula HCl. It is also known as hydrogen chloride or muriatic acid.

Chemical properties of chlorine gas

1. Effect on litmus: Dry chlorine gas has no effect on litmus but the moist chlorine do have the effect, as it turns blue litmus red due to formation of HCl.
2. Reaction with metals and non metals: It reacts with metals and non metals to form respective chlorides that is given below:

Uses (Chlorine)

• It is used to get rid of the smell of putrefaction
• It is used as a disinfectant
• Chlorine is used in the treatment of drinking water to kill bacteria
• It is used to clean swimming pools
• It is used in the production of paper and paper products
• It is used as an antiseptic
• It is used to produce drugs
• It is used in the manufacture of dyes and plastics

The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group. Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups.

Introduction
Halogens form diatomic molecules (of the form X2​, where X denotes a halogen atom) in their elemental states. The bonds in these diatomic molecules are non-polar covalent single bonds. However, halogens readily combine with most elements and are never seen uncombined in nature. As a general rule, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form Group 1 salts with similar properties. In these compounds, halogens are present as halide anions with charge of -1 (e.g. Cl-, Br-, etc.). Replacing the -ine ending with an -ide ending indicates the presence of halide anions; for example, Cl- is named "chloride." In addition, halogens act as oxidizing agents—they exhibit the property to oxidize metals. Therefore, most of the chemical reactions that involve halogens are oxidation-reduction reactions in aqueous solution. The halogens often form single bonds, when in the -1 oxidation state, with carbon or nitrogen in organic compounds. When a halogen atom is substituted for a covalently-bonded hydrogen atom in an organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro-, chloro-, bromo-, or iodo- can be used for specific halogen substitutions. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds.

Chlorine (Cl2) was the first halogen to be discovered in 1774, followed by iodine (I2), bromine (Br2), fluorine (F2), and astatine (At, discovered last in 1940). The name "halogen" is derived from the Greek roots hal- ("salt") and -gen ("to form"). Together these words combine to mean "salt former", referencing the fact that halogens form salts when they react with metals. Halite is the mineral name for rock salt, a natural mineral consisting essentially of sodium chloride (NaCl). Lastly, the halogens are also relevant in daily life, whether it be the fluoride that goes in toothpaste, the chlorine that disinfects drinking water, or the iodine that facilitates the production of thyroid hormones in one's body.

The halogens are located on the left of the noble gases on the periodic table. These five toxic, non-metallic elements make up Group 17 of the periodic table and consist of: fluorine (F), chlorine (Cl), bromine (Br), iodine (I), and astatine (At). Although astatine is radioactive and only has short-lived isotopes, it behaves similar to iodine and is often included in the halogen group. Because the halogen elements have seven valence electrons, they only require one additional electron to form a full octet. This characteristic makes them more reactive than other non-metal groups. Halogens form diatomic molecules (of the form X₂​, where X denotes a halogen atom) in their elemental states. The bonds in these diatomic molecules are non-polar covalent single bonds. However, halogens readily combine with most elements and are never seen uncombined in nature. As a general rule, fluorine is the most reactive halogen and astatine is the least reactive. All halogens form Group 1 salts with similar properties. In these compounds, halogens are present as halide anions with charge of -1 (e.g. Cl^-, Br^-, etc.). Replacing the -ine ending with an -ide ending indicates the presence of halide anions; for example, Cl^- is named "chloride." In addition, halogens act as oxidizing agents—they exhibit the property to oxidize metals. Therefore, most of the chemical reactions that involve halogens are oxidation-reduction reactions in aqueous solution. The halogens often form single bonds, when in the -1 oxidation state, with carbon or nitrogen in organic compounds. When a halogen atom is substituted for a covalently-bonded hydrogen atom in an organic compound, the prefix halo- can be used in a general sense, or the prefixes fluoro-, chloro-, bromo-, or iodo- can be used for specific halogen substitutions. Halogen elements can cross-link to form diatomic molecules with polar covalent single bonds.

Fluorine differs from rest of the members of its group because of its small size, high electro negativity and non availability of d orbitals in the valence shell.

The halogen are reactive because of the outermost electrons of its shell, all atom wants to complete the electrons of the outermost shell with eight electrons or two. For the halogens it has 1 electron at it’s outermost shell and it wants to get rid of it that means it loses electrons easily to complete the octet rule. When it loses electrons easily, other atoms reacts or gains the electrons easily. Remember that gaining electrons releases energy.

 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.