p-Block Elements Class 12 Notes CBSE Chemistry Chapter 7 | Free PDF Download

Group 15 Elements: The Nitrogen Family

Nitrogen phosphorus, arsenic, antimony, and bismuth are all members of Group 15. There is a transition from nonmetallic to metallic through metalloid property as we move down the group. Non-metals are nitrogen and phosphorus, metalloids are arsenic and antimony, and bismuth is a typical metal.

Electronic Configuration

The electrical configuration of these elements' valence shells is \[{\text{n}}{{\text{s}}^2}{\text{n}}{{\text{p}}^3}\]. The s orbital in these elements is totally filled, whereas the p orbitals are half-filled. It results in an electronic configuration extremely stable structure

Atomic and Ionic Radii

Radii of covalent and ionic (in a certain state) compounds grow in larger as they progress through the group. From N to P, the covalent radius increases dramatically. However, there is only a modest increase in covalent radius from As to Bi. This is owing to the presence of totally filled d and / or f orbitals in heavier members.

Ionisation Enthalpy

Due to the progressive rise in atomic size, the ionisation enthalpy drops down the group. The ionisation enthaply of group 15 elements is substantially greater than that of group 14 elements in the equivalent periods due to the extra stable half-filled p-orbital electronic configuration and smaller size. As expected, the order of consecutive ionisation enthalpies is \[{\Delta _1}{{\text{H}}_1} < {\Delta _1}{{\text{H}}_2} < {\Delta _1}{{\text{H}}_3}\].

Electronegativity

With increasing atomic size, the electronegativity value generally drops down the group. However, the gap is less significant among the heavier elements.

Physical Properties

This group's elements are all polyatomic. All other elements are solids except dinitrogen, which is a diatomic gas. The group's metallic aspect grows stronger as it progresses. Bismuth is a metalloid, while nitrogen and phosphorus are non-metals. Arsenic and antimony are metalloids, and bismuth is a metal. This is owing to an increase in atomic size and a decrease in ionisation enthalpy. In general, boiling points rise from top to bottom in the group, although melting points rise until arsenic and then fall until bismuth. Except for nitrogen, all elements exhibit allotropy.

Atomic & Physical Properties

Chemical Properties

State of Oxidation and Developments in Chemical Reactivity

These elements' most common oxidation states are – 3, + 3, and + 5. Down the group, the likelihood to exhibit – 3 oxidation state declines, and bismuth rarely produces any –3 oxidation state compounds. The stability of the + 5 oxidation state reduces as you progress through the group. \[{\text{Bi}}{{\text{F}}_5}\] is the only \[{\text{Bi}}\] (V) compound that has been thoroughly studied. The stability of the +5-oxidation state reduces as the group progresses, while the stability of the +3 state improves (owing to the inert pair effect). When nitrogen combines with oxygen, it has oxidation states of +1, + 2, and + 4. Some oxoacids have +1 and +4 oxidation states in phosphorous. In the case of nitrogen, in acid solution, all oxidation states from +1 to +4 tend to be disproportionate. For example

\[3{\text{HN}}{{\text{O}}_3} \to {\text{HN}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O + 2NO}}\]

Similarly, in alkali and acid, practically all intermediate oxidation levels of phosphorus disproportionate into +5 and –3. However, in the case of arsenic, antimony, and bismuth, the +3-oxidation state becomes progressively stable with respect to disproportionation.

Because only four orbitals (one s and three p) are available for bonding, nitrogen can only have a maximum covalency of four. Heavy elements contain unoccupied d orbitals in their outermost shells that can be utilised for bonding (covalency) and therefore enlarge their covalence, as in \[{\text{P}}{{\text{F}}_6}^- \]

Anomalous Properties of Nitrogen

Nitrogen is distinguished from the other members of this family by its tiny size, strong electronegativity, high ionisation enthalpy, and lack of d orbitals. Nitrogen has a one-of-a-kind ability to create p–p multiple bonds with itself and other tiny, electronegativity-rich elements \[({\text{C, O}})\]Because the atomic orbitals of heavier elements in this group are too vast and diffuse to overlap effectively, they do not form p – p bonds. Thus, nitrogen is a diatomic molecule having a triple bond between the two atoms (one s and two p). As a result, the bond enthalpy \[(941.1{\text{ KJ mo}}{{\text{l}}^{ - 1}})\] is condition. Because of the significant interelectronic repulsion of the non-bonding electrons and the short bond length, the single N – N bond is weaker than the single P – P connection. As a result, in nitrogen, the catenation tendency is weaker. The absence of d orbitals in nitrogen's valence shell is another aspect that impacts its chemistry. Nitrogen cannot form d – p bonds like the heavier elements, e.g., \[{{\text{R}}_3}{\text{P = 0}}\], since its covalency is limited to four. When their compounds like \[{\text{P(}}{{\text{C}}_2}{{\text{H}}_5}{)_3}\] and \[{\text{As(}}{{\text{C}}_6}{{\text{H}}_5}{)_3}\] act as ligands, phosphorus and arsenic can form d – p bonds with transition metals extremely high. Phosphorus, arsenic, and antimony, on the other hand, form metallic connections in their elemental.

(i) Reactivity Towards Hydrogen

All the elements in Group 15 produce ${\text{E}}{{\text{H}}_{\text{3}}}$ hydrides, where \[{\text{E = N, P, As, Sb or Bi}}\]. Table shows some of the features of these hydrides. The characteristics of the hydrides show a steady progression. The bond dissociation enthalpy of hydrides decreases from ${\text{N}}{{\text{H}}_{\text{3}}}$ to ${\text{Bi}}{{\text{H}}_{\text{3}}}$, indicating that their stability decreases. As a result, the hydrides' reducing nature improves. Ammonia is a weak reducing agent, but BiH3 is the most powerful hydride reducing agent. The order of basicity also declines.

\[{\text{N}}{{\text{H}}_3}{\text{ > P}}{{\text{H}}_3}{\text{ > As}}{{\text{H}}_3}{\text{ > Sb}}{{\text{H}}_3} \geqslant {\text{Bi}}{{\text{H}}_3}\]

Properties of Hydrides of Group 15 Elements

(ii) Reactivity Towards Oxygen

\[{{\text{E}}_2}{{\text{O}}_3}\] and \[{{\text{E}}_2}{{\text{O}}_5}\] are the two forms of oxides formed by these elements. The oxide in the element's higher oxidation state is more acidic than the oxide in the lower oxidation state. Their acidity lowers as they progress through the group. Nitrogen and phosphorus oxides are entirely acidic, arsenic and antimony oxides are amphoteric, and bismuth oxides are largely basic.

(iii) Reactivity Towards Halogens

These elements combine to generate two halide series: ${\text{E}}{{\text{H}}_{\text{3}}}$ and ${\text{E}}{{\text{X}}_5}$. Because the d orbitals in nitrogen's valence shell are not available, it does not form pentahalide. Pentahalides have a higher covalent bonding strength than trihalides. Except for nitrogen's trihalides, all these elements' trihalides are stable. Only \[{\text{N}}{{\text{F}}_3}\] is known to be stable in the situation of nitrogen. Except for \[{\text{Bi}}{{\text{F}}_3}\], trihalides are generally covalent in nature.

(iv) Reactivity Towards Metals

These elements combine with metals to generate binary compounds having an oxidation state of –3 such as \[{\text{C}}{{\text{a}}_3}{{\text{N}}_2}\](Calcium nitride), \[{\text{C}}{{\text{a}}_3}{{\text{P}}_2}\](calcium phosphide), \[{\text{N}}{{\text{a}}_3}{\text{A}}{{\text{s}}_2}\] (sodium arsenide), \[{\text{Z}}{{\text{n}}_3}{\text{S}}{{\text{b}}_2}\](zinc antimonide), and \[{\text{M}}{{\text{g}}_3}{\text{B}}{{\text{i}}_2}\] (magnesium bismuthide).

Nitrogen (n):

(i) A combination of \[{\text{N}}{{\text{H}}_4}{\text{Cl}}\] and \[{\text{NaN}}{{\text{O}}_2}\] is heated. The downward displacement of water collects \[{{\text{N}}_2}\]

\[{\text{N}}{{\text{H}}_4}{\text{Cl + NaN}}{{\text{O}}_2}\xrightarrow{\Delta }{\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_2} + {\text{NaCl; N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_2}\xrightarrow{\Delta }{{\text{N}}_2} \uparrow  + 2{{\text{H}}_2}{\text{O}}\]

An aqueous solution of ammonium chloride is treated with sodium nitrate. It is a process of preparation that is done in a laboratory.

\[{\text{N}}{{\text{H}}_4}{\text{Cl(aq) + NaN}}{{\text{O}}_2}({\text{aq)}} \to {{\text{N}}_2}({\text{g)}} + {{\text{H}}_2}{\text{O(}}\ell {\text{) + NaC}}{{\text{l}}_{({\text{aq)}}}}\]

(ii) By heating ammonium dichromate:

\[{{\text{(N}}{{\text{H}}_4}{\text{)}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_7}\xrightarrow{\Delta }{{\text{N}}_2} \uparrow  + 4{{\text{H}}_2}{\text{O + C}}{{\text{r}}_2}{{\text{O}}_3}\]

(iii) By oxidation of ammonia

(A) At lower temperature

\[{\text{8N}}{{\text{H}}_3}{\text{(}}\ell {\text{) + 3C}}{{\text{l}}_2}({\text{g) }} \to 6{\text{N}}{{\text{H}}_4}{\text{Cl}} + {{\text{N}}_2} \uparrow \]

If there is an excess of \[{\text{C}}{{\text{l}}_2}\] in this reaction, nitrogen trichloride is produced as shown below.

\[{\text{N}}{{\text{H}}_3}{\text{ + 3C}}{{\text{l}}_2} \to {\text{NC}}{{\text{l}}_3} + 3{\text{HCl}}\]

Nitrogen trichloride is a very explosive gas.

(b) Ammonia is produced by reacting calcium hypochlorite or \[{\text{B}}{{\text{r}}_2}\] with ammonia.

\[{\text{4N}}{{\text{H}}_3}{\text{ + 3Ca(OCl}}{{\text{)}}_2} \to 3{\text{CaC}}{{\text{l}}_2}{\text{ + }}{{\text{N}}_2} + {{\text{H}}_2}{\text{O}}\]

(B) At Higher Temperature

Ammonia is passed through heated cupric oxide or PbO.

\[{\text{2N}}{{\text{H}}_3}{\text{ + 3CuO}} \to {\text{ + }}{{\text{N}}_2} \uparrow  + 3{\text{Cu}} + 3{{\text{H}}_2}{\text{O}}\]

(iv) In the presence of dilute \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\], heating urea with a nitrite

\[{\text{N}}{{\text{H}}_2}{\text{CON}}{{\text{H}}_2}{\text{ + 2NaN}}{{\text{O}}_2}{\text{ + }}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\xrightarrow{\Delta }{\text{ + N}}{{\text{a}}_2}{\text{S}}{{\text{O}}_4} + 2{{\text{N}}_2} \uparrow  + 3{{\text{H}}_2}{\text{O + C}}{{\text{O}}_2} \uparrow \]

(v) Using sodium hypobromite to heat urea solution:

\[{\text{N}}{{\text{H}}_2}{\text{CON}}{{\text{H}}_2}{\text{ + 3NaOB}}{{\text{r}}_2}\xrightarrow{\Delta }{{\text{N}}_2} \uparrow  + 3{\text{NaBr + 2}}{{\text{H}}_2}{\text{O + C}}{{\text{O}}_2} \uparrow \]

(vi) Nitric oxide and\[{\text{N}}{{\text{H}}_3}\] are passed via red hot copper gauze.

\[{\text{4N}}{{\text{H}}_3} + 6{\text{NO}} \to 5{{\text{N}}_2} \uparrow  + 6{{\text{H}}_2}{\text{O}}\]

(vii) Using \[{\text{HN}}{{\text{O}}_3}\] vapours to cool red-hot copper

\[5{\text{Cu}} + 2{\text{HN}}{{\text{O}}_3} \to 5{\text{CuO + }}{{\text{N}}_2} \uparrow  + {{\text{H}}_2}{\text{O}}\]

(viii) Very pure nitrogen \[{\text{Ba(}}{{\text{N}}_3}{)_2}\xrightarrow{\Delta }{\text{Ba + 3}}{{\text{N}}_2}\]

On heating, sodium azide also produces \[{{\text{N}}_2}\].

Industrial Methods of Preparation:

(i) Fractional distillation from liquefied air: Because the boiling points of \[{{\text{N}}_2}\] and oxygen are \[ - {190^o}{\text{C}}\] and \[ - {183^o}{\text{C}}\], they can be separated.

(ii) From furnace gas produced by producers: \[{\text{CO}}\]and \[{{\text{N}}_2}\] are mixed in the producer gas. When \[{\text{CO}}\] and \[{{\text{N}}_2}\] are passed over heated \[{\text{CuO}}\], the \[{\text{CO}}\] gas is oxidised to \[{\text{C}}{{\text{O}}_2}\], which is absorbed by alkalies, while the \[{{\text{N}}_2}\] gas is collected in gas cylinders.

Properties:

(i) \[{{\text{N}}_2}\] is a colourless, odourless gas that is water insoluble.

(ii) It is a neutral, non-polar covalent molecule.

(iii) It is neither flammable nor combustible-supporting.

(iv) It is absorbed by \[{\text{Mg}}\] and \[{\text{Al}}\] when they are heated. As a result, the nitrides generated react with water to form \[{\text{N}}{{\text{H}}_3}\].

$ 3{\text{Mg + }}{{\text{N}}_2} \to {\text{M}}{{\text{g}}_3}{{\text{N}}_2}( + \ 6{{\text{H}}_2}{\text{O)}} \to {\text{3Mg(OH}}{{\text{)}}_3} + 2{\text{N}}{{\text{H}}_3} \uparrow  \ \\ $

 $ 2{\text{Al + }}{{\text{N}}_2} \to 2{\text{AlN}} \ ( + \ 6{{\text{H}}_2}{\text{O)}} \to 2{\text{Al(OH}}{{\text{)}}_3} + 2{\text{N}}{{\text{H}}_3} \uparrow  \ \\ $

(v) Reaction with \[{{\text{H}}_2}\]: In the presence of an iron catalyst and a molybdenum promoter, \[{{\text{N}}_2}\] reacts reversibly with \[{{\text{H}}_2}\] to generate ammonia at 200 atm and \[{500^0}{\text{C}}\]. The procedure is known as Haber's Process, and it is used in the industrial production of ammonia. It's an exothermic process.

\[{{\text{N}}_2} + 3{{\text{H}}_2} \to 2{\text{N}}{{\text{H}}_3}\]

(vi) Reaction with oxygen: Nitric oxide is generated when air devoid of \[{\text{C}}{{\text{O}}_2}\] and moisture is passed through an electric arc at around 2000 K. This is an endothermic process.

\[{{\text{N}}_2} + {{\text{O}}_2} \to 2{\text{NO}}\]

(vii) Reaction with \[{\text{Ca}}{{\text{C}}_2}\] and \[{\text{Ba}}{{\text{C}}_2}\]: These carbides react with \[{{\text{N}}_2}\] at \[{1100^0}{\text{C}}\] to create \[{\text{CaC}}{{\text{N}}_2}\] and \[{\text{Ba(CN}}{{\text{)}}_2}\], respectively.

\[{\text{Ca}}{{\text{C}}_2} + {{\text{N}}_2}\xrightarrow{\Delta }{\text{CaC}}{{\text{N}}_2} + {\text{C}}\](nitrolim, a fertilizer)

\[{\text{Ba}}{{\text{C}}_2} + {{\text{N}}_2}\xrightarrow{\Delta }{\text{Ba(CN}}{{\text{)}}_2}\]

In the soil,\[{\text{CaC}}{{\text{N}}_2}\] interacts with \[{{\text{H}}_2}{\text{O}}\] to form \[{\text{N}}{{\text{H}}_3}\] gas. Nitrates are formed when \[{\text{N}}{{\text{H}}_3}\] gas is transformed to nitrates by nitrating bacteria in the soil. (Nitrates are easily absorbed by plants and provide them with the nitrogen they require.)

USES:

(i) to provide an inert environment during various industrial procedures that require the absence of air or oxygen.

(ii) for the Haber process, which produces \[{\text{N}}{{\text{H}}_3}\].

(iii) for the Birkeland-Eyde process, which produces \[{\text{HN}}{{\text{O}}_3}\].

(iv) to produce nitrolim.

Compounds of Nitrogen

Ammonia

Preparation:

(i) By any base or alkali acting on any ammonium salt:

$ {\text{N}}{{\text{H}}_4}{\text{NaOH}}\xrightarrow{\Delta }{\text{N}}{{\text{H}}_3} \uparrow  + \ {\text{NaCl + }}{{\text{H}}_2}{\text{O}} \ \\ $

$ {{\text{(N}}{{\text{H}}_4})_2} + \ 2{\text{NaOH}}\xrightarrow{\Delta }2{\text{N}}{{\text{H}}_3} \uparrow  + \ {\text{N}}{{\text{a}}_2}{\text{S}}{{\text{O}}_4} + 2{{\text{H}}_2}{\text{O}} \ \\ $

$ {\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} + \ {\text{NaOH}}\xrightarrow{\Delta }{\text{N}}{{\text{H}}_3} \uparrow  + \ {\text{NaN}}{{\text{O}}_3}{\text{ + }}{{\text{H}}_2}{\text{O}} \ \\ $

$ {{\text{(N}}{{\text{H}}_4})_3}{\text{P}}{{\text{O}}_4} + \ 3{\text{NaOH}}\xrightarrow{\Delta }2{\text{N}}{{\text{H}}_3} \uparrow  + \ {\text{N}}{{\text{a}}_3}{\text{P}}{{\text{O}}_4} + \ 3{{\text{H}}_2}{\text{O}} \ \\ $

$ {{\text{(N}}{{\text{H}}_4})_2}{\text{S}}{{\text{O}}_4} + \ {\text{CaO}}\xrightarrow{\Delta }2{\text{N}}{{\text{H}}_3} \uparrow  + \ {\text{CaS}}{{\text{O}}_4} + {{\text{H}}_2}{\text{O}} \ \\ $

This is a general method for determining ammonium salts.

(ii) Metal nitrides, such as \[{\text{AlN}}\] or \[{\text{M}}{{\text{g}}_3}{{\text{N}}_2}\], are hydrolyzed.

(iii) From nitrogen oxides: When nitrogen oxides are combined with \[{{\text{H}}_2}\] and passed over a heated platinum catalyst, \[{\text{N}}{{\text{H}}_3}\] gas is produced.

$ 2{\text{NO + 5}}{{\text{H}}_2} \to 2{\text{N}}{{\text{H}}_3} \uparrow  + 2{{\text{H}}_2}{\text{O}} \ \\ $

${\text{2N}}{{\text{O}}_2} + \ 7{\text{N}}{{\text{H}}_3} \to 2{\text{N}}{{\text{H}}_3} \uparrow  + \ 4{{\text{H}}_2}{\text{O}} \ \\ $ 

(iv) From organic amides: Ammonia is produced when an organic amide is burned in a \[{\text{NaOH}}\] solution.

\[{\text{C}}{{\text{H}}_3}{\text{CON}}{{\text{H}}_2} + \ {\text{NaOH}}\xrightarrow{\Delta }{\text{C}}{{\text{H}}_3}{\text{COONa + N}}{{\text{H}}_3} \uparrow \]

(v) Nitrogen compounds (nitrates and nitrites): Ammonia is produced when a metal nitrate or nitrite is heated with zinc powder and a strong \[{\text{NaOH}}\] solution. The reactions are varied.

$  {\text{NaN}}{{\text{O}}_3} + \ {\text{7NaOH + 4Zn}} \to {\text{4N}}{{\text{a}}_2}{\text{Zn}}{{\text{O}}_2} + \ {\text{N}}{{\text{H}}_3} \uparrow  + \ 2{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{NaN}}{{\text{O}}_2} + \ 3{\text{Zn + 5NaOH}} \to 3{\text{N}}{{\text{a}}_2}{\text{Zn}}{{\text{O}}_2} + \ {{\text{H}}_2}{\text{O}} + \ {\text{N}}{{\text{H}}_3} \uparrow  \ \\ $

Thus, this reaction can identify a nitrite or a nitrate, but it cannot distinguish between the two. The ammonia produced is dried by passing it through quick lime and collected by air displacement downward. Because NH3 interacts with all of these, ammonia cannot be dried using \[{\text{CaC}}{{\text{l}}_2}\], \[{{\text{P}}_2}{{\text{O}}_5}\], or conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\].

Industrial Methods of Preparation:

(i) Haber’s process: \[{{\text{N}}_2} + \ 3{{\text{H}}_2}\xrightarrow[{{\text{Iron oxide + }}{{\text{K}}_2}{\text{O\& A}}{{\text{l}}_2}{{\text{O}}_3}}]{{{{500}^0}{\text{C,200atm}}}}2{\text{N}}{{\text{H}}_3}\]

(ii) From coal destructive distillation: Three compounds are produced when coal is burned at a high temperature in an iron retort and the distillate is bubbles in water:

(a) Tarry black pitch,

(b) Liquor ammonia

(c) Coal gas

A concentrated solution containing ammonia and ammonium salts is called liquor ammonia. It emits ammonia when heated. When all the free \[{\text{N}}{{\text{H}}_3}\] has been recovered, the leftover liquid is heated with \[{\text{Ca(OH}}{{\text{)}}_2}\] to breakdown the ammonium ions and release further ammonia.

(iii) Cyanamide process: 

$ {\text{CaO + 2C + }}{{\text{N}}_2}\xrightarrow{{{{2000}^o} {\text{C}}}}{\text{CaC}}{{\text{N}}_2} + \ {\text{CO}} \uparrow  \ \\ $

$  {\text{CaC}}{{\text{N}}_2} + 3{{\text{H}}_2}{\text{O}} \to {\text{CaC}}{{\text{O}}_3} + 2{\text{N}}{{\text{H}}_3} \uparrow  \ \\ $ 

Properties:

(i) A colourless, lighter-than-air gas. In nature, being basic turns red litmus blue. It serves as a Lewis foundation.

(ii) Water soluble to a high degree. Ammonium hydroxide solution is the name of the solution.

(iii) \[{\text{Na + N}}{{\text{H}}_3}\xrightarrow{\Delta }{\text{NaN}}{{\text{H}}_2} + 1/2{{\text{H}}_2}\]

Amides decompose back with water to form \[{\text{N}}{{\text{H}}_3}\] and \[{\text{NaOH}}\].

(iv) \[{\text{4N}}{{\text{H}}_3} + 5{{\text{O}}_2}\xrightarrow{{{\text{Pt,55}}{{\text{0}}^0}{\text{C}}}}{\text{4NO}} + 6{{\text{H}}_2}{\text{O}}\] (Ostwald's \[{\text{HN}}{{\text{O}}_3}\] manufacturing technique)

(v) Nitrogen gas is generated when Cl2 is bubbled in liquid ammonia.

\[{\text{8N}}{{\text{H}}_3} + 3{\text{C}}{{\text{l}}_2} \to {\text{6N}}{{\text{H}}_4}{\text{Cl}} + {{\text{N}}_2} \uparrow \]

It is transformed to an explosive compound, nitrogen trichloride, when there is an excess of \[{\text{C}}{{\text{l}}_2}\].

\[{\text{N}}{{\text{H}}_3} + 3{\text{C}}{{\text{l}}_2} \to {\text{NC}}{{\text{l}}_3} + {\text{3HCl}}\]

(vi) When \[{\text{N}}{{\text{H}}_3}\] is passed over heated \[{\text{CuO}}\] and \[{\text{PbO}}\], the oxides are converted to metal.

${\text{3CuO + 2N}}{{\text{H}}_3} \to 3{\text{Cu + 3}}{{\text{H}}_2}{\text{O + }}{{\text{N}}_2} \uparrow  \ \\ $

$  3{\text{PbO + 2N}}{{\text{H}}_3} \to 3{\text{Pb + 3}}{{\text{H}}_2}{\text{O + }}{{\text{N}}_2} \uparrow  \ \\ $

(vii)

${\text{CuS}}{{\text{O}}_4} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to {\text{Cu(OH}}{{\text{)}}_2} \downarrow ({\text{blue) + \ (N}}{{\text{H}}_4}{)_2}{\text{S}}{{\text{O}}_4} \ \\ $

$  {\text{Cu(OH}}{{\text{)}}_2} + \ {({\text{N}}{{\text{H}}_4})_2}{\text{S}}{{\text{O}}_4} + \ 3{\text{N}}{{\text{H}}_4}{\text{OH(excess)}} \to \left[ {{\text{Cu(N}}{{\text{H}}_3}{)_4}} \right]{\text{S}}{{\text{O}}_4}({\text{deep blue solution) + 4}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{ZnS}}{{\text{O}}_4} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to {\text{Zn(OH}}{{\text{)}}_2} \downarrow ({\text{white) + (N}}{{\text{H}}_4}{)_2}{\text{S}}{{\text{O}}_4} \ \\ $

$  {\text{Zn(OH}}{{\text{)}}_2} + \ {({\text{N}}{{\text{H}}_4})_2}{\text{S}}{{\text{O}}_4} + \ 2{\text{N}}{{\text{H}}_4}{\text{OH(excess)}} \to \left[ {{\text{Zn(N}}{{\text{H}}_3}{)_4}} \right]{\text{S}}{{\text{O}}_4}(colourless{\text{ solution) + 4}}{{\text{H}}_2}{\text{O}} \ \\ $ 

\[{\text{CdS}}{{\text{O}}_4}\] solution undergoes similar reactions.

(viii) 

${\text{M(N}}{{\text{O}}_3}{)_2} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to {\text{M(OH}}{{\text{)}}_2} \downarrow (white) + \ 2{\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} \ \\ $

$  {\text{MC}}{{\text{l}}_2} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to {\text{M(OH}}{{\text{)}}_2} \downarrow (white) + \ 2{\text{N}}{{\text{H}}_4}{\text{Cl}} \ \\ $

$  {\text{(M = Mg, Ca, Sr, Ba, Ra, Sn, Pb)}} \ \\ $ 

(ix) A brown ppt. is obtained when \[{\text{N}}{{\text{H}}_4}{\text{OH}}\] solution is introduced to \[{\text{AgN}}{{\text{O}}_3}\]solution.

\[{\text{2AgN}}{{\text{O}}_3} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to {\text{A}}{{\text{g}}_2}{\text{O}} \downarrow {\text{(brown) + 2N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}}\]

Excess ammonium hydroxide dissolves the brown portion of silver oxide, yielding a soluble complex.

\[{\text{A}}{{\text{g}}_2}{\text{O}} + 2{\text{N}}{{\text{H}}_4}{\text{OH}} \to \left[ {{\text{A}}{{\text{g}}_2}{{{\text{(N}}{{\text{H}}_3}{\text{)}}}_2}} \right]{\text{OH(colourless solution)}} + 2{{\text{H}}_2}{\text{O}}\]

Similarly, with mercuric salts, \[{\text{N}}{{\text{H}}_4}{\text{OH}}\] forms a white precipitate

\[{\text{HgC}}{{\text{l}}_2}({\text{aq}}{\text{.) + 2N}}{{\text{H}}_4}{\text{OH}} \to {\text{HgN}}{{\text{H}}_2}{\text{Cl}} \downarrow {\text{(white) + N}}{{\text{H}}_4}{\text{Cl + }}{{\text{H}}_2}{\text{O}}\]

(x) Nitrogen gas is generated when liquid ammonia is dripped on heated bleaching powder.

\[3{\text{Ca(OCl)Cl + 2N}}{{\text{H}}_3} \to 3{\text{CaC}}{{\text{l}}_2} + 3{{\text{H}}_2}{\text{O + }}{{\text{N}}_2} \uparrow \]

(xi) 

${\text{2N}}{{\text{H}}_3} + {\text{C}}{{\text{O}}_2} + {{\text{H}}_2}{\text{O}} \to {({\text{N}}{{\text{H}}_4})_2}{\text{C}}{{\text{O}}_3} \ \\ $

$  {\text{2N}}{{\text{H}}_3} + {\text{C}}{{\text{O}}_2}\xrightarrow[\Delta ]{{{\text{high pressure}}}}{\text{N}}{{\text{H}}_2}{\text{CON}}{{\text{H}}_2}({\text{urea)}} + {{\text{H}}_2}{\text{O}} \ \\ $ 

(xii) A brown precipitate or colouring is created when \[{\text{N}}{{\text{H}}_3}\] gas is introduced into a colourless solution of Nessler's reagent. This is an \[{\text{N}}{{\text{H}}_3}\] gas test.

(xiii) \[{{\text{H}}_2}{\text{PtC}}{{\text{l}}_6} + 2{\text{N}}{{\text{H}}_4}{\text{Cl}} \to {{\text{(N}}{{\text{H}}_4})_2}\left[ {{\text{PtC}}{{\text{l}}_6}} \right] \downarrow {\text{yellow + 2HCl}}\]

Uses: 

(i) It's a refrigerant fluid.

(ii) Ammonium fertilisers such as ammonium sulphate, ammonium phosphate, ammonium nitrate, and urea are produced.

(iii) To get rid of grease because \[{\text{N}}{{\text{H}}_4}{\text{OH}}\] dissolves it.

(iv) For the Ostwald procedure of producing \[{\text{HN}}{{\text{O}}_3}\].

(v) As a reagent in a laboratory.

(vi) In the manufacture of artificial rayon, silk, nylon, and other similar materials.

(2) Oxides of Nitrogen:

Nitrogen creates a variety of oxides, including \[{{\text{N}}_2}{\text{O}}\], \[{\text{NO}}\], \[{{\text{N}}_2}{{\text{O}}_3}\], \[{\text{N}}{{\text{O}}_2}\] or \[{{\text{N}}_2}{{\text{O}}_4}\] and \[{{\text{N}}_2}{{\text{O}}_5}\], as well as the highly unstable \[{\text{N}}{{\text{O}}_3}\] and \[{{\text{N}}_2}{{\text{O}}_6}\]. All these nitrogen oxides have numerous p-p bonds between nitrogen and oxygen.

Preparation

(i) \[{{\text{N}}_2}{\text{O}}\] is often generated by carefully heating \[{\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3}\].

${\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} \to {{\text{N}}_2}{\text{O + 2}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{2}}{{\text{N}}_2}{\text{O + 2}}{{\text{H}}_2}{\text{S}}{{\text{O}}_3} \to {{\text{N}}_2}{\text{O + 2}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \ \\ $ 

(ii) The best way to make \[{\text{NO}}\] is to reduce 8 \[{\text{M HN}}{{\text{O}}_3}\] using reducing agents like \[{\text{Cu}}\], or to reduce nitrous acid or nitrites with \[{\text{F}}{{\text{e}}^{2 + }}\] or –ions.

${\text{3Cu + 8HN}}{{\text{O}}_3} \to 3{\text{Cu}}{\left( {{\text{N}}{{\text{O}}_3}} \right)_2}{\text{ + 2NO + 4}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{2NaNO + 2FeS}}{{\text{O}}_4} \to {\text{4NaHS}}{{\text{O}}_4}{\text{Fe(S}}{{\text{O}}_4}{{\text{)}}_3} + 2{\text{NO + 2}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{2NaN}}{{\text{O}}_2} + 2{\text{NaI + 4}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{4NaHS}}{{\text{O}}_4} + 2{\text{NO + }}{{\text{I}}_2} + 2{{\text{H}}_2}{\text{O}} \ \\ $ 

(iii) After cooling an equimolar mixture of \[{\text{NO}}\] and \[{\text{N}}{{\text{O}}_2}\] to 250 K, \[{{\text{N}}_2}{{\text{O}}_3}\] is formed as an intense blue liquid or a pale blue solid.

\[{\text{NO + N}}{{\text{O}}_2} \to {{\text{N}}_2}{{\text{O}}_3}\]

It’s colour diminishes as it warms due to dissociation into these two oxides.

(iv) \[{\text{N}}{{\text{O}}_2}\] can be made by heating heavy metal nitrates or reducing concentrated \[{\text{HN}}{{\text{O}}_3}\] with \[{\text{Cu}}\].

${\text{Cu + 4HN}}{{\text{O}}_3} \to {\text{Cu(N}}{{\text{O}}_3}{)_2} + 2{\text{N}}{{\text{O}}_2} + 2{{\text{H}}_2}{\text{O}} \ \\ $  ${\text{2Pb(N}}{{\text{O}}_3}{)_2}\xrightarrow{{673{\text{K}}}}2{\text{PbO + }}2{\text{N}}{{\text{O}}_2} + {{\text{O}}_2} \ \\ $ 

(v) The anhydride of \[{\text{HN}}{{\text{O}}_3}\] is \[{{\text{N}}_2}{{\text{O}}_5}\]. The best way to make it is to dehydrate \[{\text{HN}}{{\text{O}}_3}\] with \[{{\text{P}}_4}{{\text{O}}_{10}}\] at low temperatures.

\[{\text{4HN}}{{\text{O}}_3} + {{\text{P}}_4}{{\text{O}}_{10}}\xrightarrow{{250{\text{K}}}}{\text{2}}{{\text{N}}_5}{{\text{O}}_5} + 4{\text{HP}}{{\text{O}}_3}\]

Properties

All nitrogen oxides are oxidising agents, with \[{{\text{N}}_2}{\text{O}}\] even assisting in the burning of \[{\text{S}}\]and\[{\text{P}}\]. The combustion of \[{\text{Mg}}\] and \[{\text{P}}\] is supported by \[{\text{NO}}\], which is thermally more stable, but not of \[{\text{S}}\]. The sulphur flame is insufficiently hot to disintegrate it. \[{{\text{N}}_2}{\text{O}}\] has a linear structure and is isoelectronic with \[{\text{C}}{{\text{O}}_2}\]. Unlike \[{\text{C}}{{\text{O}}_2}\], however, \[{{\text{N}}_2}{\text{O}}\] has a modest dipole moment. The total number of electrons in \[{\text{NO}}\] is 15. This is an odd electron molecule since it is impossible for all of them to be paired. It is paramagnetic in its gaseous state. The liquid and solid states, on the other hand, are diamagnetic due to the formation of loose dimmers that balance out the magnetic effects of unpaired electrons. The production of a compound of iron, \[{\left[ {{\text{Fe(}}{{\text{H}}_2}{\text{O}}{{\text{)}}_5}{\text{NO}}} \right]^{2 + }}\], results in the brown ring seen in the nitrate test. With 23 electrons, \[{\text{N}}{{\text{O}}_2}\] is another another unusual electron molecule. It is paramagnetic in its gaseous state. The gas condenses to a brown liquid and then to a colourless solid when it cools, both of which are diamagnetic due to dimerization. Because liquid \[{{\text{N}}_2}{{\text{O}}_4}\] self-ionizes to create \[{\text{N}}{{\text{O}}^ + }\] and \[{\text{N}}{{\text{O}}_3}\] -ions, it has been researched extensively as a nonaqueous solvent.

\[{{\text{N}}_2}{\text{O}}\]

(a) Reduction: \[{\text{Cu(hot) + }}{{\text{N}}_2}{\text{O}} \to {\text{CuO + }}{{\text{N}}_2}\]

(b) Oxidation: \[{\text{2KMn}}{{\text{O}}_4} + 3{{\text{H}}_2}{\text{S}}{{\text{O}}_4} + 5{{\text{N}}_2}{\text{O}} \to {{\text{K}}_2}{\text{S}}{{\text{O}}_4} + 2{\text{MnS}}{{\text{O}}_4} + 3{{\text{H}}_2}{\text{O + 10NO}}\]

(c) Supporter of combustion: \[{\text{Mg + }}{{\text{N}}_2}{\text{O}} \to {\text{MgO + }}{{\text{N}}_2}\]

\[{\text{NO}}\]

(a) Supporter of combustion: \[{\text{S + 2NO}} \to {\text{S}}{{\text{O}}_2}{\text{ + }}{{\text{N}}_2}\]

(b) Oxidising properties (Reduction of NO): 

$5{{\text{H}}_2} + {\text{2NO}}\xrightarrow{{{\text{Pt - black}}}}2{\text{N}}{{\text{H}}_3} + 2{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{S}}{{\text{O}}_2}{\text{ + }}{{\text{H}}_2}{\text{O + 2NO}} \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ + }}{{\text{N}}_2}{\text{O}} \ \\ $

$  {{\text{H}}_2}{\text{S + 2NO}} \to {{\text{H}}_2}{\text{O}} + {{\text{N}}_2}{\text{O}} + {\text{S}} \ \\ $ 

(c) Reducing properties (oxidation of NO):

$ {\text{2NO + }}{{\text{X}}_2} \to 2{\text{NOX}} \ \\ $

$  6{\text{KMn}}{{\text{O}}_4} + 9{{\text{H}}_2}{\text{S}}{{\text{O}}_4} + 10{\text{NO}} \to {\text{3}}{{\text{K}}_2}{\text{S}}{{\text{O}}_4} + 6{\text{MnS}}{{\text{O}}_4} + 4{{\text{H}}_2}{\text{O + 10HN}}{{\text{O}}_3} \ \\ $ 

\[{{\text{N}}_2}{{\text{O}}_3}\]

(a) \[{{\text{N}}_2}{{\text{O}}_3} + {\text{KOH}} \to {\text{2KN}}{{\text{O}}_2} + 2{\text{KN}}{{\text{O}}_2} + {{\text{H}}_2}{\text{O}}\]

(b) It is an \[{\text{HN}}{{\text{O}}_2}\] anhydride.

\[{\text{2HN}}{{\text{O}}_2} \to {{\text{N}}_2}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}}\]

(c) produce nitrosyl salts with concentrated acids

\[{{\text{N}}_2}{{\text{O}}_3} + 2{\text{HCl}}{{\text{O}}_4} \to 2{\text{NO}}\left[ {{\text{Cl}}{{\text{O}}_4}} \right] + {{\text{H}}_2}{\text{O}}\]

\[{\text{N}}{{\text{O}}_2}\]

It has the properties of both \[{\text{HN}}{{\text{O}}_2}\] and \[{\text{HN}}{{\text{O}}_3}\]. According to the reactions below, it behaves like \[{\text{HN}}{{\text{O}}_2}\] as a reducing agent and like \[{\text{HN}}{{\text{O}}_3}\] as an oxidising agent.

$  {\text{2KMn}}{{\text{O}}_4} + 3{{\text{H}}_2}{\text{S}}{{\text{O}}_4} + \ 10{\text{N}}{{\text{O}}_2} + \ 2{{\text{H}}_2}{\text{O}} \to {{\text{K}}_2}{\text{S}}{{\text{O}}_4} + \  2{\text{MnS}}{{\text{O}}_4}{\text{ + 10HN}}{{\text{O}}_3} \ \\ $

$  {\text{S}}{{\text{O}}_4} + {{\text{H}}_2}{\text{O + N}}{{\text{O}}_2} \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4} + {\text{NO}} \ \\ $ 

\[{{\text{N}}_2}{{\text{O}}_4}\] is a mixture of \[{\text{HN}}{{\text{O}}_3}\] and \[{\text{HN}}{{\text{O}}_2}\] anhydride.

\[{{\text{N}}_2}{{\text{O}}_5}\]

(a) \[2{{\text{N}}_2}{{\text{O}}_5}\xrightarrow{\Delta }2{{\text{N}}_2}{{\text{O}}_4} + {{\text{O}}_2}\]

(b) \[{{\text{N}}_2}{{\text{O}}_5} + 2{\text{NaOH}} \to 2{{\text{N}}_2}{\text{N}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}}\]

(c) \[{{\text{N}}_2}{{\text{O}}_5} + {{\text{I}}_2} \to 10{\text{N}}{{\text{O}}_2} + {{\text{I}}_2}{{\text{O}}_5}\]

(3) Oxyacids of Nitrogen:

(a) Nitrous Acid (\[{\text{HN}}{{\text{O}}_2}\]):

Preparation:

(i) By acidifying a nitrite aqueous solution

\[{\text{Ba(N}}{{\text{O}}_3}{)_2} + {{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to 2{\text{HN}}{{\text{O}}_3} + {\text{BaS}}{{\text{O}}_4} \downarrow \]

(ii) Injecting water with an equimolar combination of \[{\text{NO}}\] and \[{\text{N}}{{\text{O}}_2}\].

Properties:

(i) It is a weak, unstable acid that can only be found in aqueous solution.

(ii) When attempting to concentrate, the acid decomposes into the following.

\[{\text{3HN}}{{\text{O}}_3} \to {\text{HN}}{{\text{O}}_3} + 2{\text{NO + }}{{\text{H}}_2}{\text{O}}\]

(iii) Nitrous acid and nitrites are effective oxidizers, converting iodides to iodine, ferrous salts to ferric, stannous to stannic, and sulphites to sulphates, for example.

\[{\text{2KI + 2HN}}{{\text{O}}_2} + 2{\text{HCl}} \to {\text{2}}{{\text{H}}_2}{\text{O + 2NO + 2KCl + }}{{\text{I}}_2}\]

(iv) Nitrous acid and nitrites act as reducing agents when exposed to strong oxidising agents such as \[{\text{KMn}}{{\text{O}}_4}\] and are reduced to \[{\text{N}}{{\text{O}}_3}\] -ions:

\[{\text{2KMn}}{{\text{O}}_4}{\text{ + 5KN}}{{\text{O}}_2} + 6{\text{HCl}} \to {\text{2MnC}}{{\text{l}}_2}{\text{ + 5KN}}{{\text{O}}_3}{\text{ + 2}}{{\text{H}}_2}{\text{O + 2KCl}}\]

(iii) The ion nitrite is an excellent coordinating agent. Lone pairs in both nitrogen and oxygen can form coordinate bonds with metal ions. The nitrite ion can coordinate through either \[{\text{N}}\] or \[{\text{O}}\]. This results in linkage isomerism (it is an ambidentate ligand). The nitrites, \[{\text{RONO}}\], and the nitro compounds \[{\text{RN}}{{\text{O}}_2}\], where \[{\text{R}}\] is any alkyl or aryl group, are analogous organic derivatives.

(B) Nitric Acid (\[{\text{HN}}{{\text{O}}_3}\])

PREPARATION

\[{\text{NaN}}{{\text{o}}_3} + {{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{NaHS}}{{\text{O}}_4} + {\text{HN}}{{\text{O}}_3}\]

(ii) The Ostwald technique is now virtually entirely used to produce \[{\text{HN}}{{\text{O}}_3}\]. \[{\text{N}}{{\text{H}}_3}\] is catalytically oxidised to \[{\text{NO}}\] over a Pt-Rh catalyst at 1200K in this procedure.

\[4{\text{N}}{{\text{H}}_3} + 5{{\text{O}}_2} \to 4{\text{NO + 6}}{{\text{H}}_2}{\text{O }}\Delta {\text{H  =   - 904 kJ}}\]

\[{\text{N}}{{\text{H}}_3}\] is transformed to \[{\text{NO}}\] at a rate of 96 to 98 percent. After then, the mixture is diluted with air. \[{\text{N}}{{\text{O}}_2}\] is formed when \[{\text{NO}}\] interacts with \[{{\text{O}}_2}\] and is absorbed by water, yielding \[{\text{HN}}{{\text{O}}_3}\] and \[{\text{NO}}\], which is subsequently recycled.

$  {\text{2NO + }}{{\text{O}}_2} \to 2{\text{N}}{{\text{O}}_2} \ \\ $

$  {\text{3N}}{{\text{O}}_2}{\text{ + }}{{\text{H}}_2}{\text{O }} \to {\text{2HN}}{{\text{O}}_3} + {\text{NO}} \ \\ $ 

When a steady boiling mixture is generated, nitric acid can be concentrated to 68 percent by distillation. Distilling the mixture with concentrated sulphuric acid yields a more concentrated acid.

Properties

(i) Nitric acid in its purest form is a colourless liquid (bp 359oC). It decomposes quickly in light, resulting in the creation of nitrogen dioxide, which gives it a yellow colour. It is a powerful acid that dissociates almost completely into ions in solution.

(ii) Thermal Stability

\[{\text{4HN}}{{\text{O}}_3}\xrightarrow{\Delta }{\text{2}}{{\text{H}}_2}{\text{O  + 4N}}{{\text{O}}_2} + {{\text{O}}_2}\]

(iii) Oxidising properties

${\text{2HN}}{{\text{O}}_3}{\text{(conc}}{\text{.)}} \to {\text{H}_2}{\text{O + 2N}}{\text{O}_2} + \left [ {\text{O}} \right ] \ \\ $

$  {\text{2HN}}{{\text{O}}_3}(dilute{\text{)}} \to {{\text{H}}_2}{\text{O  + 2NO}} + 3\left[ {\text{O}} \right] \ \\ $ 

(a) Oxidises \[{{\text{H}}_2}{\text{S}}\] to sulphur

$ {{\text{H}}_2}{\text{S  + 2HN}}{{\text{O}}_3}(conc.) \to 2{{\text{H}}_2}{\text{O + 2N}}{{\text{O}}_2} + {\text{S}} \downarrow  \ \\ $ ${\text{3}}{{\text{H}}_2}{\text{S  + 2HN}}{{\text{O}}_3}(dilute) \to 4{{\text{H}}_2}{\text{O + 2NO}} + 3{\text{S}} \ \\ $ 

(b) Oxidises \[{\text{S}}{{\text{O}}_2}\] to \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\]

$ {\text{S}}{{\text{O}}_2} + 2{\text{HN}}{{\text{O}}_3}(conc.) \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ + 2NO}} \uparrow  \ \\ $

$  {\text{S}}{{\text{O}}_2} + 2{{\text{H}}_2}{\text{O + }}2{\text{HN}}{{\text{O}}_3}(dilute) \to 3{{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ + 2NO}} \ \\ $ 

Ferrous salts are oxidised to ferric salts, while halogen acids are oxidised to their respective halogens.

(iv) Reaction with non-metals.

$ {\text{C + 4HN}}{{\text{O}}_3} \to {{\text{H}}_2}{\text{C}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O + 4N}}{{\text{O}}_2} \ \\ $

$  {\text{S + 6HN}}{{\text{O}}_3} \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4} + 2{{\text{H}}_2}{\text{O + 6N}}{{\text{O}}_2} \ \\ $

$  {{\text{I}}_2} + 10{\text{HN}}{{\text{O}}_3} \to 2{\text{HI}}{{\text{O}}_3} + 4{{\text{H}}_2}{\text{O + 10N}}{{\text{O}}_2} \ \\ $

$  {{\text{P}}_2} + 5{\text{HN}}{{\text{O}}_3} \to 2{{\text{H}}_3}{\text{P}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O + 5N}}{{\text{O}}_2} \ \\ $ 

Conc. \[{\text{HN}}{{\text{O}}_3}\] is used in these processes.

(v) Reaction with metals

(A) Metals with a higher electropositive charge than hydrogen.

(a) Action on zinc or \[{\text{Fe}}\]:

$ {\text{Zn + 4HN}}{{\text{O}}_3}({\text{conc}}{\text{.)}} \to {\text{Zn(N}}{{\text{O}}_3}{)_2} + 2{{\text{H}}_2}{\text{O + 2N}}{{\text{O}}_2} \ \\ $

$ {\text{4Zn + 10HN}}{{\text{O}}_3}(dil.{\text{)}} \to 4{\text{Zn(N}}{{\text{O}}_3}{)_2} + 5{{\text{H}}_2}{\text{O + N}}{{\text{O}}_2} \ \\ $

$  {\text{4Zn + 10HN}}{{\text{O}}_3}(v.dil.{\text{)}} \to 4{\text{Zn(N}}{{\text{O}}_3}{)_2} + {\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{4Mg + 10HN}}{{\text{O}}_3}(v.dil.{\text{)}} \to 4{\text{Mg(N}}{{\text{O}}_3}{)_2} + {\text{N}}{{\text{H}}_4}{\text{N}}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \ \\ $ 

(b) Action on tin:

$  {\text{Sn + 4HN}}{{\text{O}}_3}({\text{conc}}{\text{.)}} \to {{\text{H}}_2}{\text{Sn}}{{\text{O}}_3}{\text{ + 4N}}{{\text{O}}_2} + {{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{Sn + 4HN}}{{\text{O}}_3}(dil{\text{.)}} \to 4{\text{Sn(N}}{{\text{O}}_3}{{\text{)}}_2}{\text{ + 2NO}} + 4{{\text{H}}_2}{\text{O}} \ \\ $ 

(c) Action on lead

${\text{Pb + 4HN}}{{\text{O}}_3}({\text{conc}}{\text{.)}} \to {\text{Pb(N}}{{\text{O}}_3}{{\text{)}}_2}{\text{ + 2N}}{{\text{O}}_2} + 2{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{3Pb + 8HN}}{{\text{O}}_3}(dil{\text{.)}} \to 4{\text{Pb(N}}{{\text{O}}_3}{{\text{)}}_2}{\text{ + 2NO}} + 4{{\text{H}}_2}{\text{O}} \ \\ $ 

(B) Metals with a lower electrostatic charge than hydrogen.

(i) Action on copper

$ {\text{Cu + 4HN}}{{\text{O}}_3}({\text{conc}}{\text{.)}} \to {\text{Cu(N}}{{\text{O}}_3}{{\text{)}}_2}{\text{ + 2N}}{{\text{O}}_2} + 2{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{3Cu + 8HN}}{{\text{O}}_3}(v.dil{\text{.)}} \to 4{\text{Cu(N}}{{\text{O}}_3}{{\text{)}}_2}{\text{ + NO}} + 4{{\text{H}}_2}{\text{O}} \ \\ $ 

(C) Metalloids: Sb and As

\[{\text{Sb + 5HN}}{{\text{O}}_3}({\text{conc}}{\text{.)}} \to {{\text{H}}_3}{\text{Sb}}{{\text{O}}_4}({\text{antimonic acid) + 5N}}{{\text{O}}_2} + \ {{\text{H}}_2}{\text{O}}\]

The only metals that create hydrogen gas with cold (1–2%) \[{\text{HN}}{{\text{O}}_3}\] are magnesium and manganese.

${\text{Mg + 2HN}}{{\text{O}}_3} \to {\text{Mg(N}}{{\text{O}}_3}{)_2} + {{\text{H}}_2} \uparrow  \ \\ $

$  {\text{Mn + 2HN}}{{\text{O}}_3} \to {\text{Mn(N}}{{\text{O}}_3}{)_2} + {{\text{H}}_2} \uparrow  \ \\ $ 

As a result of the creation of a persistent layer of insoluble oxide on the metal surface, concentrated nitric acid (80 percent) acts as an oxidising agent, rendering metals such as \[{\text{Al,Fe,Cr}}\], and others inactive.

Nitric acid has little effect on noble metals such as \[{\text{Au,Pt,Rh}}\], and Ir. Aqua regia, a 1:3 mixture of concentrated \[{\text{HN}}{{\text{O}}_3}\] and concentrated \[{\text{HCl}}\], dissolves \[{\text{Au}}\] and \[{\text{Pt}}\] because it includes free(atomic) chlorine:

${\text{HN}}{{\text{O}}_3} + 3{\text{HCl}} \to {\text{2}}{{\text{H}}_2}{\text{O + 2Cl + NOCl}} \ \\ $

$  {\text{Au + 3Cl + HCl}} \to {\text{HAuC}}{{\text{l}}_4} \ \\ $

$  {\text{Pt + 4Cl + 2HCl}} \to {{\text{H}}_2}{\text{PtC}}{{\text{l}}_6} \ \\ $ 

Brown ring test

$ {\text{2HN}}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{S}}{{\text{O}}_4} + 6{\text{FeS}}{{\text{O}}_4} \to 3{\text{F}}{{\text{e}}_2}{({\text{S}}{{\text{O}}_4})_3} + 2{\text{NO + 4}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{FeS}}{{\text{O}}_4} + {\text{NO + 5}}{{\text{H}}_2}{\text{O}} \to \left[ {{\text{Fe(}}{{\text{H}}_2}{\text{O}}{{\text{)}}_5}{\text{NO}}} \right]{\text{S}}{{\text{O}}_4} \ \\ $ 

2. Phosphorus:

It's a non-metal with a high reactivity. In the air, it catches fire. It can be found in nature as stable phosphates. (Calcium phosphate is also found in animal bones) (58 percent). The following minerals are essential

(i) Phosphorite \[{\text{C}}{{\text{a}}_3}{({\text{P}}{{\text{O}}_4})_2}\]

(ii) Chloraptite \[{\text{C}}{{\text{a}}_3}{({\text{P}}{{\text{O}}_4})_2}{\text{CaC}}{{\text{l}}_2}\]

(iii) Fluoraptite \[{\text{C}}{{\text{a}}_3}{({\text{P}}{{\text{O}}_4})_2}{\text{Ca}}{{\text{F}}_2}\]

(iv) Vivianite 

(v) Redonda phosphate \[{\text{AlP}}{{\text{O}}_4}\]

Allotropic Forms of Phosphorus:

(i) White or Yellow Phosphorus (P4 ):

Preparation:

\[2{\text{C}}{{\text{a}}_3}{({\text{P}}{{\text{O}}_4})_2}({\text{From bone - ash) + 10C + 6CaSi}}{{\text{O}}_3} + 10{\text{C + }}{{\text{P}}_4}({\text{s)}}\]

Properties:

It's a soft waxy substance that's white to transparent. At 20°C, it has a density of 1.8 g/cc. Its mp and bp are 44 and 287 degrees Celsius, respectively. It is water insoluble but soluble in CS2. Slow oxidation produces a yellowish-green light that glows in the dark. The term for this phenomenon is phosphorescence.

\[{{\text{P}}_4} + 5{{\text{O}}_2} \to {{\text{P}}_4}{{\text{O}}_{10}}\]

Phosphorus in the form of white phosphorus is toxic. After a while, it turns yellow and is known as yellow phosphorus. It oxidises in the presence of air, gradually raising its temperature until, due to its low ignition temperature (30oC), it spontaneously catches fire after a few moments. It is stored under water because of this.

Acts as a reducing agent since it is easily oxidised.

$ {{\text{P}}_4} + 20{\text{HN}}{{\text{O}}_3} \to 4{{\text{H}}_3}{\text{P}}{{\text{O}}_4} + 20{\text{N}}{{\text{O}}_2} + 4{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{3CaO + 8P + 9}}{{\text{H}}_2}{\text{O}} \to {\text{3Ca(}}{{\text{H}}_2}{\text{P}}{{\text{O}}_2}{{\text{)}}_2}{\text{ + 2P}}{{\text{H}}_3} \ \\ $

$  {\text{C}}{{\text{u}}_3}{{\text{P}}_2} + 5{\text{CuS}}{{\text{O}}_4} + 8{{\text{H}}_2}{\text{O}} \to 8{\text{Cu + 5}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ + 2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_4} \ \\ $ 

A colloidal gold solution can be made by reducing a gold chloride solution with phosphorus dissolved in ether.

(ii) Red phosphorus:

Red phosphorus is produced when white phosphorus is heated in the presence of \[{\text{C}}{{\text{O}}_2}\] or coal gas at 573 K. This red phosphorus may still contain some white phosphorus, which can be eliminated by boiling the combination with \[{\text{NaOH}}\], which converts the white phosphorus to \[{\text{P}}{{\text{H}}_3}\] gas while leaving the red phosphorus inert.

\[{{\text{P}}_4} + 3{\text{NaOH + }}3{{\text{H}}_2}{\text{O}} \to {\text{P}}{{\text{H}}_3}(g) + 3Na{H_2}P{O_2}\]

It can also be made by heating white phosphorus with a few crystals of iodine catalyst at 250°C in the absence of air under high pressure.

Properties

It has a density of 2.2 g/cc and is a red crystalline solid. It has a lower reactivity than white phosphorus and will not dissolve in \[{\text{C}}{{\text{S}}_2}\] liquid. Because its ignition temperature is 260oC, it does not catch fire at room temperature.

This is a polymeric material that forms linear chains.

(iii) Black phosphorus: It comes in two varieties: \[\alpha \]- black phosphorus and \[\beta \]- black phosphorous.

(a)  α-black Phosphorous

\[{\text{P(red)}}\xrightarrow[{{\text{tube 803K}}}]{{{\text{insulated}}}}{\text{P(}}\alpha {\text{ - black)}}\]

The structure of α black phosphorous is amorphous and does not conduct electricity.

(b) β-black Phosphorous: 

\[{\text{P(white)}}\xrightarrow[{{\text{High pressure}}}]{{{\text{473K}}}}{\text{P(}}\beta {\text{ - black)}}\]

β-black phosphorous is an electrical conductor that is like graphite in terms of flakiness and shine. In \[{\text{C}}{{\text{S}}_2}\], it is insoluble. Like graphite, it has a layered structure.

(iv) Brown phosphorus: \[{{\text{P}}_4}\] molecules begin to break into \[{{\text{P}}_2}\] molecules at temperatures over 1600oC. Brown phosphorus is formed when this vapour is rapidly cooled, and it contains \[{{\text{P}}_2}\] molecules.

Chemical Properties of Phosphorus

The reactivity of various allotropic forms of phosphorus to other chemicals declines in the following order: brown > white > red > black, with the last one being nearly inert.

Aside from the differences in reactivity, all the forms are chemically equivalent.

(i) Action of air

Phosphorus trioxide and pentoxide are formed when white phosphorus is burned in the air.

$ {{\text{P}}_4} + 5{{\text{O}}_2} \to 2{{\text{P}}_2}{{\text{O}}_5} \ \\ $

$  {{\text{P}}_4} + 3{{\text{O}}_2} \to 2{{\text{P}}_2}{{\text{O}}_3} \ \\ $ 

Phosphorus in red and other forms also burns in air or oxygen, but only when heated.

(ii) Action of non-metals

Phosphorus creates compounds when heated with non-metals. \[{\text{P}}{{\text{X}}_3},{\text{P}}{{\text{X}}_5},{{\text{P}}_2}{{\text{S}}_3}{\text{ and }}{{\text{P}}_2}{{\text{S}}_5}\]

$2{\text{P + 3}}{{\text{X}}_2} \to 2{\text{P}}{{\text{X}}_3} \ \\ $

$  2{\text{P + 5}}{{\text{X}}_2} \to 2{\text{P}}{{\text{X}}_5} \ \\ $ 

(iii) Action with metals

When alkali metals are burned in vacuum with white phosphorus, alkali metal phosphide is formed, which reacts with water to form phosphine gas.

$ 3{\text{M + P}}\xrightarrow{\Delta }{{\text{M}}_3}{\text{P}} \ \\ $

$  {{\text{M}}_3}{\text{P}} + 3{{\text{H}}_2}{\text{O}}\xrightarrow{\Delta }3{\text{MOH + P}}{{\text{H}}_3} \uparrow  \ \\ $ 

(iv) Action of \[{\text{NaOH}}\]

Phosphine gas is produced when white phosphorus is cooked in a \[{\text{NaOH}}\] solution.

\[{{\text{P}}_4} + 3{\text{NaOH + 3}}{{\text{H}}_2}{\text{O}} \to {\text{3Na}}{{\text{H}}_2}{\text{P}}{{\text{O}}_2} + {\text{P}}{{\text{H}}_3} \uparrow \]

(v) Action of conc. \[{\text{HN}}{{\text{O}}_3}\]

Phosphorus is oxidised to \[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\] when heated with conc. \[{\text{HN}}{{\text{O}}_3}\].

\[{{\text{P}}_4} + 5{\text{HN}}{{\text{O}}_3} \to {{\text{H}}_3}{\text{P}}{{\text{O}}_4} + 5{\text{N}}{{\text{O}}_2} \uparrow  + {{\text{H}}_2}{\text{O}}\]

(vi) Action of conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\]

When heated with conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] phosphorus is oxidized to \[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\].

\[{\text{2P}} + 5{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\xrightarrow{\Delta }{\text{2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_4} + 5{\text{S}}{{\text{O}}_2} \uparrow  + 2{{\text{H}}_2}{\text{O}}\]

White phosphorus is converted to red phosphorus, and red phosphorus is converted to white phosphorus. Heat to 250oC and high pressure in the absence of air in the presence of 2 catalyst.

White phosphorus \[ \rightleftharpoons \]Red phosphorus

In the presence of an inert gas, heat over 250°C and condense in water.

Compounds of Phosphorus

(1) Phosphine

 Preparation

(i) In the presence of coal gas, white phosphorus is heated with a \[{\text{NaOH}}\] solution. The downward displacement of water collects phosphorus gas.

\[{{\text{P}}_2}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \to 2{{\text{H}}_2}{\text{P}}{{\text{O}}_3}\]

To keep phosphine from oxidising, coal gas is employed. The phsophine gas is polluted with \[{{\text{P}}_2}{{\text{H}}_4}\], a flammable gas. It is separated from \[{\text{P}}{{\text{H}}_3}\] by passing the gaseous combination through a freezing mixture, where \[{{\text{P}}_2}{{\text{H}}_4}\] condenses to a liquid and \[{\text{P}}{{\text{H}}_3}\] is recovered by forcing air downward. In the absence of oxygen, pure \[{\text{P}}{{\text{H}}_3}\] does not burn.

(ii) As a result of alkali action on phosphonium salts:

\[{\text{P}}{{\text{H}}_4}{\text{I + NaOH}}\xrightarrow{\Delta }{\text{NaI + P}}{{\text{H}}_3} \uparrow  + {{\text{H}}_2}{\text{O}}\]

(iii) As a result of the action of dil. On metal phosphides, use \[{\text{HCl}}\] or diluted \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\].

$ {\text{N}}{{\text{a}}_3}{\text{P + 3HCl}} \to {\text{3NaCl + P}}{{\text{H}}_3} \uparrow  \ \\ $

$  {\text{AlP + 3HCl}} \to {\text{AlC}}{{\text{l}}_3}{\text{ + P}}{{\text{H}}_3} \uparrow  \ \\ $

$  {\text{2N}}{{\text{a}}_3}{\text{P + 3}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{3N}}{{\text{a}}_2}{\text{S}}{{\text{O}}_4}{\text{ + 2P}}{{\text{H}}_3} \uparrow  \ \\ $ 

(iv) \[{{\text{H}}_3}{\text{P}}{{\text{O}}_2}{\text{ + 4H}}\xrightarrow{{{\text{Zn/HCl}}}}{\text{P}}{{\text{H}}_3}{\text{ + 2}}{{\text{H}}_2}{\text{O}}\]

Properties

(i) It is a colourless gas with a rotten fish odour that is litmus paper neutral; it is heavier than air and only slightly soluble in water; it is a toxic gas that functions as a Lewis base; and it is heavier than air and only slightly soluble in water.

(ii) Action of chlorine

\[{\text{PC}}{{\text{l}}_5}\] is formed when it interacts with \[{\text{C}}{{\text{l}}_2}{\text{ }}\]

\[{\text{P}}{{\text{H}}_3}{\text{ + 4C}}{{\text{l}}_2} \to {\text{PC}}{{\text{l}}_5}{\text{ + 3HCl}}\]

(iv) Action on \[{\text{CuS}}{{\text{O}}_4}\] solution

When \[{\text{P}}{{\text{H}}_3}\] is bubbled in an acidic copper sulphate solution, a black copper phosphide precipitate form.

\[{\text{3CuS}}{{\text{O}}_4}{\text{ + 2P}}{{\text{H}}_3} \to {\text{C}}{{\text{u}}_3}{{\text{P}}_2} \downarrow \left\{ {{\text{Black}}} \right\} + 3{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\]

(v) Reaction with \[{\text{AgN}}{{\text{O}}_3}\] solution

When \[{\text{P}}{{\text{H}}_3}\] gas is bubbled into an \[{\text{AgN}}{{\text{O}}_3}\] solution, a yellow silver phosphide precipitate, \[{\text{A}}{{\text{g}}_3}{\text{P}}\], forms first, then decomposes to black \[{\text{Ag}}\].

$ {\text{3AgN}}{{\text{O}}_3} + {\text{P}}{{\text{H}}_3} \to {\text{A}}{{\text{g}}_3}{\text{P}} \downarrow {\text{(yellow)  + 3HN}}{{\text{O}}_3} \ \\ $

$  {\text{A}}{{\text{g}}_3}{\text{P}} + {\text{3AgN}}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \to {\text{6Ag}} \downarrow {\text{(black) + 3HN}}{{\text{O}}_3} + {{\text{H}}_3}{\text{P}}{{\text{O}}_3} \ \\ $ 

(vi) Reaction with mercuric chloride solution

A brownish black precipitate of mercuric phosphide is generated when mercuric chloride solution is treated with \[{\text{P}}{{\text{H}}_3}\] gas.

\[{\text{3HgC}}{{\text{l}}_2}{\text{ + 2P}}{{\text{H}}_3} \to {\text{H}}{{\text{g}}_3}{{\text{P}}_2} \downarrow {\text{(brownish - black)  + 6HCl}}\]

Quick lime or \[{\text{NaOH}}\] sticks can be used to dry \[{\text{P}}{{\text{H}}_3}\] samples. Because of its reactivity with conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\], it cannot be dried with it.

Uses

It's utilised in battlefields to make smoke signals and create smoke screens.

2. Oxides of Phosphorus  

(a) Phosphorus Trioxide (\[{{\text{P}}_2}{{\text{O}}_3}\]) 

Preparation

When phosphorus is burned with a limited supply of oxygen, gaseous \[{{\text{P}}_4}{{\text{O}}_{10}}\] and \[{{\text{P}}_4}{{\text{O}}_6}\] are generated. \[{{\text{P}}_4}{{\text{O}}_6}\] remains gaseous when the temperature is lowered by a condenser, but \[{{\text{P}}_4}{{\text{O}}_{10}}\] condenses as a solid and is prevented by glasswool. When the residual gaseous mixture is passed through the freezing mixture, it transforms into colourless \[{{\text{P}}_4}{{\text{O}}_6}\] crystals.

Properties

(i) It is a colourless crystalline solid with a melting point of 23.8°C and a boiling point of 178°C.

(ii) Phosphorus acid is formed when it dissolves in cold water. As a result, it's phosphorus acid anhydride.

\[{{\text{P}}_2}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \to 2{{\text{H}}_2}{\text{P}}{{\text{O}}_3}\]

(iii) In hot water, it dissolves, releasing \[{\text{P}}{{\text{H}}_3}\].

\[{\text{2}}{{\text{P}}_2}{{\text{O}}_3} + 6{{\text{H}}_2}{\text{O}} \to 3{{\text{H}}_3}{\text{P}}{{\text{O}}_4} + {\text{P}}{{\text{H}}_3}\]

(iv) In the presence of air, it progressively oxidises to create \[{{\text{P}}_2}{{\text{O}}_5}\].

\[{{\text{P}}_2}{{\text{O}}_3} + {{\text{O}}_2} \to {{\text{P}}_2}{{\text{O}}_5}\]

(v) It decomposes into phosphorus oxytrichloride (\[{\text{POC}}{{\text{l}}_3}\]) and phosphoryl chloride (\[{\text{P}}{{\text{O}}_2}{\text{Cl}}\]) when exposed to \[{\text{C}}{{\text{l}}_2}\].

\[{{\text{P}}_2}{{\text{O}}_3} + 2{\text{C}}{{\text{l}}_2} \to {\text{POC}}{{\text{l}}_3} + {\text{P}}{{\text{O}}_2}{\text{Cl}}\]

(B) Phosphorus Pentoxide (\[{{\text{P}}_2}{{\text{O}}_5}\]) 

Preparation 

Phosphorus is obtained by burning it in the air.

\[{{\text{P}}_4}{\text{ + 5}}{{\text{O}}_2} \to {{\text{P}}_4}{{\text{O}}_{10}}\]

Properties

(i) It's an acidic white powder that's the anhydride of orthophosphoric acid. \[{{\text{P}}_2}{{\text{O}}_5}\] is its empirical formula, and \[{{\text{P}}_4}{{\text{O}}_{10}}\] is its molecular formula.

(ii) When heated to 250°C, it becomes sublime.

(iii) Action of Water

With a hissing sound, it dissolves in water, creating metaphosphoric acid and then orthophosphoric acid.

$ {{\text{P}}_4}{{\text{O}}_{10}}{\text{ + 2}}{{\text{H}}_2}{\text{O}} \to {\text{4HP}}{{\text{O}}_3} \ \\ $

$  {\text{HP}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}} \to {{\text{H}}_3}{\text{P}}{{\text{O}}_4} \ \\ $ 

It converts concentrated \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] and concentrated \[{\text{HN}}{{\text{O}}_3}\] to \[{\text{S}}{{\text{O}}_3}\] and \[{{\text{N}}_2}{{\text{O}}_5}\], respectively.

$  {\text{HN}}{{\text{O}}_3} + {{\text{P}}_2}{{\text{O}}_5}\xrightarrow{{{\text{distillation}}}}{\text{2HP}}{{\text{O}}_3} + {{\text{N}}_2}{{\text{O}}_5} \ \\ $

$  {{\text{H}}_2}{\text{S}}{{\text{O}}_4} + {{\text{P}}_2}{{\text{O}}_5}\xrightarrow{{{\text{distillation}}}}{\text{2HP}}{{\text{O}}_3} + {\text{S}}{{\text{O}}_3} \ \\ $ 

Uses

(i) Dehumidification of acidic gases

(ii) It can be used as a dehydrating agent.

(iii) For the manufacture of \[{\text{S}}{{\text{O}}_3}\] and \[{{\text{N}}_2}{{\text{O}}_5}\]

(iv) In order to make phosphoric acid

3. Oxy-Acids of Phosphorus 

(A) Phosphorus Acid (\[{{\text{H}}_3}{\text{P}}{{\text{O}}_3}\]) 

Preparation

(i) \[{{\text{P}}_2}{{\text{O}}_3}\] is dissolved in water.

\[{{\text{P}}_2}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \to {\text{2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3}\]

(ii) Using water to hydrolyze \[{\text{PC}}{{\text{l}}_3}\].

$ {{\text{P}}_2}{{\text{O}}_3} + 3{{\text{H}}_2}{\text{O}} \to {\text{2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3} \ \\ $

$  {\text{PC}}{{\text{l}}_3} + 3{{\text{H}}_2}{\text{O      }}{{\text{H}}_3}{\text{P}}{{\text{O}}_3} + 3{\text{HCl}} \ \\ $ 

The \[{{\text{H}}_3}{\text{P}}{{\text{O}}_3}\] and \[{\text{HCl}}\] solution is heated to 180°C, and the \[{\text{HCl}}\] gas is forced out. On crystallisation, the resultant solution yields white \[{{\text{H}}_3}{\text{P}}{{\text{O}}_3}\] crystals.

(iii) Using hypophosphorus acid that has been heated

\[{\text{3}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3}({\text{concentrated solution)}}\xrightarrow{{{{40}^0}{\text{ or more }}}}{\text{P}}{{\text{H}}_3}{\text{ + 2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3}\]

Properties

(i) It's a white crystalline solid that's water soluble and has a melting point of 74°C.

(ii) It is a reducing agent and a weak acid.

(iii) It generates unstable neutral salts termed phosphites when neutralised with bases or alkalies.

\[{\text{3}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3} + 3{\text{NaOH}} \to {\text{3}}{{\text{H}}_3}{\text{P}}{{\text{O}}_3} + {\text{P}}{{\text{H}}_3}({\text{Disproportionation)}}\]

(v) \[{{\text{H}}_3}{\text{P}}{{\text{O}}_3} + 3{\text{PC}}{{\text{l}}_5} \to {\text{PC}}{{\text{l}}_3} + 3{\text{POC}}{{\text{l}}_3} + 3{\text{HCl}}\]

(vi) It has a high reducing power.

$2{\text{AgN}}{{\text{O}}_3} + {{\text{H}}_3}{\text{P}}{{\text{O}}_3}{\text{ + }}{{\text{H}}_2}{\text{O}} \to {\text{2Ag + 2HN}}{{\text{O}}_3}{\text{ + }}{{\text{H}}_3}{\text{P}}{{\text{O}}_4} \ \\ $

$  2{\text{HgC}}{{\text{l}}_2}{\text{ + }}{{\text{H}}_3}{\text{P}}{{\text{O}}_3}{\text{ + }}{{\text{H}}_2}{\text{O}} \to {\text{H}}{{\text{g}}_2}{\text{C}}{{\text{l}}_2}{\text{ + 2HCl + }}{{\text{H}}_3}{\text{P}}{{\text{O}}_4} \ \\ $ 

(B) Orthophosphoric Acid (\[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\]) 

Preparation

(i) Using conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] to heat calcium phosphate

\[{\text{Ca(P}}{{\text{O}}_4}{{\text{)}}_2}{\text{ + 3}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{2CaS}}{{\text{O}}_4}{\text{ + 2}}{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\]

\[{\text{CaS}}{{\text{O}}_4}\] is insoluble in water. \[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\] solution is isolated from \[{\text{CaS}}{{\text{O}}_4}\] solution. It is then concentrated by evaporating it at 180°C and dehydrated by placing conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] in a vacuum desiccator and cooling the combination. As a result, white \[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\] crystals form.

(ii) hydrolysis of \[{\text{Pl}}{{\text{C}}_5}\]

\[{\text{Pl}}{{\text{C}}_5} + 4{{\text{H}}_2}{\text{O}} \to {{\text{H}}_3}{\text{P}}{{\text{O}}_4}{\text{ + 5HCl}}\]

(iii) Using conc. \[{\text{HN}}{{\text{O}}_3}\] to heat white phosphorus

\[{\text{P}} + 5{\text{HN}}{{\text{O}}_3} \to {{\text{H}}_3}{\text{P}}{{\text{O}}_4}{\text{ + 5N}}{{\text{O}}_2} + {{\text{H}}_2}{\text{O}}\]

Properties

(i) Pure orthophosphoric acid is a white crystalline solid with a melting point of 42°C that is readily soluble in water. It's a very weak acid. Two acid salts and one normal salt are formed. \[{\text{Na}}{{\text{H}}_2}{\text{P}}{{\text{O}}_4}\], \[{\text{N}}{{\text{a}}_2}{\text{HP}}{{\text{O}}_4}\], and \[{\text{N}}{{\text{a}}_3}{\text{P}}{{\text{O}}_4}\] are sodium dihydrogen phosphate, sodium hydrogen phosphate, and sodium orthophosphate, respectively.

(ii) Action of heat

${{\text{H}}_3}{\text{P}}{{\text{O}}_4}\xrightarrow{{{{220}^0}{\text{C}}}}{{\text{H}}_4}{{\text{P}}_2}{{\text{O}}_7}({\text{pyrophosphoric acid)}} \ \\ $

${{\text{H}}_4}{{\text{P}}_2}{{\text{O}}_7}\xrightarrow{{{{316}^0}{\text{C}}}}{\text{HP}}{{\text{O}}_3}({\text{metaphosphoric acid)}} \ \\ $ 

(iii) Neutralization with alkalies or bases

$ {{\text{H}}_3}{\text{P}}{{\text{O}}_4}\xrightarrow[{ - {{\text{H}}_2}{\text{O}}}]{{{\text{NaOH}}}}{\text{Na}}{{\text{H}}_2}{\text{P}}{{\text{O}}_4}(pri.{\text{phosphate)}}\xrightarrow[{ - {{\text{H}}_2}{\text{O}}}]{{{\text{NaOH}}}}{\text{NaHP}}{{\text{O}}_4}(sec. \ {\text{phosphate)}}\xrightarrow[{ - {{\text{H}}_2}{\text{O}}}]{{{\text{NaOH}}}}{\text{N}}{{\text{a}}_3}{\text{P}}{{\text{O}}_4}(tert.{\text{phosphate)}} \ \\ $

$  {\text{Na}}{{\text{H}}_2}{\text{P}}{{\text{O}}_4}\xrightarrow{\Delta }{\text{NaP}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}} \ \\ $ 

Group Sixteen Elements: The Oxygen Family

Group 16 of the periodic table includes oxygen, sulphur, selenium, tellurium, and polonium. This is also known as a chalcogen group. Sulphur and its congeners are linked to copper, as seen by the term, which is derived from the Greek word for brass. Most copper minerals contain either oxygen or sulphur, as well as other group members.

Occurrence

The element oxygen is the most prevalent on the planet. By mass, oxygen makes up 46.6 percent of the earth's crust. By volume, dry air contains 20.946 percent oxygen. The amount of sulphur in the earth's crust, on the other hand, is merely 0.03-0.1%. Sulphates such as gypsum \[{\text{CaS}}{{\text{O}}_4}{\text{.2}}{{\text{H}}_2}{\text{O}}\], epsom salt \[{\text{MgS}}{{\text{O}}_4}{\text{.7}}{{\text{H}}_2}{\text{O}}\], baryta \[{\text{BaS}}{{\text{O}}_4}\] and sulphides such as galena \[{\text{PbS}}\], zinc blende \[{\text{ZnS}}\], and copper pyrites \[{\text{CuFe}}{{\text{S}}_2}\] are the most common forms of combined sulphur. In volcanoes, traces of sulphur appear as hydrogen sulphide.

In sulphide ores, selenium and tellurium are found as metal selenides and tellurides. Polonium is a decay product of thorium and uranium minerals found in nature.

Electronic Configuration

Group 16 elements have six electrons in the outermost shell and a common electrical configuration of \[{\text{n}}{{\text{s}}^2}{\text{ n}}{{\text{p}}^4}\].

Atomic and Ionic Radii

Atomic and ionic radii rise from top to bottom in the group as the number of shells increases. The size of oxygen atoms, on the other hand, is extremely small.

Ionisation Enthalpy

The enthalpy of ionisation decreases as you progress through the group. It's because of the growth in size. In the corresponding periods, however, the elements of this group have lower ionisation enthalpy values than those of group 15. This is owing to the additional stable half-filled p orbitals elelctronic configurations found in group 15 elements. 

Electron Gain Enthalpy

The negative electron gain enthalpy of oxygen is lower than that of sulphur due to its compact structure. From sulphur onwards, the value grows less negative until it reaches polonium.

Electronegativity

Oxygen, along with fluorine, has the greatest electronegativity value of all the elements. With an increase in atomic number, electronegativity decreases within the group. This means that from oxygen to polonium, the metallic character increases.

Physical Properties

Sulphur and oxygen are non-metals, while selenium and tellurium are metalloids, and polonium is a metal. Polonium is a radioactive element with a short half-life (Half-life 13.8 days). All of these elements are allotropic. The melting and boiling points rise as the atomic number decreases in the group. The bigger disparity in melting and boiling temperatures between oxygen and sulphur can be explained by their atomicity; oxygen is made up of diatomic molecules (O2), whereas sulphur is made up of polyatomic molecules (S8).

Atomic & Physical Properties

Chemical Properties

Oxidation states and chemical reactivity trends: Group 16 elements have a variety of oxidation states. The stability of the -2-oxidation state reduces as you progress through the group.

Polonium rarely exhibits oxidation states of -2. Because oxygen's electronegativity is so high, it only has negative oxidation states like -2, except for OF2, which has positive oxidation states like + 2. Other elements in the group have oxidation states of + 2 + 4 + 6, however + 4 and + 6 are more frequent.

In their compounds with oxygen, sulphur, selenium, and tellurium commonly show +4 oxidation and +6 oxidation with fluorine. The stability of the +6-oxidation state reduces as the group progresses, while the stability of the +4- oxidation state rises (inert pair effect). In the + 4 and + 6 oxidation states, most bonds are covalent.

Anomalous Behaviour of Oxygen

Because of its tiny size and high electronegativity, oxygen, like other members of the p-block present in the second period, exhibits unusual behaviour. The occurrence of strong hydrogen bonding in \[{{\text{H}}_2}{\text{O}}\], which is absent in \[{{\text{H}}_2}{\text{S}}\], is an example of the impact of small size and high electronegativity.

Because oxygen lacks d orbitals, its covalency is limited to four, and it rarely exceeds two in practise. In the case of other group members, however, the valence shell can be enlarged, and covalence approaches four.

(i) Reactivity with hydrogen: Group 16 elements all form hydrides of the type \[{{\text{H}}_2}{\text{E}}\] (E = S, Se, Te, Po). Table lists some of the features of hydrides. From \[{{\text{H}}_2}{\text{O}}\] to \[{{\text{H}}_2}{\text{Te}}\], they get increasingly acidic. The decrease in bond (H-E) dissociation enthalpy down the group can explain the increase in acidic character. The thermal stability of hydrides falls from \[{{\text{H}}_2}{\text{O}}\] to \[{{\text{H}}_2}{\text{Po}}\] as the bond (H-E) dissociation enthalpy lowers down the group. Except for water, all hydrides have a decreasing property, which increases from \[{{\text{H}}_2}{\text{S}}\] to \[{{\text{H}}_2}{\text{Te}}\].

Table: Properties of Hydrides of Group 16 Elements

(ii) Reactivity with oxygen: S, Se, Te, or Po. Sulfur dioxide (SO2) and ozone (O3) are gases, but selenium dioxide \[{\text{(Se}}{{\text{O}}_2})\]is a solid. From \[{\text{S}}{{\text{O}}_2}\] to \[{\text{Te}}{{\text{O}}_2}\], the reducing property of dioxide declines; \[{\text{S}}{{\text{O}}_2}\] is a reducing agent, whereas TeO2 is an oxidising agent. Selenium and tellurium, in addition to \[{\text{E}}{{\text{O}}_2}\]type sulphur, generate \[{\text{E}}{{\text{O}}_3}\] type oxides \[{\text{(S}}{{\text{O}}_3},Se{O_3}{\text{ and Te}}{{\text{O}}_3})\]. Both forms of oxides have an acidic pH. Group 16 elements form a greater number of halides of the types \[{\text{E}}{{\text{X}}_6},{\text{E}}{{\text{X}}_4},{\text{ and E}}{{\text{X}}_2}\], where \[{\text{E}}\] is a group element and \[{\text{X}}\] is a halogen. The halides' stabilities decline in the following order: \[{\text{F > Cl > Br > I}}\]. Hexafluorides are the only stable halides among the hexahalides. In nature, all hexafluorides are gaseous. They are octahedral in shape. For steric reasons, shulphur hexafluoride \[{\text{S}}{{\text{F}}_6}\]is extremely stable. \[{\text{S}}{{\text{F}}_4}\] is a gas, \[{\text{Se}}{{\text{F}}_4}\] is a liquid, and \[{\text{Te}}{{\text{F}}_4}\] is a solid among the terrafluorides. These fluorides contain sp3d hybridisation and so have a trigonal bipyramidal structure with a lone pair of electrons in one of the equatorial positions. See-saw geometry is another name for this type of geometry.

Except for selenium, all elements create dichlorides and dibromides. The tetrahedral structure of these dihalides is due to sp3 hybridisation. \[{{\text{S}}_2}{{\text{F}}_2},{\text{ }}{{\text{S}}_2}{\text{C}}{{\text{l}}_2},{\text{ }}{{\text{S}}_2}{\text{B}}{{\text{r}}_2},{\text{ S}}{{\text{e}}_2}{\text{C}}{{\text{l}}_2}{\text{ and S}}{{\text{e}}_2}{\text{B}}{{\text{r}}_2}\]are examples of well-known monohalides that are dimeric in nature. The disproportionation of these dimeric halides is as follows:

\[2{\text{S}}{{\text{e}}_2}{\text{C}}{{\text{l}}_2} \to {\text{SeC}}{{\text{l}}_4}{\text{ + 3Se}}\]

1. Dioxygen (\[{{\text{O}}_2}\]):

It is distinguished from the other members of the VIth group by the following characteristics.

(A) small size 

(B) high electronegativity and 

(C) non-availability of d-orbitals.

Preparation

(i) By the thermal breakdown of metal oxides.

$2{\text{HgO}}\xrightarrow{{{{450}^0}{\text{C}}}}2{\text{Hg + }}{{\text{O}}_2} \ \\ $

$2{\text{A}}{{\text{g}}_2}{\text{O}}\xrightarrow{{{{350}^0}{\text{C}}}}4{\text{Ag + }}{{\text{O}}_2} \ \\ $

$  3{\text{Mn}}{{\text{O}}_2}\xrightarrow{\Delta }{\text{M}}{{\text{n}}_3}{{\text{O}}_4} + {{\text{O}}_2} \ \\ $

$  2{\text{P}}{{\text{b}}_3}{{\text{O}}_4}\xrightarrow{\Delta }6{\text{PbO + }}{{\text{O}}_2} \ \\ $ 

(ii) Oxygen-rich compounds decompose thermally.

$2{\text{NaN}}{{\text{O}}_3}\xrightarrow{\Delta }2{\text{NaN}}{{\text{O}}_2} + {{\text{O}}_2} \ \\ $

$  2{\text{KCl}}{{\text{O}}_3}\xrightarrow{\Delta }2{\text{KCl + 3}}{{\text{O}}_2}({\text{laboratory method)}} \ \\ $

${\text{4}}{{\text{K}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_7}\xrightarrow{\Delta }4{{\text{K}}_2}{\text{Cr}}{{\text{O}}_4} + 2{\text{C}}{{\text{r}}_2}{{\text{O}}_3} + 3{{\text{O}}_2} \ \\ $

$  2{\text{KMn}}{{\text{O}}_4}\xrightarrow{\Delta }{{\text{K}}_2}{\text{Mn}}{{\text{O}}_4} + {\text{Mn}}{{\text{O}}_2} + {{\text{O}}_2} \ \\ $ 

(iii) By the action of conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] on \[{\text{Mn}}{{\text{O}}_2}\].

\[2{\text{Mn}}{{\text{O}}_4}{\text{ + 2}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{2 MnS}}{{\text{O}}_4} + 2{{\text{H}}_2}{\text{O + }}{{\text{O}}_2}\]

(iv) By the action of water on \[{\text{N}}{{\text{a}}_2}{{\text{O}}_2}\].

\[{\text{2N}}{{\text{a}}_2}{{\text{O}}_2} + 2{{\text{H}}_2}{\text{O}} \to {\text{4NaOH + }}{{\text{O}}_2}\]

(v) By the action of conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] on \[{\text{KMn}}{{\text{O}}_4}\] or \[{{\text{K}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_7}\].

$2{\text{KMn}}{{\text{O}}_4}{\text{ + 6}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {{\text{K}}_2}{\text{S}}{{\text{O}}_4} + 4{\text{MnS}}{{\text{O}}_4} + 6{{\text{H}}_2}{\text{O}} + 5{{\text{O}}_2} \ \\ $

$  {\text{2}}{{\text{K}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_7} + 8{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to 4{{\text{K}}_2}{\text{S}}{{\text{O}}_4} + 2{\text{C}}{{\text{r}}_2}{{\text{(S}}{{\text{O}}_4})_3} + 8{{\text{H}}_2}{\text{O}} + 3{{\text{O}}_2} \ \\ $ 

(vi) By Brins process (mfg.)

$2{\text{BaO + }}{{\text{O}}_2}({\text{air)}}\xrightarrow{{{{500}^0}{\text{C}}}}2{\text{Ba}}{{\text{O}}_2} \ \\ $

$  2{\text{Ba}}{{\text{O}}_2}{\text{ + }}{{\text{O}}_2}({\text{air)}}\xrightarrow{{{{800}^0}{\text{C}}}}2{\text{Ba}}{{\text{O}}_2} + {{\text{O}}_2} \ \\ $ 

(vii) From air (mfg.)

Liquefaction of air followed by fractional distillation yields oxygen.

Properties

Gas that is colourless, odourless, and tasteless. It has allotropy and is paramagnetic. \[{}_8^{16}{\text{O, }}{}_8^{17}{\text{O and }}{}_8^{18}{\text{O}}\] are three oxygen isotopes. Although oxygen does not burn, it is a powerful promoter of combustion.

Uses

(i) For artificial respiration, oxygen combined with helium or CO2 is employed.

(ii) In rocket fuels, liquid oxygen is used as an oxidising agent.

(iii) Oxygen is utilised to create oxy-hydrogen or oxy-acetylene flames, which are used for cutting and welding.

2. Oxides

(i) Acidic Oxides

They generate oxyacids when they dissolve in water, e.g., \[{\text{C}}{{\text{O}}_2}{\text{, S}}{{\text{O}}_2}{\text{, S}}{{\text{O}}_3}{\text{, }}{{\text{N}}_2}{{\text{O}}_5}{\text{, }}{{\text{N}}_2}{{\text{O}}_3}{\text{, }}{{\text{P}}_4}{{\text{O}}_6}{\text{, }}{{\text{P}}_4}{{\text{O}}_{10}}{\text{, C}}{{\text{l}}_2}{{\text{O}}_7}{\text{, Cr}}{{\text{O}}_3}{\text{, M}}{{\text{n}}_2}{{\text{O}}_7}{\text{, }}{{\text{V}}_2}{{\text{O}}_5}\]

$ {\text{C}}{{\text{l}}_2}{{\text{O}}_7}{\text{ + }}{{\text{H}}_2}{\text{O}} \to 2{\text{HCl}}{{\text{O}}_4} \ \\ $

$  {\text{M}}{{\text{n}}_2}{{\text{O}}_7}{\text{ + }}{{\text{H}}_2}{\text{O}} \to 2{\text{HMn}}{{\text{O}}_4} \ \\ $ 

(ii) Basic Oxides

They either dissolve in water to form alkalies, react with acids to form salts, or react with acidic oxides to form salts. e.g., \[{\text{N}}{{\text{a}}_2}{\text{O, CaO, CuO, FeO, BaO}}\]etc.,

$ {\text{N}}{{\text{a}}_2}{\text{O + }}{{\text{H}}_2}{\text{O}} \to {\text{2NaOH}} \ \\ $

$  {\text{CaO + }}{{\text{H}}_2}{\text{O}} \to {\text{2Ca(OH}}{{\text{)}}_2} \ \\ $

$  {\text{CuO + }}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{CuS}}{{\text{O}}_4}{\text{ + }}{{\text{H}}_2}{\text{O}} \ \\ $ 

(iii) Neutral Oxides

They do not produce salts when combined with acids or bases. E.g., \[{\text{CO, }}{{\text{N}}_2}{\text{O, NO}}\]

(iv) Amphoteric Oxides

These can react with both acids and bases, for example. \[{\text{ZnO, A}}{{\text{l}}_2}{{\text{O}}_3}{\text{ , BeO, S}}{{\text{b}}_2}{{\text{O}}_3}{\text{ , C}}{{\text{r}}_2}{{\text{O}}_3}{\text{ , PbO}}\] etc.,

$  {\text{PbO  +  2NaOH }} \to {\text{ N}}{{\text{a}}_2}{\text{Pb}}{{\text{O}}_2}{\text{  +  }}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{PbO  +  }}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ }} \to {\text{ PbS}}{{\text{O}}_4}{\text{  +  }}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{C}}{{\text{r}}_2}{{\text{O}}_3}{\text{  +  2NaOH }} \to {\text{ N}}{{\text{a}}_2}{\text{C}}{{\text{r}}_2}{{\text{O}}_4}{\text{  +  }}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{C}}{{\text{r}}_2}{{\text{O}}_3}{\text{  +  3}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ }} \to {\text{ C}}{{\text{r}}_2}{\left( {{\text{S}}{{\text{O}}_4}} \right)_3}{\text{ +  3}}{{\text{H}}_2}{\text{O}} \ \\ $

(v) Mixed Oxides

They have the properties of a combination of two simple oxides.

${\text{P}}{{\text{b}}_3}{{\text{O}}_4}\left( {{\text{2PbO  +  Pb}}{{\text{O}}_2}} \right) \ \\ $

$  {\text{F}}{{\text{e}}_3}{{\text{O}}_4}\left( {{\text{FeO  +  F}}{{\text{e}}_2}{{\text{O}}_3}{\text{ }}} \right) \ \\ $

$  {\text{M}}{{\text{n}}_3}{{\text{O}}_4}\left( {{\text{2MnO  +  Mn}}{{\text{O}}_2}{\text{ }}} \right) \ \\ $

(vi) Peroxides

They create \[{{\text{H}}_2}{{\text{O}}_2}\] when they react with dilute acids. For eg., \[{\text{N}}{{\text{a}}_2}{{\text{O}}_2}{\text{ , }}{{\text{K}}_2}{{\text{O}}_2}{\text{ , Ba}}{{\text{O}}_2}\] etc.,

\[{\text{N}}{{\text{a}}_2}{{\text{O}}_2}{\text{ +  }}{{\text{H}}_2}{\text{S}}{{\text{O}}_4} \to {\text{ N}}{{\text{a}}_2}{\text{S}}{{\text{O}}_4}{\text{ +  }}{{\text{H}}_2}{{\text{O}}_2}\]

They create \[{{\text{O}}_2}\] when they react with water.

\[{\text{N}}{{\text{a}}_2}{{\text{O}}_2}{\text{ +  }}{{\text{H}}_2}{\text{O }} \to {\text{ 2NaOH  +  1/2}}{{\text{O}}_2}\]

(vii) Dioxides

They have an overabundance of oxygen, like peroxide, but do not produce \[{{\text{H}}_2}{{\text{O}}_2}\] when mixed with dilute acids. For example, \[{\text{Pb}}{{\text{O}}_2}{\text{ , Mn}}{{\text{O}}_2}\]etc.,

They produce \[{\text{C}}{{\text{l}}_2}\] from concentrated \[{\text{HCl}}\] and \[{{\text{O}}_2}\] from concentrated \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\].

$  {\text{Mn}}{{\text{O}}_2}{\text{ +  4HCl}} \to {\text{MnC}}{{\text{l}}_2}{\text{ +  C}}{{\text{l}}_2}{\text{ + 2}}{{\text{H}}_2}{\text{O}} \ \\ $

$  {\text{2Mn}}{{\text{O}}_2}{\text{ + 2}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}{\text{ }} \to {\text{2MnS}}{{\text{O}}_4}{\text{ +  }}{{\text{O}}_2}{\text{ + 2}}{{\text{H}}_2}{\text{O}} \ \\ $ 

(viii) Super Oxides

They contain \[{{\text{O}}_2}^ - {\text{ }}\]ions, such as \[{\text{K}}{{\text{O}}_2}{\text{ , Rb}}{{\text{O}}_2}{\text{ and Cs}}{{\text{O}}_2}\]. These oxides react with water to generate \[{{\text{H}}_2}{{\text{O}}_2}{\text{ and }}{{\text{O}}_2}\] compounds.

\[{\text{2 K}}{{\text{O}}_2}{\text{  +  2}}{{\text{H}}_2}{\text{O }} \to {\text{ 2KOH  +  }}{{\text{H}}_2}{{\text{O}}_2}{\text{  +  }}{{\text{O}}_2}\]

(ix) Sub Oxides

They have less oxygen than would be expected based on the element's typical valency, for example, \[{{\text{C}}_3}{{\text{O}}_2}{\text{ , }}{{\text{N}}_2}{\text{O, P}}{{\text{b}}_2}{\text{O, H}}{{\text{g}}_2}{\text{O}}\]

$ {{\text{C}}_3}{{\text{O}}_2} \to {\text{ O = C = C = C = O}} \ \\ $

$  {\text{                    sp  sp  sp  }} \ \\ $ 

2. Ozone (\[{{\text{O}}_3}\])

Preparation

It's made by sending a quiet electric discharge through dry, pure oxygen.

$  {{\text{O}}_2}\xrightarrow{{{\text{Energy}}}}{\text{ O  +  O}} \ \\ $

$  {{\text{O}}_2}{\text{  +  O }} \to {\text{ }}{{\text{O}}_3} \ \\ $

$  \Delta {\text{H  =  2845 kJ mo}}{{\text{l}}^{{\text{ - 1}}}} \ \\ $ 

\[{\text{3}}{{\text{O}}_{\text{2}}} \rightleftharpoons {\text{2}}{{\text{O}}_{\text{3}}}\]

The resulting mixture is called ozonised oxygen since it contains 5-10% ozone by volume. The device used for this is known as an ozoniser.

(i) Simen’s and (ii) Brodie’s ozonisers

Properties

(i) Pale blue gas that cools into a blue liquid and solidifies into violet black crystals. It has a distinct fishy odour and is soluble in turpentine oil, glacial acetic acid, or \[{\text{CC}}{{\text{l}}_{{\text{4 }}}}\] but not in water. Although the \[{{\text{O}}_{\text{3}}}\]molecule is diamagnetic, the \[{{\text{O}}_{\text{3}}}^{{\text{-- }}}\] molecule is paramagnetic.

(i) Oxidising agent

$  {{\text{O}}_{\text{3}}}{\text{  +  2}}{{\text{H}}^{\text{ + }}}{\text{  +  2}}{{\text{e}}^ - }{\text{ }} \to {{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O; SRP  =   +  2}}{\text{.07 v}} \ \\ $

$  {{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  2}}{{\text{e}}^ - } \to {\text{ }}{{\text{O}}_{\text{2}}}{\text{  +  2O}}{{\text{H}}^{\text{--}}};{\text{SRP  =   +  1}}{\text{.24 v}} \ \\ $ 

As a result, ozone is a powerful oxidising agent in acidic environments.

(a) It oxidises — to \[{{\text{I}}_{\text{2}}}\] (from a neutral K solution).

\[{{\text{O}}_{\text{3}}}{\text{ }} \to {\text{ }}{{\text{O}}_{\text{2}}}{\text{  +  }}\left[ {\text{O}} \right]\]

$  {\text{Kn  +  3}}{{\text{O}}_{\text{3}}} \to {\text{ Kn}}{{\text{O}}_{\text{3}}}{\text{  +  3}}{{\text{O}}_{\text{2}}} \ \\ $

$  {\text{Kn  +  4}}{{\text{O}}_{\text{3}}}{\text{ }} \to {\text{ Kn}}{{\text{O}}_{\text{4}}}{\text{  +  4}}{{\text{O}}_{\text{2}}} \ \\ $ 

Similarly, \[{{\text{S}}^{{\text{2-- }}}}{\text{ S}}{{\text{O}}_{{\text{4 }}}}^{{\text{2--}}}\], (but not \[{{\text{H}}_2}{\text{S}}\]),\[{\text{N}}{{\text{O}}^{{\text{2 --}}}}{\text{,N}}{{\text{O}}^{{\text{3 -- }}}}{\text{, S}}{{\text{O}}^{{\text{3 --}}}}{\text{, S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{, As}}{{\text{O}}_{\text{3}}}^{{\text{3--}}}{\text{, As}}{{\text{O}}_{\text{4}}}^{{\text{3--}}}{\text{ , Mn}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{,Mn}}{{\text{O}}_{\text{4}}}^{\text{--}}{\text{ , S}}{{\text{n}}^{{\text{2 + }}}}\]\[ \to {\text{ S}}{{\text{n}}^{{\text{4 + }}}}{\text{ }}{\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]^{{\text{4--}}}} \to {\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]^{{\text{3-- }}}}\]in acidic medium.

(b) It converts wet S, P, and As to oxy acids.

\[{{\text{O}}_{\text{3}}}{\text{ }} \to {{\text{O}}_{{\text{2 }}}}{\text{ + }}\left[ {\text{O}} \right]{\text{  \times  3}}\]

\[{\text{S  +  3}}\left[ {\text{O}} \right] \to {\text{S}}{{\text{O}}_{\text{3}}}\]

\[\dfrac{{{\text{S}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}} \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4}}}{{{\text{S + 3}}{{\text{O}}_3} + {{\text{H}}_2}{\text{O}} \to {{\text{H}}_2}{\text{S}}{{\text{O}}_4} + 3{{\text{O}}_2}}}\]

(c) It oxidises \[{{\text{H}}_2}{\text{S}}\] to \[{\text{S}}\]

\[{{\text{H}}_2}{\text{S +  }}{{\text{O}}_{\text{3}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S }} \downarrow \]

(ii) Reaction with dry \[{{\text{I}}_2}\]

\[{\text{2}}{{\text{I}}_{{\text{2 }}}}{\text{ +  9}}\left[ {{{\text{O}}_{\text{3}}}} \right] \to {{\text{I}}_{\text{4}}}{{\text{O}}_{\text{9}}}{\text{  +  9}}{{\text{O}}_{\text{2}}}\]

\[{{\text{I}}^{{\text{ + 3 }}}}{\left( {{\text{I}}{{\text{O}}_{\text{3}}}^{{\text{-- }}}} \right)_{\text{3}}}\] is the composition of \[{{\text{I}}_{\text{4}}}{{\text{O}}_{\text{9}}}{\text{ }}\] yellow solid. 3. The formation of this molecule is concrete evidence for I2's basic nature (i.e. its tendency to form cations)

(iii) Reaction with moist iodine

$  {{\text{O}}_{\text{3}}} \to {{\text{O}}_{\text{2}}}{\text{ +  }}\left[ {\text{O}} \right]{\text{  \times  5}} \ \\ $

$  {{\text{I}}_{\text{2}}}{\text{  +  5}}\left[ {\text{O}} \right]{\text{ }} \to {{\text{I}}_{\text{2}}}{{\text{O}}_{\text{5}}} \ \\ $ 

\[\dfrac{{{{\text{I}}_2}{{\text{O}}_5} + {{\text{H}}_2}{\text{O}} \to 2{\text{HI}}{{\text{O}}_3}}}{{{\text{5}}{{\text{O}}_3} + {\text{I + }}{{\text{H}}_2}{\text{O}} \to 2{\text{HI}}{{\text{O}}_3} + 5{{\text{O}}_2}}}\]

(iv) Reaction with Silver

When silver objects come into touch with ozone, they turn black.

\[{\text{Ag  +  }}{{\text{O}}_{\text{3}}} \to {\text{A}}{{\text{g}}_{\text{2}}}{\text{O}} \downarrow \left( {{\text{black}}} \right){\text{  +  }}{{\text{O}}_{\text{2}}}\]

(v) Reaction with \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\]

\[{{\text{2}}^{{\text{e-- }}}}{\text{ +  }}{{\text{2}}^{{\text{H +  }}}}{\text{ +  }}{{\text{O}}_{\text{3}}} \to {{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O}}\]

\[\dfrac{{{{\text{H}}_2}{{\text{O}}_2} \to {{\text{O}}_2} + 2{{\text{H}}^ + }{\text{ + 2}}{{\text{e}}^ - }}}{{{{\text{O}}_3} + {{\text{H}}_2}{{\text{O}}_2} \to 2{{\text{O}}_2} + {{\text{H}}_2}{\text{O}}}}\]

The fact that the SRP of ozone is higher (+2.07) than the SRP of hydrogen peroxide (+1.77) supports this theory. As a result, ozone is a more powerful oxidizer than hydrogen peroxide.

(vi) Bleaching Action

Through oxidation, \[{{\text{O}}_3}\] also bleaches coloured things.

(vii) Ozonolysis

Ozonides are formed when alkenes and alkynes react with ozone.

(viii) Reaction with \[{\text{KOH}}\]

Potassium ozonide, an orange-colored compound, is formed.

\[{\text{2 KOH  +  5}}{{\text{O}}_{\text{3}}} \to {\text{2K}}{{\text{O}}_3}^ - {\text{ + 5}}{{\text{O}}_2} + {{\text{H}}_2}{\text{O}}\]

Tests for Ozone

(i) When a filter paper soaked in alcoholic benzidine comes into touch with \[{{\text{O}}_{\text{3}}}\], it turns brown.

(ii) Tailing of mercury

Due to the breakdown of \[{\text{H}}{{\text{g}}_{\text{2}}}{\text{O}}\] (mercury sub-oxide) in Hg, its mobility slows when it meets \[{{\text{O}}_3}\], and it begins to attach to the glass surface, forming a type of tail.

\[{\text{2 Hg  +  }}{{\text{O}}_3} \to {\text{H}}{{\text{g}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_{\text{2}}}\]

Uses

(i)For sterilising water and enhancing the atmosphere of crowded locations as a germicide and disinfectant.

(ii) Detecting the double bond position in unsaturated chemical molecules.

(iii) In the manufacture of artificial silk, synthetic camphor, \[{\text{KMn}}{{\text{O}}_4}\], and other similar materials.

3. Hydrogen Peroxide (H2O2) 

Preparation 

(i) Laboratory method

\[{\text{Ba}}{{\text{O}}_{\text{2}}}{\text{.8}}{{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ }}\left( {{\text{cold}}} \right) \to {\text{BaS}}{{\text{O}}_{\text{4}}} \downarrow \left( {{\text{white}}} \right){\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  8}}{{\text{H}}_{\text{2}}}{\text{O}}\]

Aqueous hydrogen peroxide is obtained by filtering \[{\text{BaS}}{{\text{O}}_{\text{4}}}\].

$ {\text{Ba}}{\left( {{\text{OH}}} \right)_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  6}}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{Ba}}{{\text{O}}_{\text{2}}}{\text{.8}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{Ba}}{{\text{O}}_{\text{2}}}{\text{ +  2HCl }}\left( {{\text{ice cold}}} \right) \to {\text{BaC}}{{\text{l}}_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \ \\ $ 

Because \[{\text{BaC}}{{\text{l}}_{{\text{2 }}}}\] is soluble in water, it is impossible to extract \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] from this solution.

Due to the creation of a protective layer of \[{\text{BaS}}{{\text{O}}_{\text{4}}}\]on \[{\text{Ba}}{{\text{O}}_{\text{2}}}\], the interaction between anhydrous \[{\text{Ba}}{{\text{O}}_{\text{2}}}\] and \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] is sluggish and eventually stops.

Because \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] can decompose \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] at a higher temperature, the reaction should take place at a low temperature, or \[{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}}\] can be substituted for \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\].

$  {\text{3Ba}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}} \to {\text{B}}{{\text{a}}_{\text{3}}}\left( {{\text{P}}{{\text{O}}_{\text{4}}}} \right){\text{2}} \downarrow {\text{ +  3}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ }} \ \\ $

$  {\text{B}}{{\text{a}}_{\text{3}}}{\left( {{\text{P}}{{\text{O}}_{\text{4}}}} \right)_{\text{2}}}{\text{ +  3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{3BaS}}{{\text{O}}_{\text{4}}} \downarrow {\text{ +  2}}{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}} \ \\ $ 

\[{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}}\] can be utilised once more.

(ii) Using inert electrodes to electrolyze conc. \[{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\] at 00C (platinum)

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \rightleftharpoons {{\text{H}}^ \oplus }{\text{3HS}}{{\text{O}}_{\text{4}}}^ - \]

\[{{\text{H}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{8}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O }}\dfrac{{80 - {{90}^0}{\text{C}}}}{{{\text{distillation}}}}{\text{2}}{{\text{H}}_2}{\text{S}}{{\text{O}}_4}\xrightarrow{{{\text{BaC}}{{\text{l}}_{\text{2}}}}}{\text{BaS}}{{\text{O}}_{{\text{4 }}}} \downarrow \left( {{\text{white}}} \right){\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\left( {{\text{aq}}} \right)\]

Filtration removes \[{\text{BaS}}{{\text{O}}_{{\text{4 }}}}\] to yield aqueous \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\].

(iii) Industrial method (Auto oxidation)

PROPERTIES

(i) A colourless viscous liquid that appears blue in large amounts and is soluble in water (owing to H-bonding) in all quantities, forming the hydrate \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\]. \[{{\text{H}}_{\text{2}}}{\text{O}}\] (mp 221 K) 

\[{{\text{H}}_{\text{2}}}{\text{O}}\] (mp 221 K)  The dielectric constant and density are also higher than in \[{{\text{H}}_{\text{2}}}{\text{O}}\].

(iii) Its aqueous solution is more stable than the anhydrous liquid, which decomposes slowly into water and oxygen when exposed to light.

\[{\text{2}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ }} \to {\text{2}}{{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_2}\]

Because residues of alkali metal ions from the glass can catalyse the explosive disintegration of \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\], it is not kept in glass containers.

As a result, aqueous solutions are maintained in plastic or wax-lined glass containers with urea, phosphoric acid, or glycerol added since these substances have been discovered to act as negative catalysts in the decomposition of \[{{\text{H}}_{\text{2}}}{\text{O }}\].

According to the following equation, it behaves as a weak acid.

$ {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\left( {{\text{aq}}} \right){\text{ }}{{\text{H}}^{\text{ + }}} \rightleftharpoons {\text{H}}{{\text{O}}_2}^ - {\text{ }} \ \\ $

$  {{\text{K}}_{\text{a}}}{\text{  =  1}}{\text{.5  \times  1}}{{\text{0}}^{{\text{ - 12}}}}{\text{ at 2}}{{\text{5}}^{\text{0}}}{\text{ C}} \ \\ $ 

\[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] in an aqueous solution turns blue litmus crimson, which is then bleached by \[{{\text{H}}_{\text{2}}}{\text{O }}\]'s oxidising property.

$  {\text{N}}{{\text{a}}_{\text{2}}}{\text{C}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{N}}{{\text{a}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{O}}_{\text{2}}} \ \\ $

$  {\text{Ba}}{\left( {{\text{OH}}} \right)_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{  +  6}}{{\text{H}}_{\text{2}}}{\text{O }} \to {\text{Ba}}{{\text{O}}_{\text{2}}}{\text{ }}{\text{. 8}}{{\text{H}}_{\text{2}}}{\text{O }} \downarrow  \ \\ $ 

The pH of a 30 percent \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] solution is 4.

(v) Oxidising Agent

$  {\text{2 }}{{\text{e}}^ - }{\text{  +  2}}{{\text{H}}^{\text{ + }}}{\text{  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ }} \to {\text{2}}{{\text{H}}_{\text{2}}}{\text{O ; SRP  =   +  1}}{\text{.77 v}} \ \\ $

$  {\text{2 }}{{\text{e}}^ - }{\text{  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2O}}{{\text{H}}^{\text{ - }}}{\text{ ; SRP  =   +  0}}{\text{.87 v}} \ \\ $ 

We can conclude that \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] is a strong oxidising agent in acidic medium based on the above potentials, but that reactions in basic medium are faster kinetically.

(A) In Acidic Medium

(a) It oxidises \[{\text{PbS}}\] to \[{\text{PbS}}{{\text{O}}_{\text{4}}}\]

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  }}\left[ {\text{O}} \right]{\text{  \times  4}} \ \\ $

$  {\text{PbS  +  4}}\left[ {\text{O}} \right] \to {\text{PbS}}{{\text{O}}_{\text{4}}}{\text{ }} \ \\ $

$  {\text{PbS  +  4}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{PbS}}{{\text{O}}_{\text{4}}}{\text{  +  4}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

This feature is used to restore the white colours in old paintings that have turned black because of atmospheric \[{{\text{H}}_{\text{2}}}{\text{S}}\] becoming \[{\text{PbS}}\].

(b) \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] converts \[{{\text{H}}_{\text{2}}}{\text{S}}\] to sulphur.

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  }}\left[ {\text{O}} \right] \ \\ $

$  {\text{H2S  +  }}\left[ {\text{O}} \right]{\text{ }} \to {{\text{H}}_{\text{2}}}{\text{O  +  S}} \downarrow  \ \\ $

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {\text{2}}{{\text{H}}_{\text{2}}}{\text{O  +  S}} \downarrow  \ \\ $ 

\[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\] in acidic do oxidise \[{\text{As}}{{\text{O}}_{\text{3}}}^{{\text{3 - }}}{\text{ }} \to {\text{As}}{{\text{O}}_{\text{4}}}^{{\text{3 - }}}{\text{, S}}{{\text{O}}_{\text{3}}}^{{\text{2 -  }}} \to {\text{ S}}{{\text{O}}_{{\text{4 }}}}^{{\text{2 -  }}}{\text{, KI }} \to {{\text{I}}_{\text{2}}}{\text{, }}{{\text{S}}^{{\text{2 -  }}}} \to {\text{S}}{{\text{O}}_{\text{4}}}^{{\text{2 - }}}{\text{,}}\]also.

\[{\text{FeS}}{{\text{O}}_{\text{4}}} \to {\text{F}}{{\text{e}}_{\text{2}}}{\left( {{\text{S}}{{\text{O}}_{\text{4}}}} \right)_{\text{3}}}{\text{ \&  }}{\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{{\text{6 }}}}} \right]^{{\text{4 - }}}} \to {\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]^{{\text{3 - }}}}\]

(c) \[{\text{N}}{{\text{H}}_{\text{2}}}{\text{  -  N}}{{\text{H}}_{\text{2}}}\left( {{\text{hydrazine}}} \right){\text{ +  2}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {{\text{N}}_{\text{2}}}{\text{ +  4}}{{\text{H}}_{\text{2}}}{\text{O}}\]

d)

(B) In Alkaline Medium

(a) \[{\text{Cr}}{\left( {{\text{OH}}} \right)_{\text{3}}}\left( {\text{s}} \right){\text{ +  4NaOH  +  3}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2N}}{{\text{a}}_{\text{2}}}{\text{Cr}}{{\text{O}}_{\text{4}}}\left( {{\text{aq}}{\text{.}}} \right){\text{ +  8}}{{\text{H}}_{\text{2}}}{\text{O}}\]

or

\[{\text{10 O}}{{\text{H}}^{{\text{-- }}}}{\text{ +  3}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ + 2C}}{{\text{r}}^{{\text{3 + }}}} \to {\text{2Cr}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{ +  8}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(b) \[{\text{2NaB}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  6}}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{N}}{{\text{a}}_{\text{2}}}\left[ {{{\left( {{\text{OH}}} \right)}_{{\text{2 }}}}{\text{B}}{{\left( {{\text{O - O}}} \right)}_{\text{2}}}{\text{B}}{{\left( {{\text{OH}}} \right)}_{\text{2}}}} \right]{\text{ 6}}{{\text{H}}_{\text{2}}}{\text{O}}\](sodium per oxoborate)

In washing powder, it's used as a brightener.

(vi) Reducing Agent

It operates as a reducing agent in the presence of a strong oxidising agent.

\[{{\text{H}}_{\text{2}}}{{\text{O}}_{{\text{2 }}}} \to {{\text{O}}_{\text{2}}}{\text{ + 2}}{{\text{H}}^{\text{ + }}}{\text{ + 2}}{{\text{e}}^{\text{ - }}}\]

Its reducing nature is stronger in alkaline solution than in acidic medium.

\[{\text{2 O}}{{\text{H}}^{\text{ - }}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  2}}{{\text{e}}^ - }\]

(a) \[{\text{A}}{{\text{g}}_{\text{2}}}{\text{O }}\] is decomposed into \[{\text{Ag}}\].

\[{\text{A}}{{\text{g}}_{\text{2}}}{\text{O  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2Ag  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_{\text{2}}}\]

(b) It converts \[{{\text{O}}_{\text{3}}}\] to \[{{\text{O}}_{\text{2}}}\].

\[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{O}}_{\text{3}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  2}}{{\text{O}}_{\text{2}}}\]

(c) Ferric cyanide is converted to ferrous cyanide (basic medium)

$ {\text{2 }}{{\text{K}}_{\text{3}}}\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]{\text{  +  2KOH}} \to {{\text{K}}_{\text{4}}}\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]{\text{ + }}{{\text{H}}_{\text{2}}}{\text{O  +  O}} \ \\ $

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{  +  O }} \to {{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_{\text{2}}}{\text{ }} \ \\ $

$  {\text{2}}{{\text{K}}_{{\text{3 }}}}\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]{\text{  +  2KOH  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2}}{{\text{K}}_{\text{4}}}\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}} \right]{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_{\text{2}}} \ \\ $ 

It reduces \[{\text{Mn}}{{\text{O}}_{{\text{4 }}}}^{{\text{ -  }}} \to {\text{ M}}{{\text{n}}^{{\text{2 + }}}}\](acidic medium) 

\[{\text{Mn}}{{\text{O}}_{\text{4}}}^{\text{ - }} \to {\text{Mn}}{{\text{O}}_{\text{2}}}\](basic medium)

\[{\text{OC}}{{\text{l}}^{{\text{-- }}}} \to {\text{C}}{{\text{l}}^{\text{--}}}{\text{ , I}}{{\text{O}}_{\text{4}}}^{\text{--}} \to {\text{I}}{{\text{O}}_{{\text{3 }}}}^{\text{--}}{\text{C}}{{\text{l}}_{\text{2}}} \to {\text{C}}{{\text{l}}^{\text{--}}}\]

Tests For \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\]

(i) with \[{{\text{K}}_{\text{2}}}{\text{C}}{{\text{r}}_{\text{2}}}{{\text{O}}_{\text{7}}}{\text{ }}\]

\[{{\text{K}}_{\text{2}}}{\text{C}}{{\text{r}}_{\text{2}}}{{\text{O}}_{\text{7}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{{\text{4 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2Cr}}{{\text{O}}_{\text{5}}}{\text{  +  5}}{{\text{H}}_{\text{2}}}{\text{O}}\]

or

\[{\text{C}}{{\text{r}}_{\text{2}}}{{\text{O}}_{\text{7}}}^{{\text{2 -  }}}{\text{ +  2}}{{\text{H}}^{\text{ + }}}{\text{  +  4}}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2Cr}}{{\text{O}}_{\text{5}}}{\text{ + 5}}{{\text{H}}_{\text{2}}}{\text{O}}\]

\[{\text{Cr}}{{\text{O}}_{\text{5}}}\] is a vivid blue chemical that is soluble in ether.

\[{\text{Cr}}{{\text{O}}_{\text{5}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{{\text{4 }}}} \to {\text{2C}}{{\text{r}}_{\text{2}}}{\left( {{\text{S}}{{\text{O}}_{\text{4}}}} \right)_{\text{3}}}{\text{ +  6}}{{\text{H}}_{\text{2}}}{\text{O  +  7}}{{\text{O}}_{\text{2}}}\]

(ii) \[{\text{2 HCHO  +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\dfrac{{{\text{ O}}{{\text{H}}^ - }}}{{{\text{pyrogallol}}}}{\text{ > 2HCOOH  +  }}{{\text{H}}_{\text{2}}}\]

When this reaction is carried out in the dark, it causes light to be emitted (yellow coloured). Chemiluminescene is a type of chemiluminescene.

(iii) With \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}\], an acidified titanium salt solution turns yellow or orange.

\[{\text{T}}{{\text{i}}^{{\text{ + 4}}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O }} \to {{\text{H}}_{\text{2}}}{\text{Ti}}{{\text{O}}_{\text{4}}}\left( {{\text{yellow/orange}}} \right){\text{ +  4}}{{\text{H}}^{\text{ + }}}\]

Uses

(i) When bleaching delicate materials such as silk, wool, cotton, ivory, and other similar materials.

(ii) Perhydrol, a good antiseptic and germicide used to clean wounds, teeth, and ears.

(iii) As an antichlor to eliminate chlorine and hypochlorite residues.

(iv) In rocket fuels as an oxidising agent 

4. Sulphur (S)

Sulphur Allotropic Forms

Sulphur exists in a variety of allotropes, the most prominent of which are yellow rhombic (- sulphur) and monoclinic (- sulphur). At ambient temperature, rhombic sulphur is stable, but when heated over 369 K, it transforms into monoclinic sulphur.

Rhombic Sulphur (α - sulphur)

This allotrope has a yellow colour, a melting point of 385.8 K, and a specific gravity of 2.06. When roll sulphur solution in \[{\text{C}}{{\text{S}}_{\text{2}}}\] is evaporated, rhombic sulphur crystals emerge. It is insoluble in water but soluble in benzene, alcohol, and ether to some extent. In \[{\text{C}}{{\text{S}}_{\text{2}}}\], it dissolves quickly.

Monoclinic Sulphur (β - sulphur)

It has a melting point of 393 K and a specific gravity of 1.98. It dissolves in CS2. Sulphur in this form is made by melting rhombic sulphur in a dish and chilling it until a crust forms. The remaining liquid is poured out through two holes in the crust. Colorless needle-shaped crystals of - sulphur occur after the crust is removed. Above 369 K, it is stable, but below that temperature, it converts into - sulphur. - sulphur, on the other hand, is stable below 369 K and transforms into - sulphur above that temperature. Both types are stable around 369 K. This is referred to as the transi

S8 molecules can be found in both rhombic and monoclinic sulphur. These S8 molecules are crammed together to form a variety of crystal shapes. In both variants, the S8 ring is puckered and has a crown shape. Figure 1 shows the molecular dimensions temperature.

In the previous two decades, several additional sulphur modifications with 6-20 sulphur atoms per ring have been synthesised.

The ring in cyclo- \[{{\text{S}}_6}\] has a chair shape, and the molecular dimensions are as indicated in fig. (b) \[{{\text{S}}_2}\] is the dominating species at high temperatures (> 1000 K) and is paramagnetic like \[{{\text{O}}_2}\].

3. Compounds of Sulphur 

(A) SODIUM THIOSULPHATE (\[{\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ }}{\text{.5}}{{\text{H}}_{\text{2}}}{\text{O}}\]) 

 PREPARATION

(i) \[{\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  S}}\dfrac{{{\text{boiled}}}}{{{\text{in absence of air}}}}{\text{ N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}\]

(ii) $  {\text{N}}{{\text{a}}_{\text{2}}}{\text{C}}{{\text{O}}_{\text{3}}}{\text{  +  2SO2 }}\left( {{\text{excess}}} \right){\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{2NaHS}}{{\text{O}}_{\text{3}}}{\text{ +  C}}{{\text{O}}_{\text{2}}}{\text{ ;}} \ \\ $

$  {\text{2NaHS}}{{\text{O}}_{\text{3}}}{\text{  +  N}}{{\text{a}}_{\text{2}}}{\text{C}}{{\text{O}}_{\text{3}}} \to {\text{2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{O}}_{\text{2}}}{\text{ }} \ \\ $ 

(iii) \[{\text{2 NaHS  +  4NaHS}}{{\text{O}}_{\text{3}}} \to {\text{3N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ + 3}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(iv) \[{\text{N}}{{\text{a}}_{\text{2}}}{\text{S  +  N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{{\text{3 }}}}{\text{ +  }}{{\text{I}}_{\text{2}}} \to {\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  2NaI}}\]

(v) \[{\text{2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{3}}}{\text{  +  3}}{{\text{O}}_{\text{2}}}{\text{ }}\left( {{\text{from air}}} \right)\xrightarrow{\Delta }{\text{ 2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  2S}}\]

Properties

(i) It's a colourless crystalline substance that's soluble in water and loses water of crystallisation when heated to high temperatures.

(ii) As antichlor

According to the following reaction, it eliminates chlorine from the surface of fibres (during dyeing).

\[{\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  C}}{{\text{l}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  2HCl  +  S}}\]

As a result, it is referred to as antichlor.

(iii) Reaction with HCl

\[{\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  HCl}} \to {\text{2NaCl  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  S  +  }}{{\text{H}}_{\text{2}}}{\text{O}}\]

This test helps in distinguishing between \[{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}^{{\text{2 - }}}{\text{ and S}}{{\text{O}}_{\text{3}}}^{{\text{2 - }}}{\text{ }}\]

(iv) Complex formation reactions

(a) Reaction with silver salts \[\left( {{\text{AgN}}{{\text{O}}_{\text{3}}}{\text{ , AgCl, AgBr or AgI}}} \right)\]

${\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  2AgN}}{{\text{O}}_{\text{3}}} \to {\text{A}}{{\text{g}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \downarrow \left( {{\text{white}}} \right){\text{  +  2NaN}}{{\text{O}}_{\text{3}}} \ \\ $

${\text{A}}{{\text{g}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O }} \to {\text{A}}{{\text{g}}_{\text{2}}}{\text{S }} \downarrow \left( {{\text{Black}}} \right){\text{ + }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \ \\ $ 

If you have too much hypo, you'll get a soluble complex.

\[{\text{2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  AgN}}{{\text{O}}_{\text{3}}} \to {\text{N}}{{\text{a}}_{\text{3}}}\left[ {{\text{Ag}}{{\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}} \right)}_{\text{2}}}} \right]\left( {{\text{soluble complex}}} \right){\text{ +  NaN}}{{\text{O}}_{\text{3}}}\]

When hypo is employed as a fixer in photography, this reaction is used.

(b) Reaction with \[{\text{FeC}}{{\text{l}}_3}\]

It takes on a pink or violet hue, which fades quickly as the reaction progresses.

$  {\text{F}}{{\text{e}}^{{\text{3 + }}}}{\text{ +  2}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}^{{\text{2--}}}{\left[ {{\text{Fe}}{{\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ }}} \right)}_{{\text{2 }}}}} \right]^{\text{--}}} \ \\ $

$  {\left[ {{\text{Fe}}{{\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{{\text{3 }}}}} \right)}_{\text{2}}}} \right]^{\text{--}}}{\text{ +  F}}{{\text{e}}^{{\text{3 +  }}}} \to {\text{2F}}{{\text{e}}^{{\text{2 + }}}}{\text{ +  }}{{\text{S}}_{\text{4}}}{{\text{O}}_{{\text{6 }}}}^{{\text{2--}}} \ \\ $ 

(c) Reaction with \[{\text{AuC}}{{\text{l}}_{\text{3}}}{\text{ }}\]

$ {\text{AuC}}{{\text{l}}_{\text{3}}}{\text{  +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{AuCl}} \downarrow {\text{ +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{4}}}{{\text{O}}_{\text{6}}}{\text{  +  2HCl}} \ \\ $

$  {\text{AuCl  +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{N}}{{\text{a}}_{\text{3}}}{\text{ }}\left[ {{\text{Au}}{{\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}} \right)}_{\text{2}}}} \right]\left( {{\text{soluble complex}}} \right){\text{  +  NaCl}} \ \\ $ 

(d) Reaction with CuCl2

$  {\text{2CuC}}{{\text{l}}_{\text{2}}}{\text{ + 2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{2CuCl}} \downarrow {\text{ + N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{4}}}{{\text{O}}_{\text{6}}}{\text{ + 2NaCl}} \ \\ $

$  {\text{CuCl  +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{C}}{{\text{u}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \downarrow {\text{ +  2NaCl}} \ \\ $

  ${\text{3C}}{{\text{u}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  2N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{N}}{{\text{a}}_{\text{4}}}\left[ {{\text{C}}{{\text{u}}_{\text{6}}}\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}} \right){\text{5}}} \right]\left( {{\text{soluble complex}}} \right) \ \\ $ 

(e) Reaction with Bismuth

\[{\text{B}}{{\text{i}}^{{\text{3 + }}}}{\text{ + 3N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {\text{N}}{{\text{a}}_{\text{3}}}\left[ {{\text{Bi}}{{\left( {{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}} \right)}_{\text{3}}}} \right]\left( {{\text{soluble complex}}} \right){\text{ + 3N}}{{\text{a}}^{\text{ + }}}\]

However, it quickly decomposes into dark \[\left[ {{\text{B}}{{\text{i}}_2}{{\text{S}}_3}} \right]\] ppt.

(v) Reaction with \[{\text{HgC}}{{\text{l}}_{\text{2}}}\]

${\text{N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{  +  HgC}}{{\text{l}}_{\text{2}}} \to {\text{H}}{{\text{g}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ +  2NaCl}} \ \\ $

$  {\text{                                         }} \downarrow {\text{ + H2O}} \ \\ $

$  {\text{HgC}}{{\text{l}}_{\text{2}}}{\text{ }}{\text{. 2HgS }}\xleftarrow{{{\text{HgC}}{{\text{l}}_2}}}{\text{HgS}} \downarrow \left( {{\text{Black}}} \right) \ \\ $ 

(vi) As reducing agent in iodometric titration

(a) \[{{\text{I}}_{\text{2}}}{\text{  +  2}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}^{{\text{2--}}} \to {{\text{2}}^{{\text{I--}}}}{\text{ +  }}{{\text{S}}_{\text{4}}}{{\text{O}}_{\text{6}}}^{{\text{2--}}}\]

(b) \[{\text{2 KMn}}{{\text{O}}_{\text{4}}}{\text{ +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + M}}{{\text{n}}_{\text{2}}}{{\text{O}}_{\text{3}}}\]

(vii) \[{\text{4N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{.5H2O }}\dfrac{{{\text{21}}{{\text{5}}^0}{\text{C}}}}{{{\text{ - H2O(All}})}}{\text{4N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{{\text{3 }}}}\xrightarrow{{{{220}^0}{\text{C}}}}{\text{3N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  N}}{{\text{a}}_{\text{2}}}{{\text{S}}_{\text{5}}}\]

Uses

(i) To eliminate excess chlorine from bleached materials as a ‘antichlor.'

(ii) As a fixer in photography.

(iii) In idometric and idiometric titrations as a reagent.

(B) Hydrogen Sulphide (H2S) 

Preparation:

(i) \[{\text{FeS  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{FeS}}{{\text{O}}_{\text{4}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}\]

It's made with Kipp's apparatus.

(ii) Pure \[{{\text{H}}_{\text{2}}}{\text{S}}\] gas preparation

\[{\text{S}}{{\text{b}}_{\text{2}}}{{\text{S}}_{\text{3}}}\left( {{\text{pure}}} \right){\text{ + 6HCl}}\left( {{\text{pure}}} \right) \to {\text{2SbC}}{{\text{l}}_{\text{3}}}{\text{ + 3}}{{\text{H}}_{\text{2}}}{\text{S}}\]

Properties

(i) Gas that is colourless and smells like rotten eggs

(ii) Moderately soluble in water; however, as temperature rises, solubility diminishes.

(iii) Reducing Agent

As it decomposes developing hydrogen, it acts as a strong reducing agent.

(a) \[{{\text{H}}_{\text{2}}}{\text{S  +  }}{{\text{X}}_{\text{2}}} \to {\text{2HX  +  S}}\]

(b) \[{{\text{H}}_{\text{2}}}{\text{S  + S}}{{\text{O}}_{\text{2}}}\xrightarrow{{{\text{moisture}}}}{{\text{H}}_{\text{2}}}{\text{O  +  S}}\]

(c) \[{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S  +  }}{{\text{O}}_{\text{2}}}\]

(d) \[{\text{2HN}}{{\text{O}}_{\text{3}}} \to {{\text{H}}_{\text{2}}}{\text{O + 2N}}{{\text{O}}_{\text{2}}}{\text{ + }}\left[ {\text{O}} \right]\]

\[\dfrac{{{{\text{H}}_2}{\text{S + }}\left[ {\text{O}} \right]{{\text{H}}_2}{\text{O  +  S}}}}{{{\text{2HN}}{{\text{O}}_3}{\text{  +  }}{{\text{H}}_2}{\text{S}} \to {\text{2}}{{\text{H}}_2}{\text{O  +  N}}{{\text{O}}_2}{\text{ +  S}}}}\]

It reduces 

$  {\text{KMn}}{{\text{O}}_{{\text{4 }}}} \to {\text{M}}{{\text{n}}^{{\text{2 + }}}} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{S}}{{\text{O}}_{\text{2}}}{\text{  }} \ \\ $

 ${{\text{K}}_{\text{2}}}{\text{C}}{{\text{r}}_{\text{2}}}{{\text{O}}_{\text{7}}} \to {\text{C}}{{\text{r}}^{{\text{3 + }}}} \ \\ $ 

(iv) Acidic Nature

According to the reaction below, its aqueous solution behaves as a weak dibasic acid.

\[{{\text{H}}_{\text{2}}}{\text{S}} \rightleftharpoons {\text{HS  +  }}{{\text{H}}^{\text{ + }}} \rightleftharpoons {{\text{S}}^{{\text{2 - }}}}{\text{ +  2}}{{\text{H}}^{\text{ + }}}\]

As a result, it produces two salt series, as shown below.

$  {\text{NaOH  +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {\text{NaHS  +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  {\text{NaOH  +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S  + 2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(v) Formation of Polysulphides

They're made by passing hydrogen sulphide gas through metal hydroxides.

$  {\text{Ca}}{\left( {{\text{OH}}} \right)_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {\text{CaS  +  2}}{{\text{H}}_{\text{2}}}{\text{O ;}} \ \\ $

$  {\text{CaS  +  4}}{{\text{H}}_{\text{2}}}{\text{S}} \to {\text{Ca}}{{\text{S}}_{\text{5}}}{\text{  +  4}}{{\text{H}}_{\text{2}}} \ \\ $

$  {\text{ N}}{{\text{H}}_{\text{4}}}{\text{OH  +  }}{{\text{H}}_{\text{2}}}{\text{S}} \to {\left( {{\text{N}}{{\text{H}}_{\text{4}}}} \right)_{\text{2}}}{\text{S  +  2}}{{\text{H}}_{\text{2}}}{\text{O;}} \ \\ $

$  {\left( {{\text{NH4}}} \right)_{\text{2}}}{\text{S  +  }}{{\text{H}}_{\text{2}}}{\text{S }}\left( {{\text{excess}}} \right) \to {\left( {{\text{N}}{{\text{H}}_{\text{4}}}} \right)_{\text{2}}}{{\text{S}}_{{\text{x + 1 }}}}{\text{ +  x}}{{\text{H}}_{\text{2}}} \ \\ $ 

yellow ammonium sulphide

Tests For \[{{\text{H}}_{\text{2}}}{\text{S}}\]

(i) With sodium nitropruside solution, turns acidified lead acetate paper black.

(ii) With sodium nitropruside solution, gives violet or purple coloration.

Uses

(i) In qualitative analysis, as a laboratory reagent for the detection of basic radicals.

(ii) In the role of a reducing agent.

(c) Sulphur Dioxide 

PREPARATION

(i) \[{\text{S  +  }}{{\text{O}}_{\text{2}}}{\text{ or air}}\xrightarrow{{{\text{Burn}}}}{\text{S}}{{\text{O}}_{\text{2}}}\]

(ii) \[{\text{S  +  2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{conc}}{\text{.}}} \right)\xrightarrow{\Delta }{\text{3S}}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(iii) \[{\text{Cu}}\] or \[{\text{Ag}}\] is heated with conc. \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\]

\[{\text{Cu  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{CuS}}{{\text{O}}_{\text{4}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{2}}}\]

(iv) Metal sulphites react with dil. \[{\text{HCl}}\] to produce dil. \[{\text{HCl}}\].

\[{\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}{\text{  +  2HCl}} \to {\text{2NaCl  +  S}}{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}}\]

Bisulphites, on the other hand, produce \[{\text{S}}{{\text{O}}_{\text{2}}}\] when mixed with dil. \[{\text{HCl}}\]

\[{\text{NaHS}}{{\text{O}}_{\text{3}}}{\text{  +  HCl}} \to {\text{NaCl  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O}}\]

(v) Sulphides are heated more than air.

\[{\text{2 ZnS  +  3}}{{\text{O}}_{{\text{2 }}}} \to {\text{2ZnO  +  2S}}{{\text{O}}_{\text{2}}}\]

(vi) \[{\text{CaS}}{{\text{O}}_{\text{4}}}\left( {{\text{gypsum}}} \right){\text{ +  C}}\dfrac{\Delta }{{{{1000}^0}{\text{C}}}}{\text{ 2CaO  +  S}}{{\text{O}}_{\text{2}}}{\text{ +  C}}{{\text{O}}_2}\]

\[{\text{S}}{{\text{O}}_{\text{2}}}\] is obtained in enormous quantities using this process.

Properties

(i) A colourless gas having a sulphur-burning odour.

(ii) It is denser than air and highly water soluble.

(iii) It neither burns nor aids in the burning of magnesium and potassium, yet they continue to burn in the atmosphere.

$

  {\text{3Mg  +  S}}{{\text{O}}_{\text{2}}} \to {\text{2 MgO  +  MgS }} \ \\ $

$  {\text{4K  +  3S}}{{\text{O}}_{\text{2}}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{K}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}} \ \\ $

(iv) Acidic Nature

Sulphurous acid is formed when acidic oxide dissolves in water.

\[{\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}\]

(v) Addition Reaction

$ {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  C}}{{\text{l}}_{\text{2}}}\xrightarrow{{{\text{Sun light}}}}{\text{S}}{{\text{O}}_{\text{2}}}{\text{C}}{{\text{l}}_{\text{2}}}\left( {{\text{sulphuryl chloride}}} \right) \ \\ $

$  {\text{S}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{O}}_{\text{2}}} \rightleftharpoons {\text{ S}}{{\text{O}}_{\text{3}}}{\text{ }} \ \\ $

$  {\text{Pb}}{{\text{O}}_{\text{2}}}{\text{ +  S}}{{\text{O}}_{\text{2}}} \to {\text{PbS}}{{\text{O}}_{\text{4}}} \ \\ $

(vi) Reducing Nature

$  {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{2}}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}{\text{ }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2H}} \ \\ $ 

Because to the liberation of nascent hydrogen, the character is reducing.

(a) Halogens are reduced to their corresponding halides.

$ {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + 2H}} \ \\ $

$  {\text{2H  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2HCl}} \ \\ $

$  {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{l}}_{\text{2}}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + 2HCl}} \ \\ $ 

(b) Iodates that have been acidified are converted to iodine.

$ {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2H]  \times  5 }} \ \\ $

$  {\text{2KI}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2HI}}{{\text{O}}_{\text{3}}}{\text{ }} \ \\ $

$  {\text{2HI}}{{\text{O}}_{\text{3}}}{\text{ + 10H}} \to {{\text{I}}_{\text{2}}}{\text{ +  6}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{2KI}}{{\text{O}}_{\text{3}}}{\text{ +  5S}}{{\text{O}}_{\text{2}}}{\text{ +  4}}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  4}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  }}{{\text{I}}_{\text{2}}} \ \\ $ 

It reduces acidification 

\[{\text{KMn}}{{\text{O}}_{\text{4}}} \to {\text{M}}{{\text{n}}^{{\text{2 + }}}}{\text{ }}\]

Acidified \[{{\text{K}}_{\text{2}}}{\text{C}}{{\text{r}}_{\text{2}}}{{\text{O}}_{\text{7}}} \to {\text{C}}{{\text{r}}^{{\text{3 + }}}}\]+ (green coloured solution) & Ferric Sulphate ⎯→⎯ Ferrous sulphate

(vii) Oxidising nature

With a strong reducing agent, it acts as an oxidising agent.

(a) \[{\text{2}}{{\text{H}}_{\text{2}}}{\text{S  +  S}}{{\text{O}}_{\text{2}}}\xrightarrow{{{\text{moisture}}}}{\text{2}}{{\text{H}}_{\text{2}}}{\text{O  +  3S}}\]

(b) \[{\text{2SnC}}{{\text{l}}_{\text{2}}}{\text{ +  S}}{{\text{O}}_{\text{2}}}{\text{  +  4HCl}} \to {\text{2SnC}}{{\text{l}}_{\text{4}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  S}}\]

(c) \[{\text{2H}}{{\text{g}}_{\text{2}}}{\text{C}}{{\text{l}}_{\text{2}}}{\text{  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  4HCl}} \to {\text{2HgC}}{{\text{l}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  S}}\]

(d) \[{\text{2CO  +  S}}{{\text{O}}_{\text{2}}} \to {\text{2C}}{{\text{O}}_{{\text{2 }}}}{\text{ +  S}}\]

(e) \[{\text{2Fe  +  S}}{{\text{O}}_{\text{2}}} \to {\text{2FeO  +  FeS}}\]

(viii) Bleaching Action

\[{\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2H}}\]

This is because \[{\text{S}}{{\text{O}}_{\text{2}}}\] is a reducing gas.

Coloured matter + H \[ \rightleftharpoons \]Air oxidation colourless matter.

As a result, bleaching is only a temporary solution.

Uses

(i) Used in the production of \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] and paper made from wood pulp.

(ii) Bleaching agent for delicate materials such as wool, silk, and straw.

(iii) Used in petroleum and sugar refining.

(D) Sulphur Trioxide (\[{\text{S}}{{\text{O}}_{\text{3}}}\]) 

Preparation

(i) \[{\text{6}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  }}{{\text{P}}_{\text{4}}}{{\text{O}}_{{\text{10}}}} \to {\text{6S}}{{\text{O}}_{\text{3}}}{\text{ +  4}}{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}}\]

\[{{\text{P}}_{\text{4}}}{{\text{O}}_{{\text{10}}}}\] is a dehydrating substance.

(ii) \[{\text{F}}{{\text{e}}_{\text{2}}}{\left( {{\text{S}}{{\text{O}}_{\text{4}}}} \right)_{\text{3}}}\xrightarrow{\Delta }{\text{F}}{{\text{e}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ + 3S}}{{\text{O}}_{\text{3}}}\]

(iii) \[{\text{2S}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{O}}_{\text{2}}} \rightleftharpoons {\text{2S}}{{\text{O}}_{\text{3}}}\]

Properties

(i) Acidic Nature

Sulphuric acid is formed when it dissolves in water.

\[{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ }}\]

(ii) \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  S}}{{\text{O}}_{\text{3}}} \to {{\text{H}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{7}}}{\text{ }}\](oleum)

(iii) \[{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  HCl}} \to {\text{S}}{{\text{O}}_{\text{2}}}\left( {{\text{OH}}} \right){\text{Cl }}\](Chlorosulphuric acid)

(iv) Oxidising Nature

(a) \[{\text{2S}}{{\text{O}}_{\text{3}}}{\text{ +  S}}\xrightarrow{{{{100}^0}{\text{C}}}}{\text{3S}}{{\text{O}}_{\text{2}}}\]

(b) \[{\text{5S}}{{\text{O}}_{\text{3}}}{\text{ +  2P}} \to {\text{5S}}{{\text{O}}_{\text{2}}}{\text{ +  P2}}{{\text{O}}_{\text{5}}}{\text{ }}\]

(c) \[{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  PC}}{{\text{l}}_{\text{5}}} \to {\text{POC}}{{\text{l}}_{\text{3}}}{\text{ +  S}}{{\text{O}}_{\text{2}}}{\text{ +  C}}{{\text{l}}_{\text{2}}}{\text{ }}\]

(d) \[{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  2HBr}} \to {{\text{H}}_{\text{2}}}{\text{O  +  B}}{{\text{r}}_{\text{2}}}{\text{  +  S}}{{\text{O}}_{\text{2}}}\]

Uses

(i) It's used to make \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] and oleum.

(ii) It is used as a gas drying agent.

(E) Sulphuric Acid (\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\]) 

PREPARATION

(i) \[{\text{2FeS}}{{\text{O}}_{{\text{4 }}}}{\text{. 7}}{{\text{H}}_{\text{2}}}{\text{O}}\xrightarrow{{{\text{dist}}}}{\text{F}}{{\text{e}}_{\text{2}}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  S}}{{\text{O}}_{\text{2}}}{\text{ + 13}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(ii) Lead Chamber Process (Industrial method)

\[{\text{2S}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{O}}_{\text{2}}}\left( {{\text{air}}} \right){\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O + }}\left[ {{\text{NO}}} \right]\left( {{\text{catalyst}}} \right) \to {\text{2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + }}\left[ {{\text{NO}}} \right]\left( {{\text{catalyst}}} \right)\]

The acid obtained is known as brown oil of vitriol and is 80 percent pure.

(iii) Contact process (Industrial method)

\[{{\text{O}}_{\text{2}}}{\text{ + 2S}}{{\text{O}}_{\text{2}}} \rightleftharpoons {\text{2S}}{{\text{O}}_{\text{3}}}\]

Platinum, ferric oxide, and vanadium pentoxide are the most often used catalysts. \[{{\text{V}}_2}{{\text{O}}_{\text{5}}}\] is favoured since it is less expensive and does not contain contaminants.

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{58\% }}} \right){\text{ +  S}}{{\text{O}}_{\text{3}}} \to {{\text{H}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{7}}}\left( {{\text{oleum}}} \right)\]

By diluting oleum with water, you can get sulphuric acid in any concentration you want.

\[{{\text{H}}_{\text{2}}}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{7}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\]

Sulphur dioxide oxidation is reversible and exothermic. The favourable circumstances for a higher output of sulphur trioxide, according to the Le-chatelier principle, are.

(a) An surplus of air—the molecular proportions of \[{\text{S}}{{\text{O}}_{\text{2}}}\] and oxygen are 2:3

(b) Low temperature-optimal temperature \[{450^0}{\text{C}}\]

(c) Higher pressure-one atmosphere

Properties

(i)It's a colourless syrupy liquid (it's H-bonded) 

(ii) It's highly corrosive and fumes heavily in damp air.

(iii) Thermal decomposition

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \rightleftharpoons {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{3}}}\]

(iv) Acidic Nature

It's a dibasic acid with a high ionisation potential.

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \rightleftharpoons {{\text{H}}^{\text{ + }}}{\text{ +  HS}}{{\text{O}}_{\text{4}}}^{\text{--}}{\text{2}}{{\text{H}}^{\text{ + }}}{\text{ + S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}\]

(a) Produces two salt series

$  {\text{NaOH + }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{NaHS}}{{\text{O}}_{\text{4}}}\left( {{\text{sodium bisulphate}}} \right){\text{ + }}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{NaHS}}{{\text{O}}_{\text{4}}}{\text{ + NaOH}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{sodium sulphate}}} \right){\text{ + }}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(b) Carbonates and bicarbonates are decomposed into carbon dioxide.

$  {\text{N}}{{\text{a}}_{\text{2}}}{\text{C}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{ N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{O}}_{\text{2}}}{\text{ }} \ \\ $

$  {\text{NaHC}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{NaHS}}{{\text{O}}_{\text{4}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{O}}_{\text{2}}} \ \\ $ 

(c) Displaces the metal salts of more volatile acids.

$ {\text{2NaCl  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + HCl}} \ \\ $

$  {\text{ 2NaN}}{{\text{O}}_{\text{3}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  2HN}}{{\text{O}}_{\text{3}}}{\text{ }} \ \\ $

$  {\text{Ca}}{{\text{F}}_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{CaS}}{{\text{O}}_{\text{4}}}{\text{  +  2HF}} \ \\ $ 

(v) Oxidising Nature

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] is a powerful oxidising agent.

\[{\text{2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2}}{{\text{e}}^{\text{--}}} \to {\text{S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O +  S}}{{\text{O}}_{\text{2}}}\]

or

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{{\text{4 }}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  }}\left[ {\text{O}} \right]\]

(a) Non-metals (such as carbon and sulphur) are oxidised to their oxides.

$ {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{{\text{2 }}}}{\text{ + }}\left[ {\text{O}} \right] \times 2 \ \\ $

$  {\text{C + 2}}\left[ {\text{O}} \right] \to {\text{C}}{{\text{O}}_{\text{2}}} \ \\ $

$  {\text{C  +  2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{C}}{{\text{O}}_{\text{2}}}{\text{ +  2S}}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(b) Metals (copper, silver, mercury, and so on) are oxidised to oxides, which then react with acid to generate sulphates.

$  {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  }}\left[ {\text{O}} \right]{\text{ }} \ \\ $

$  {\text{Cu  +  }}\left[ {\text{O}} \right]{\text{ }} \to {\text{CuO }} \ \\ $

$  {\text{CuO  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{CuS}}{{\text{O}}_{\text{4}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{Cu  +  2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{CuSO4  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(c) Iodine is liberated from KI

$  {\text{2KI  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ }} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{{\text{4 }}}}{\text{ +  2HI }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {{\text{H}}_{\text{2}}}{\text{O  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  }}\left[ {\text{O}} \right]{\text{ }} \ \\ $

$  {\text{2HI  + }}\left[ {\text{O}} \right] \to {{\text{I}}_{\text{2}}}{\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  {\text{2 KI  +  2}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  S}}{{\text{O}}_{\text{2}}}{\text{  +  }}{{\text{I}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

In this case, \[{\text{HI}}\] is oxidised to \[{{\text{I}}_{\text{2}}}\]. Bromine is also released from \[{\text{KBr}}\].

(d) \[{{\text{C}}_{{\text{10}}}}{{\text{H}}_{\text{8}}}\left( {{\text{naphthalene}}} \right){\text{ +  9}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\dfrac{{{\text{Hg as}}}}{{{\text{Catalyst}}}}{\text{ }}{{\text{C}}_{\text{8}}}{{\text{H}}_{\text{6}}}{{\text{O}}_{\text{4}}}\left( {{\text{phthalic acid}}} \right){\text{  + 10}}{{\text{H}}_{\text{2}}}{\text{O  +  9S}}{{\text{O}}_{\text{2}}}{\text{ + 2C}}{{\text{O}}_{\text{2}}}\]

(vi) Dehydrating agent

Because of its affinity for water, sulphuric acid is a powerful dehydrating agent.

(a) \[{{\text{C}}_{{\text{12}}}}{{\text{H}}_{{\text{22}}}}{{\text{O}}_{{\text{11 }}}}\left( {{\text{cane sugar}}} \right){\text{ }}\dfrac{{{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}}}{{{\text{ - 11}}{{\text{H}}_{\text{2}}}{\text{O}}}}{\text{12C}}\]

(vii) Miscellaneous reactions

(a) Sulphonation of aromatic compounds

(b) Reaction with \[{\text{PC}}{{\text{l}}_5}\]

$ \left( {\text{c}} \right){\text{ }}{{\text{K}}_{\text{4}}}\left[ {{\text{Fe}}{{\left( {{\text{CN}}} \right)}_{\text{6}}}{\text{ }}} \right]{\text{  +  6}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  6}}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{2}}{{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  FeS}}{{\text{O}}_{\text{4}}}{\text{  +  3}}\left( {{\text{N}}{{\text{H}}_{\text{4}}}{\text{ }}} \right){\text{2S}}{{\text{O}}_{\text{4}}}{\text{  +  6CO }} \ \\ $

$  \left( {\text{d}} \right){\text{ 3KCl}}{{\text{O}}_{{\text{3 }}}}{\text{ +  3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\xrightarrow{\Delta }{\text{3KHS}}{{\text{O}}_{\text{4}}}{\text{ +  HCl}}{{\text{O}}_{\text{4}}}{\text{  +  2Cl}}{{\text{O}}_{{\text{2 }}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  \left( {\text{e}} \right){\text{ }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  }}{{\text{P}}_{\text{2}}}{{\text{O}}_{\text{5}}} \to {\text{2HP}}{{\text{O}}_{\text{3}}}{\text{  +  S}}{{\text{O}}_{\text{3}}} \ \\ $ 

Uses

(i) Fertilizers such as ammonium sulphate and super phosphate of lime are made.

(ii) As a vital reagent in the laboratory.

(iii) In the form of storage batteries.

(iv) In the industries of leather, textiles, paper, and dyeing.

Group 17 Elements : 

The Halogen Family

Group 17 includes fluorine, chlorine, bromine, iodine, and astatine. The halogens are the collective name for these substances (Greek halo means salt and genes born i.e., salt producers). The halogens are non-metallic elements with a high reactivity.

Electronic Configuration

These elements all have seven electrons in their outermost shells (\[n{s^2}n{p^5}\]), one less than the following noble gas.

Atomic and Ionic Radii

Because of their maximal effective nuclear charge, halogens have the shortest atomic radii in their respective eras. The number of quantum shells increases as the number of atomic and ionic radii increases from fluorine to iodine.

Ionisation Enthalpy

They have a low proclivity for losing electrons. As a result, their ionisation enthalpy is extremely high. The group's ionisation enthalpy decreases as atomic size increases.

Electron Gain Enthalpy

In the same period, negative electrons obtain the most enthalpy. This is because the atoms of these elements only have one less electron than stable noble gas structures. The electron gain enthalpy of the group's elements decreases as the group progresses. Fluorine, on the other hand, has a lower negative electron gain enthalpy than chlorine. It is owing to the fluorine atom's tiny size. As a result, substantial interelectronic repulsions exist in fluorine's relatively tiny 2p orbitals, and the incoming electron receives little attraction.

Electronegativity

Their electronegativity is high. The electronegativity drops as you progress through the group. Fluorine is the periodic table's most electronegative element.

Physical Properties

Iodine is a solid while fluorine and chlorine are gases. Bromine is a liquid and fluorine is a gas. With increasing atomic number, their melting and boiling points rise. Halogen is all coloured. This is due to the absorption of visible light, which causes the outer electrons to be excited to a higher energy level. They show distinct colours by absorbing varying quanta of radiation.

 for example, \[{{\text{F}}_{\text{2}}}\], is yellow, \[{\text{C}}{{\text{l}}_{\text{2}}}\] is greenish yellow, \[{\text{B}}{{\text{r}}_{\text{2}}}\] is red, and \[{{\text{I}}_{\text{2}}}\] is violet. Water reacts with fluorine and chlorine.

In water, bromine and iodine are only slightly soluble. However, organic solvents such as chloroform, carbon tetrachloride, carbon disulphide, and hydrocarbons are soluble in them, resulting in coloured solutions. Except for the fact that \[{{\text{F}}_{\text{2}}}\] has a lower dissociation enthalpy than \[{\text{C}}{{\text{l}}_{\text{2}}}\], \[{\text{X - X}}\] bond dissocitation enthalpies from chlorine onwards follow the predicted pattern: \[{\text{Cl -- Cl  >  Br -- Br  >  I -- I}}\]. The higher electron-electron repulsion among the lone pairs in the \[{{\text{F}}_{\text{2}}}\] molecule, where they are significantly closer to one other than in the \[{\text{C}}{{\text{l}}_{\text{2}}}\] molecule, is one cause for this abnormality.

Chemical Properties 

Oxidation States and Trends in Chemical Reactivity

The oxidation state of all halogens is –1. Chlorine, bromine, and iodine, on the other hand, have oxidation states of + 1, + 3, + 5, and + 7. Chlorine, bromine, and iodine have greater oxidation states when they combine with the tiny and extremely electronegative fluorine and oxygen atoms, as in interhalogens, oxides, and oxoacids.

Because the fluorine atom's valence shell lacks d orbitals, it can't expand its octet. It has only one oxidation state since it is the most electronegative.

The halogens are all extremely reactive. They form halides when they react with metals and non-metals. The halogens' reactivity reduces as they progress through the group.

The high oxidising tendency of halogens is due to their easy uptake of an electron. \[{{\text{F}}_2}\] is the most powerful halogen oxidizer, oxidising other halide ions in solution and in solid form. The usual electrode potentials show that the halogen's oxidising activity in aqueous solution decreases as the group progresses. Water is oxidised by fluorine to produce oxygen, but chlorine and bromine react with water to produce hydrohalic and hypohalous acids. Iodine interactions with water are not spontaneous. In an acidic solution, I– can be oxidised by oxygen, which is the exact opposite of what happens with fluorine.

$  {\text{\;2}}{{\text{F}}_{\text{2}}}{\text{ }}\left( {\text{g}} \right){\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\left( \ell  \right) \to {\text{ 4}}{{\text{H}}^{\text{ + }}}\left( {{\text{aq}}} \right){\text{  +  4F-- }}\left( {{\text{aq}}} \right){\text{  +  }}{{\text{O}}_{\text{2}}}{\text{ }}\left( {\text{g}} \right) \ \\ $

  ${{\text{X}}_{\text{2}}}\left( {\text{g}} \right){\text{  +  }}{{\text{H}}_{\text{2}}}{\text{O }}\left( \ell  \right) \to {\text{HX}}\left( {{\text{aq}}} \right){\text{  +  HOX }}\left( {{\text{aq}}} \right){\text{ }} \ \\ $ \left( {{\text{where X  =  Cl or Br}}} \right){\text{ }} \ \\ $

$  {\text{4}}{{\text{I}}^{\text{--}}}\left( {{\text{aq}}} \right){\text{  +  4}}{{\text{H}}^{\text{ + }}}\left( {{\text{aq}}} \right){\text{  +  }}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right) \to {\text{2}}{{\text{I}}_{\text{2}}}\left( {\text{s}} \right){\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\left( \ell  \right) \ \\ $ 

Standard Reduction Potential (SRP)

$  {{\text{X}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}} \to {\text{2}}{{\text{X}}^{\text{--}}} \ \\ $

$  {{\text{F}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}} \to {\text{2}}{{\text{F}}^{\text{--}}}{\text{   \varepsilon ^\circ   =   +  2}}{\text{.87 V}} \ \\ $

$  {\text{B}}{{\text{r}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{{\text{ --}}}} \to {\text{2Br}}{{\text{ }}^{\text{--}}}{\text{     \varepsilon ^\circ   =   +  1}}{\text{.09 V }} \ \\ $

$  {\text{C}}{{\text{l}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}} \to {\text{2C}}{{\text{l}}^{\text{--}}}{\text{       \varepsilon ^\circ   =   +  1}}{\text{.36 V}} \ \\ $

$  {{\text{I}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}} \to {\text{2}}{{\text{e}}^{\text{--}}}{\text{     \varepsilon ^\circ   =   +  0}}{\text{.54 V}} \ \\ $ 

The (algebraically) oxidising agent becomes more potent as the SRP increases in value. As a result, the order of oxidising power is \[{{\text{F}}_{\text{2}}}{\text{  >  C}}{{\text{l}}_{\text{2}}}{\text{  >  B}}{{\text{r}}_{\text{2}}}{\text{  >  2}}\]SRP is a strogenst oxidising agent since it is the highest for \[{{\text{F}}_{\text{2}}}\] (among all elements of P.T.).

\[{{\text{F}}_{\text{2}}}\] is a more potent oxidizer than \[{{\text{O}}_{\text{3}}}\].

[Despite the fact that \[{{\text{O}}_{\text{3}}}\] has three ‘O's]

Note: The E.A. and.E. values refer to atoms in the gaseous medium, where redox processes occur.

As a result, qualities in the gas phase cannot be reflected in the solution phase in the same way.

Because electrode potential values are experimental (based on the correct situation), they would be the monitoring parameter in the solution phase.

Hydration Energy of \[{{\text{X}}^ - }\]

The higher the hydration energy, the smaller the ion.

Anomalous Behaviour of Fluorine

Fluorine's unusual behaviour is caused by its tiny size, strong electronegativity, low F-F bond dissociation enthalpy, and lack of d orbitals in the valence shell. Fluorine's reactions are mostly exothermic (owing to the tiny and strong bonds it forms with other elements). It only produces one oxoacid, whereas other halogens produce many oxoacids.

Due to strong hydrogen bonding, hydrogen fluoride is liquid (b.p. 293 K). Other hydrogen halides exist in the form of gases.

(i) Reactivity Towards Hydrogen

They all produce hydrogen halides when they react with hydrogen, although their affinity for hydrogen decreases as they progress from fluorine to iodine. Hydrohalic acids are formed when they dissolve in water.  The stability of these halides decreases when the bond (H–X) dissociation enthalpy decreases in the following order: \[{\text{HF < HCl < HBr < HI}}\] 

(ii) Reactivity towards oxygen

When halogens react with oxygen, they produce a variety of oxides, the majority of which are unstable. \[{\text{O}}{{\text{F}}_{\text{2}}}\] and \[{{\text{O}}_{\text{2}}}{{\text{F}}_{\text{2}}}\]are the two oxides formed by fluorine. At 298 K, however, only \[{\text{O}}{{\text{F}}_{\text{2}}}\] is thermally stable. Because flurorine has a higher electronegativity than oxygen, these oxides are effectively oxygen fluorides. Both are fluorinating agents with a high fluorinating power. The reactions in which \[{{\text{O}}_{\text{2}}}{{\text{F}}_{\text{2}}}\] oxidises plutonium to \[{\text{Pu}}{{\text{F}}_{\text{6}}}\] are utilised to remove plutonium as \[{\text{Pu}}{{\text{F}}_{\text{6}}}\] from spent nuclear fuel.

Chlorine, bromine, and iodine are halogens that produce oxides with oxidation states ranging from 1 to 7. The order of stability oxides generated by halogens is generally decreasing due to a combination of kinetic and thermodynamic variables. \[{\text{C}}{{\text{l}}_{\text{2}}}{\text{O, Cl}}{{\text{O}}_{\text{2}}}{\text{, C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{6}}}{\text{, and C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{7}}}\] are highly reactive oxidising chemicals that tend to burst. \[{\text{Cl}}{{\text{O}}_{\text{2}}}\] is used in water treatment and as a bleaching agent for paper pulp and textiles.

The bromine oxides, \[{\text{B}}{{\text{r}}_{\text{2}}}{\text{O, Br}}{{\text{O}}_{\text{2}}}{\text{, and Br}}{{\text{O}}_{\text{3}}}\], are the least stable of the halogen oxides and can only exist at very low temperatures. They are extremely potent oxidizers.

Iodine oxides, such as \[{{\text{I}}_{\text{2}}}{{\text{O}}_{\text{4}}}{\text{, }}{{\text{I}}_{\text{2}}}{{\text{O}}_{\text{5}}}{\text{, and }}{{\text{I}}_{\text{2}}}{{\text{O}}_{\text{7}}}\], are insoluble solids that disintegrate when heated. \[{{\text{I}}_{\text{2}}}{{\text{O}}_{\text{5}}}\] is a powerful oxidizer that is used to calculate carbon monoxide levels.

(iii) Reactivity Towards Metals

Metal halides are formed when halogen reacts with metals. Bromine, for example, interacts with magnesium to produce magnesium bromide.

(iv) Reactivity of Halogen Towards Other Halogens

Halogens mix to generate interhalogens of the types AB, AB3, AB5, and AB7, where A is a larger size halogen and B is a smaller size halogen.

1. Fluorine (\[{{\text{F}}_{\text{2}}}\]) 

Preparation:

(i) Electrolytic method: 

Electrolyte: Molten \[{\text{KH}}{{\text{F}}_{\text{2}}}\] (1 part) + \[{\text{HF}}\] (5 part)

Anode: Carbon 

Cathode: Steel 

Vessel: Monel metal

On Electrolysis

Cathode: \[{\text{2}}{{\text{H}}^{\text{ + }}}{\text{ +  2}}{{\text{e}}^{\text{--}}} \to {{\text{H}}_{\text{2}}}\left( {\text{g}} \right)\]

Anode: \[{\text{2}}{{\text{F}}^{{\text{-- }}}} \to {{\text{F}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}}\]

The resulting \[{{\text{F}}_{\text{2}}}\] gas must be free of HF, which is more corrosive than fluorine.

To remove HF from flourine, the gas is routed via \[{\text{NaF}}\], which absorbs the HF.

Because \[{{\text{F}}_{\text{2}}}\] rapidly interacts with graphite to generate a polymeric substance known as graphite fluoride, the anode of carbon should be devoid of graphide.

Otherwise, fluorine will react with water if there is any moisture in the vessel.

$  {\text{3}}{{\text{H}}_{\text{2}}}{\text{O  +  3}}{{\text{F}}_{\text{2}}} \to {\text{6HF  +  }}{{\text{O}}_{\text{3}}}{\text{ }} \ \\ $

$  {\text{2}}{{\text{F}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{4HF  +  }}{{\text{O}}_{\text{2}}} \ \\ $ 

Fluorine cannot be produced by electrolysis of aqueous \[{\text{NaF}}\] or KF solutions. This is because when aqueous KF solution is electrolyzed, two oxidation reactions compete at the anode.

$  {{\text{H}}_{\text{2}}}{\text{O}} \to {\text{1/2}}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{H}}^{\text{ + }}}{\text{  +  2}}{{\text{e}}^{\text{--}}}{\text{     SOP  =  -- 1}}{\text{.23 v}} \ \\ $

$  {{\text{F}}^{\text{--}}}{\text{ }} \to {\text{1/2}}{{\text{F}}_{\text{2}}}{\text{  +  }}{{\text{e}}^{\text{--}}}{\text{              SOP  =  -- 2}}{\text{.87v}} \ \\ $ 

As a rule, that substance will oxidise if its SOP is higher, hence water will oxidise at the anode rather than \[{{\text{F}}^{\text{--}}}\]

(ii) Chemical Method

From \[{{\text{K}}_{\text{2}}}\left[ {{\text{Mn}}{{\text{F}}_{\text{6}}}} \right]\]- potassium hexafluoromanganate (IV)

\[{{\text{K}}_{\text{2}}}\left[ {{\text{Mn}}{{\text{F}}_{\text{6}}}} \right]{\text{ +  2Sb}}{{\text{F}}_{\text{5}}} \to {\text{2K}}\left[ {{\text{Sb}}{{\text{F}}_{\text{6}}}} \right]{\text{ + Mn}}{{\text{F}}_{\text{3}}}{\text{ + }}\dfrac{{\text{1}}}{2}{\text{ }}{{\text{F}}_{\text{2}}}\]

The stronger Lewis acid, \[{\text{Sb}}{{\text{F}}_{\text{5}}}\], displaces the weaker Lewis acid, \[{\text{Mn}}{{\text{F}}_{\text{4}}}\], from its salt in this reaction. \[{\text{Mn}}{{\text{F}}_{\text{4}}}\] is a highly unstable compound that easily decomposes into \[{\text{Mn}}{{\text{F}}_{\text{3}}}\] and fluorine.

Properties

(i) Diatomic, pale green-yellow gas with an almost colourless appearance. It has a higher density than air. At – 188°C, it condenses to yellow liquid, and at – 2230°C, it condenses to yellow solute. It is exceedingly poisonous and has punged oddur.

(ii) Oxidizing properties: It is the most potent oxidizer.

$  {{\text{F}}_{\text{2}}}{\text{ +  2NaX}} \to {\text{2NaF  +  }}{{\text{X}}_{{\text{2 }}}} \ \\ $

$  \left( {{\text{X  =  Cl, Br, n}}} \right) \ \\ $ 

(a) It can convert all other halide ions to halogen molecules.

(b) It can oxidise \[{\text{Cl}}{{\text{O}}_3}^ - \] into \[{\text{Cl}}{{\text{O}}_4}^ - \] and \[{\text{I}}{{\text{O}}_3}^ - \] to \[{\text{I}}{{\text{O}}_4}^ - \]

\[{{\text{F}}_{\text{2}}}{\text{  +  Cl}}{{\text{O}}_{\text{3}}}^{\text{--}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{2}}{{\text{F}}^{\text{--}}}{\text{ +  Cl}}{{\text{O}}_{\text{4}}}^{{\text{-- }}}{\text{ +  2}}{{\text{H}}^{\text{ + }}}\]

(c) It can oxidise\[{\text{HS}}{{\text{O}}_{\text{4}}}^{{\text{-- }}}{\text{into }}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{8}}}^{{\text{2--}}}\]

\[{\text{2HS}}{{\text{O}}_{\text{4}}}^{{\text{-- }}}{\text{ + }}{{\text{F}}_2} \to {\text{2}}{{\text{F}}^ - }{\text{ + }}{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{8}}}^{{\text{2--}}} + 2{{\text{H}}^ + }\]

In the persulphate ions, some of the \[{{\text{O}}_{\text{2}}}^{{\text{-- }}}\] is converted to \[{{\text{O}}^{{\text{-- }}}}\] (bearing the \[{{\text{O}}^{{\text{-- }}}}\]). As a result, oxygen gets oxidised.

(iii) Reaction with NaOH solution: Oxygen difluoride is formed by dilute alkali, whereas oxygen difluoride is formed by concentrated alkali.

$  {\text{2}}{{\text{F}}_{\text{2}}}{\text{ + 2NaOH}}\left( {{\text{dil}}} \right) \to {\text{O}}{{\text{F}}_{\text{2}}}\left( {\text{g}} \right){\text{  +  2NaF  +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  {\text{2}}{{\text{F}}_{\text{2}}}{\text{ + 4NaOH}}\left( {{\text{conc}}{\text{.}}} \right) \to {{\text{O}}_{\text{2}}}\left( {\text{g}} \right){\text{  +  4NaF  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(iv) Reaction with \[{\text{N}}{{\text{H}}_{\text{3}}}\]: (Distinction from other halogens)

\[{\text{2N}}{{\text{H}}_{\text{3}}}{\text{ +  3}}{{\text{F}}_{\text{2}}} \to {{\text{N}}_{\text{2}}}{\text{ +  6HF}}\]

Other halogens combine with conc. \[{\text{N}}{{\text{H}}_{\text{3}}}\] (liquor ammonia) to generate explosive \[{\text{N}}{{\text{X}}_{\text{3}}}\]

(v) Reaction with H2S:

\[{{\text{H}}_{\text{2}}}{\text{S  +  }}{{\text{F}}_{\text{2}}} \to {\text{S}}{{\text{F}}_{\text{6}}}{\text{ + 2HF}}\]

\[{{\text{H}}_{\text{2}}}{\text{S}}\] burns

(vi) Reaction with SiO2:

It attacks glass at a temperature of around \[{100^0}{\text{C}}\]

\[{\text{Si}}{{\text{O}}_{\text{2}}}\left( {\text{s}} \right){\text{ + 2}}{{\text{F}}_{\text{2}}}\left( {\text{g}} \right) \to {\text{Si}}{{\text{F}}_{\text{4}}}\left( {\text{g}} \right){\text{ + }}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right)\]

With dry \[{{\text{F}}_{\text{2}}}\], the reaction is gradual.

(vii) Reaction with \[{{\text{H}}_{\text{2}}}{\text{O}}\]:

\[{\text{2}}{{\text{H}}_{\text{2}}}{\text{O  + 2}}{{\text{F}}_{\text{2}}} \to {\text{4HF  +  }}{{\text{O}}_{\text{2}}}\]

Occasionally, a small amount of \[{{\text{O}}_{\text{3}}}\] is formed.

\[{\text{3}}{{\text{H}}_{\text{2}}}{\text{O  + 3}}{{\text{F}}_{\text{2}}} \to {\text{6HF  +  }}{{\text{O}}_{\text{3}}}{\text{ }}\]

(viii) Reaction with \[{\text{Xe}}\]

(ix) Reaction with \[{{\text{H}}_2}\]

\[{{\text{H}}_{\text{2}}}{\text{ +  }}{{\text{F}}_{\text{2}}} \to {\text{2HF or }}{{\text{H}}_{\text{2}}}{\text{ }}{{\text{F}}_{\text{2}}}\]

Even in the dark, this reflex occurs.

(x) Reaction with \[{\text{S}}{{\text{O}}_{\text{3}}}\]

\[{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{F}}_{\text{2}}}\dfrac{{{{180}^0}{\text{C}}}}{{{\text{AgF}}}}{\text{ FS}}{{\text{O}}_{\text{2}}}{\text{OOS}}{{\text{O}}_{\text{2}}}{\text{F}}\]

(xi) Reaction with O2

\[{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{F}}_{\text{2}}}\dfrac{{{\text{only in presence of}}}}{{{\text{silent electric discharge}}}}{\text{ >  }}{{\text{O}}_{\text{2}}}{{\text{F}}_{\text{2}}}\]

(xii) Reaction with Metals and Non–Metals

It is compatible with most metals. In the presence of \[{{\text{F}}_{\text{2}}}\], almost all non-metals, except for \[{{\text{O}}_{\text{2}}}\] and \[{{\text{N}}_{\text{2}}}\], spontaneously ignite.

$ {\text{2Ag  +  }}{{\text{F}}_{\text{2}}} \to {\text{2AgF ; 2Al  +  3}}{{\text{F}}_{\text{2}}} \to {\text{2Al}}{{\text{F}}_{\text{3}}}{\text{ }} \ \\ $  ${\text{C  +  2 }}{{\text{F}}_{\text{2}}} \to {\text{C}}{{\text{F}}_{\text{4}}}{\text{ ; Si  +  2}}{{\text{F}}_{\text{2}}} \to {\text{Si}}{{\text{F}}_{\text{4}}} \ \\ $ 

2. Chlorine (Cl2)

PREPARATION

(i) Common method (\[{\text{C}}{{\text{l}}_{\text{2}}}\], \[{\text{B}}{{\text{r}}_{\text{2}}}\] , \[{{\text{I}}_{\text{2}}}\] )

$ {\text{2NaX  +  3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{conc}}{\text{.}}} \right){\text{ +  Mn}}{{\text{O}}_{\text{2}}}\left( {{\text{oxidising agent}}} \right)\xrightarrow{\Delta }{{\text{X}}_{\text{2}}}{\text{ +  MnS}}{{\text{O}}_{\text{4}}}{\text{ +  2NaHS}}{{\text{O}}_{\text{4}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  {\text{4}}{{\text{H}}^{\text{ + }}}{\text{  +  Mn}}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{X}}^{\text{--}}} \to {{\text{X}}_{\text{2}}}{\text{  +  M}}{{\text{n}}^{{\text{ + 2}}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(ii) Only for \[{\text{C}}{{\text{l}}_{\text{2}}}\]

$  {\text{NaCl  +  HN}}{{\text{O}}_{\text{3}}} \to {\text{NaN}}{{\text{O}}_{\text{3}}}{\text{ +  HCl  \times  3 }} \ \\ $

$  {\text{HN}}{{\text{O}}_{\text{3}}}{\text{ +  3HCl}} \to {\text{NOCl  +  C}}{{\text{l}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{3NaCl  +  4HN}}{{\text{O}}_{\text{3}}} \to {\text{3NaN}}{{\text{O}}_{\text{3}}}{\text{ +  NOCl}}\left( {{\text{nitrosyl chloride}}} \right){\text{ +  C}}{{\text{l}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$  {\text{2NOCl  +  }}{{\text{O}}_{\text{2}}} \to {\text{2N}}{{\text{O}}_{\text{2}}}{\text{  +  C}}{{\text{l}}_{\text{2}}}{\text{ }} \ \\ $

$  {\text{N}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{HN}}{{\text{O}}_{\text{3}}} \ \\ $ 

(a) When\[{\text{C}}{{\text{l}}_{\text{2}}}\] is used to chlorinate hydrocarbons, \[{\text{HCl}}\] is produced as a byproduct. Copper powder combined with rare earth chlorides is used to catalytically oxidise \[{\text{HCl}}\] into \[{{\text{H}}_{\text{2}}}{\text{O}}\] and \[{\text{C}}{{\text{l}}_{\text{2}}}\]. 

(b) \[{\text{4 HCl  +  }}{{\text{O}}_{\text{2}}}{\text{ }}\dfrac{{{\text{Cu powder + rare earth}}}}{{{\text{Chloride}}}}{\text{2}}{{\text{H}}_{\text{2}}}{\text{O  +  2C}}{{\text{l}}_{\text{2}}}\]

(c)

(d) \[{\text{2KMn}}{{\text{O}}_{\text{4}}}{\text{ + 16 HCl}} \to {\text{2KCl +  2MnC}}{{\text{l}}_{\text{2}}}{\text{ +  5C}}{{\text{l}}_{\text{2}}}{\text{ + 8}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(e)  \[{\text{Pb}}{{\text{O}}_{\text{2}}}{\text{ +  4HCl}} \to {\text{PbC}}{{\text{l}}_{\text{2}}}{\text{,  +  C}}{{\text{l}}_{\text{2}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(iii) Manufacture of chlorine

(a) When \[{\text{C}}{{\text{l}}_{\text{2}}}\] is used to chlorinate hydrocarbons, \[{\text{HCl}}\] is produced as a byproduct. Copper powder combined with rare earth chlorides is used to catalytically oxidise \[{\text{HCl}}\] into \[{{\text{H}}_{\text{2}}}{\text{O}}\] and \[{\text{C}}{{\text{l}}_{\text{2}}}\].

\[{\text{4HCl  +  }}{{\text{O}}_{\text{2}}}\xrightarrow{{{\text{CuC}}{{\text{l}}_2}}}{\text{2C}}{{\text{l}}_{\text{2}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(b) Electrolytic process: The electrolysis of brine produces chlorine (concentrated NaCl solution). At the anode, chlorine is released. Many chemical industries also produce it as a byproduct.

$ {\text{NaX}}\left( {{\text{aq}}} \right) \to {\text{N}}{{\text{a}}^{{\text{ +  }}}}\left( {{\text{aq}}} \right){\text{  +  }}{{\text{X}}^{{\text{-- }}}}\left( {{\text{aq}}} \right){\text{ }} \ \\ $

$  {\text{Anode : 2}}{{\text{X}}^{\text{--}}} \to {{\text{X}}_{\text{2}}}{\text{ +  2}}{{\text{e}}^{\text{--}}} \ \\ $ 

Properties

(i) It has a harsh and suffocating odour and is a greenish–yellow gas. It weighs around 2–5 times as much as air. It can be liquefied into a greenish–yellow liquid that boils at 239 degrees Fahrenheit. It is water soluble.

(ii) It forms a hydrate with water with the formula \[{\text{C}}{{\text{l}}_2}.8{{\text{H}}_2}{\text{O}}\], which is a clathrate chemical, at low temperatures.

(iii) \[{{\text{H}}_{\text{2}}}{\text{ +  C}}{{\text{l}}_{\text{2}}}\xrightarrow[{{\text{a zero order reaction}}}]{}{\text{2HCl }}\left( {\text{g}} \right){\text{ , }}{{\text{H}}_{\text{2}}}{\text{ +  B}}{{\text{r}}_{\text{2}}} \to {\text{ 2HBr }}\]is not a zero- order reaction.

(iv) Reaction with \[{\text{N}}{{\text{H}}_{\text{3}}}\] (common for \[{\text{C}}{{\text{l}}_{\text{2}}}\] & \[{\text{B}}{{\text{r}}_{\text{2}}}\] )

$  \left( {\text{a}} \right){\text{ 8NH3 + 3C}}{{\text{l}}_{\text{2}}} \to {{\text{N}}_{\text{2}}}{\text{ + 6N}}{{\text{H}}_{\text{4}}}{\text{Cl }} \ \\ $

$  \left( {\text{b}} \right){\text{ N}}{{\text{H}}_{\text{3}}}{\text{ + 3C}}{{\text{l}}_{\text{2}}} \to {\text{NC}}{{\text{l}}_{\text{3}}}{\text{ + 3HCl}} \ \\ $ 

(v) Reaction with alkali metal halides (KX)

$  {\text{2 KBr  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2KCl  +  B}}{{\text{r}}_{\text{2}}}{\text{ }} \ \\ $

$  {\text{2KI  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2KCl  +  }}{{\text{I}}_{\text{2}}} \ \\ $ 

\[{\text{C}}{{\text{l}}_{\text{2}}}\] can oxidise both \[{\text{B}}{{\text{r}}^ - }\] and \[{{\text{I}}^ - }\], but \[{\text{B}}{{\text{r}}_{\text{2}}}\] can only oxidise \[{{\text{I}}^ - }\].

Because \[{{\text{F}}_{\text{2}}}\] reacts with water, it is not employed in aqueous reactions.

(vi) Oxidising & bleaching properties: Because \[{\text{HCl}}\] and \[{\text{HOCl}}\] develop as chlorine water sits, it loses its yellow colour. When hypochlorous acid (\[{\text{HOCl}}\]) is created, nascent oxygen is produced, which is responsible for chlorine's oxidising and bleaching effects.

(i) Ferrous to ferric, sulphite to sulphate, sulphur dioxide to sulphuric acid, and iodine to iodic acid are all oxidised.

$ {\text{2 FeS}}{{\text{O}}_{\text{4}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{  +  C}}{{\text{l}}_{\text{2}}} \to {\text{F}}{{\text{e}}_{\text{2}}}{\left( {{\text{S}}{{\text{O}}_{\text{4}}}} \right)_{\text{3}}}{\text{ + 2HCl }} \ \\ $

$  {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{3}}}{\text{ +  C}}{{\text{l}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2HCl }} \ \\ $

$  {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O  + C}}{{\text{l}}_{\text{2}}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2HCl }} \ \\ $  ${\text{I2  +  6}}{{\text{H}}_{\text{2}}}{\text{O + 5C}}{{\text{l}}_{\text{2}}} \to {\text{2HI}}{{\text{O}}_{\text{3}}}{\text{ + 10HCl}} \ \\ $ 

(ii) It is a strong bleaching agent; oxidation is responsible for the bleaching activity.

\[{\text{C}}{{\text{l}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O}} \to {\text{2HCl  +  O}}\]

Coloured substance + O → Colourless substance

In the presence of moisture, it bleaches vegetable or organic matter. Chloride has a long-lasting bleaching effect.

However, because SO2 is produced by reduction, its bleaching effect is only brief.

$  {\text{S}}{{\text{O}}_{{\text{2 }}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O}} \to {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ + 2H }} \ \\ $

$  {\text{S}}{{\text{O}}_{\text{3}}}^{{\text{2--}}}{\text{ +  Coloured material}} \to {\text{S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{ +  Reduced coloured material }}\left( {{\text{colourless}}} \right) \ \\ $ 

Reduced Coloured material (colourless) \[\xrightarrow{{{{\text{O}}_2}{\text{ of air}}}}\]material

(vii) Reaction with \[{\text{NaOH}}\]

Common to \[{\text{B}}{{\text{r}}_{\text{2}}}\], \[{{\text{I}}_{\text{2}}}\], and \[{\text{C}}{{\text{l}}_{\text{2}}}\] (however \[{{\text{F}}_{\text{2}}}\] is distinct, as \[{\text{O}}{{\text{F}}_2}\] or \[{{\text{O}}_2}\] is formed)

$  \left( {\text{a}} \right){\text{ 2NaOH}}\left( {{\text{cold  dilute}}} \right){\text{ +  C}}{{\text{l}}_{\text{2}}} \to {\text{NaCl  +  NaClO  +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  \left( {\text{b}} \right){\text{ 6NaOH }}\left( {{\text{hot  concentrated}}} \right){\text{  +  3C}}{{\text{l}}_{\text{2}}} \to {\text{5NaCl  +  NaCl}}{{\text{O}}_{\text{3}}}{\text{ +  3}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

(viii) Reaction with Hypo solution

This reaction is common with \[{\text{C}}{{\text{l}}_{\text{2}}}\] & \[{\text{B}}{{\text{r}}_{\text{2}}}\] but with \[{{\text{I}}_{\text{2}}}\] it is different

\[{\text{N}}{{\text{a}}_2}{{\text{S}}_2}{{\text{O}}_3}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O  +  C}}{{\text{l}}_{\text{2}}} \to {\text{ N}}{{\text{a}}_2}{\text{S}}{{\text{O}}_4} + {\text{2HCl}} + {\text{S}} \downarrow \left( {{\text{colloidal}}} \right){\text{ }}\]

Thiosulphate ions are disproportionated into \[{\text{S}}{{\text{O}}_4}^{2 - }\] and \[{\text{S}}\] in this reaction, whereas \[{\text{C}}{{\text{l}}_{\text{2}}}\] is reduced to \[{\text{C}}{{\text{l}}^ - }\].

(ix) Reaction with dry slaked lime, \[{\text{CaC}}{{\text{l}}_{\text{2}}}\]

It produces bleaching powder.

\[{\text{2Ca}}{\left( {{\text{OH}}} \right)_{\text{2}}}{\text{ +  2C}}{{\text{l}}_{\text{2}}} \to {\text{Ca}}{\left( {{\text{OCl}}} \right)_{\text{2}}}{\text{ +  CaC}}{{\text{l}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(x) Reaction with metals & non–metals: The appropriate chlorides are formed.

$ {\text{2Al  +  3C}}{{\text{l}}_{\text{2}}} \to {\text{2AlC}}{{\text{l}}_{\text{3}}}{\text{ ; 2Na  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2NaCl }} \ \\ $

$  {{\text{P}}_{\text{4}}}{\text{ + 6C}}{{\text{l}}_{\text{2}}} \to {\text{4PC}}{{\text{l}}_{\text{3}}}{\text{ ; }}{{\text{S}}_{\text{8}}}{\text{ + 4C}}{{\text{l}}_{\text{2}}} \to {\text{4}}{{\text{S}}_{\text{2}}}{\text{C}}{{\text{l}}_{\text{2}}} \ \\ $ 

It has a strong attraction to hydrogen. It becomes \[{\text{HCl}}\] when it combines with hydrogen-containing substances.

$  {{\text{H}}_{\text{2}}}{\text{ +  C}}{{\text{l}}_{\text{2}}} \to {\text{2HCl ; }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{S  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2 HCl  +  S ; }} \ \\ $

$  {{\text{C}}_{{\text{10}}}}{{\text{H}}_{{\text{16}}}}{\text{ +  8 C}}{{\text{l}}_{\text{2}}} \to {\text{16 HCl  +  10 C}} \ \\ $ 

Uses

(i)  for woodpulp bleaching (required for the manufacture of paper and rayon). bleaching cotton and textiles, and 

(ii) in the production of dyes, medicines, and organic chemicals like \[{\text{CC}}{{\text{l}}_{\text{4}}}{\text{, CHC}}{{\text{l}}_{\text{3}}}\], DDT, and refrigerants, among other things.

(iii) in the gold and platinum extraction process.

(iv) in the manufacture of hazardous gases such as phosgene (\[{\text{COC}}{{\text{l}}_{\text{2}}}\]), tear gas (\[{\text{CC}}{{\text{l}}_{\text{3}}}{\text{N}}{{\text{O}}_{\text{2}}}\]), and mustard gas, and 

(v) in the sterilisation of drinking water (\[{\text{ClC}}{{\text{H}}_{\text{2}}}{\text{C}}{{\text{H}}_{\text{2}}}{\text{SC}}{{\text{H}}_{\text{2}}}{\text{C}}{{\text{H}}_{\text{2}}}{\text{Cl}}\]).

3. Bromine (\[{\text{B}}{{\text{r}}_{\text{2}}}\])

Preparation 

(i) Common method

\[{\text{2 NaBr  +  3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{conc}}{\text{.}}} \right){\text{ +  Mn}}{{\text{O}}_{\text{2}}}\xrightarrow{\Delta }{\text{B}}{{\text{r}}_{\text{2}}}{\text{ +  MnS}}{{\text{O}}_{\text{4}}}{\text{ + 2NaHS}}{{\text{O}}_{\text{4}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(ii) From Sea-water

The predominant component is \[{\text{NaCl}}\], however there is also some \[{\text{NaBr}}\] in sea water. When bromine vapours are evolved, \[{\text{C}}{{\text{l}}_{\text{2}}}\] gas is transported via sea water.

\[{\text{2 B}}{{\text{r}}^{\text{--}}}\left( {{\text{aq}}} \right){\text{ +  C}}{{\text{l}}_{\text{2}}} \to {\text{2C}}{{\text{l}}^{{\text{-- }}}}\left( {{\text{aq}}{\text{.}}} \right){\text{  +  B}}{{\text{r}}_{\text{2}}}\]

Properties

(i) A reddish-brown liquid that is fairly water soluble. It also produces hydrates, such as \[{\text{C}}{{\text{l}}_{\text{2}}}\].

\[\left( {{\text{B}}{{\text{r}}_{\text{2}}}{\text{.8}}{{\text{H}}_{\text{2}}}{\text{O}}} \right) \leftarrow {\text{Clathrate compound}}\]

(ii) The rest of the reactions are the same as when using \[{\text{C}}{{\text{l}}_{\text{2}}}\].

4. Iodine (I2)

Preparation

(i) Common Method

\[{\text{2NaI + 3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\left( {{\text{conc}}{\text{.}}} \right){\text{ + Mn}}{{\text{O}}_{\text{2}}}\xrightarrow{\Delta }{{\text{I}}_{\text{2}}}{\text{ +  MnS}}{{\text{O}}_{\text{4}}}{\text{ +  2NaHS}}{{\text{O}}_{\text{4}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

(ii) From Caliche or Crude chile Saltpetre

Iodine is mostly obtained from \[{\text{NaI}}{{\text{O}}_3}\] (sodium iodate), which is found in nature together with \[{\text{NaN}}{{\text{O}}_3}\] (chile saltpetre). A minor quantity of \[{\text{NaI}}{{\text{O}}_3}\] is found. The mother liquor contains \[{\text{NaI}}{{\text{O}}_3}\] after \[{\text{NaN}}{{\text{O}}_3}\] crystallisation (soluble). \[{\text{NaHS}}{{\text{O}}_{\text{3}}}\] is added to this solution, and 2 precipitates as a result.

\[{\text{2I}}{{\text{O}}_{\text{3}}}^{\text{--}}{\text{ +  5HS}}{{\text{O}}_{\text{3}}}^{\text{--}} \to {\text{3HS}}{{\text{O}}_{\text{4}}}^{\text{--}}{\text{ +  2S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}{\text{ +  }}{{\text{I}}_{\text{2}}}\left( {\text{s}} \right){\text{ + }}{{\text{H}}_{\text{2}}}{\text{O}}\]

(iii) From Sea-Weeds

Seaweeds are dried and burned in shallow pits, and the ash that remains is known as kelp. Chlorides, carbonates, sulphates, and iodides of sodium and potassium are dissolved in ash when it is extracted with hot water. On concentration, the solution separates off all  the iodide, leaving only the iodide in the solution. In iron retorts, the solution is combined with \[{\text{Mn}}{{\text{O}}_{\text{2}}}\] and concentrated \[{{\text{H}}_2}{\text{S}}{{\text{O}}_{\text{4}}}\]. Iodine that has been liberated is collected in aludels, a type of earthenware.

\[{\text{2NaI +  Mn}}{{\text{O}}_{\text{2}}}{\text{ + 3}}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{2NaHS}}{{\text{O}}_{\text{4}}}{\text{ +  MnS}}{{\text{O}}_{\text{4}}}{\text{ +  }}{{\text{I}}_{\text{2}}}{\text{ +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\]

$\left( {{\text{iv}}} \right){\text{ 2KI + C}}{{\text{l}}_{\text{2}}} \to {\text{2KCl + I2 }} \ \\ $

  $\left( {\text{v}} \right){\text{ 2KI +  }}{{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2KOH  +  }}{{\text{I}}_{\text{2}}}{\text{ }} \ \\ $

 $ \left( {{\text{vi}}} \right){\text{ CuS}}{{\text{O}}_{\text{4}}}{\text{ +  2KI}} \to {{\text{K}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  Cu}}{{\text{I}}_{\text{2}}}{\text{ ; 2Cu}}{{\text{I}}_{\text{2}}} \to {\text{C}}{{\text{u}}_{\text{2}}}{{\text{I}}_{\text{2}}}{\text{ +  }}{{\text{I}}_{\text{2}}} \ \\ $ 

This 2 dissolves in K, creating 3, and because \[{{\text{3}}^ - }\] ions are yellow, the solution takes on a yellow hue

Properties

(i) It is a dark violet solid that sublimates and is the least soluble (among the halogens) in water, but much more soluble in K(aq.) due to the creation of \[{{\text{K}}_3}\]. \[{\text{K}}{{\text{F}}_3}\] cannot be created in the same way since F does not have ‘d' orbitals. As a result, \[{\text{s}}{{\text{p}}^3}{\text{d}}\] hybridization with F is not conceivable.

(ii) It can be made violet by dissolving it in organic solvents such as \[{\text{CHC}}{{\text{l}}_3},{\text{CC}}{{\text{l}}_3}\] and others.

(iii) Reaction with hypo

iodometric titrations

\[{{\text{S}}_{\text{2}}}{{\text{O}}_{\text{3}}}^{{\text{2--}}}\left( {{\text{thiosulphate ions}}} \right){\text{  +  }}{{\text{I}}_{\text{2}}} \to {{\text{S}}_4}{{\text{O}}_6}^{2 - }{\text{ }}\left( {{\text{tetrathionate ions}}} \right){\text{  +  2}}{{\text{I}}^{\text{--}}}\]

This reaction is the foundation for iodometric (direct 2 titration) titration, which is used to estimate iodine levels using starch as a marker.

(iv) Reaction with \[{\text{N}}{{\text{H}}_3}\]

\[{\text{N}}{{\text{H}}_{\text{3}}}\left( {\text{g}} \right){\text{  +  }}{{\text{I}}_{\text{2}}} \to {\text{No Reaction}}\]

\[{\text{NH}}\left( {{\text{aq}}} \right){\text{  +  }}{{\text{I}}_2}\left( {\text{s}} \right)\xrightarrow[{A{\text{ }}slurry{\text{ }}is{\text{ }}formed{\text{ }}which{\text{ }}can{\text{ }}be{\text{ }}dried{\text{ }}and{\text{ }}on{\text{ }}hammering{\text{ }}it{\text{ }}explodes{\text{ }}causing{\text{ }}sound{\text{ }}\left( {crakers} \right)}]{}{\text{N}}{{\text{I}}_3}{\text{ }}{\text{. N}}{{\text{H}}_3}{\text{  + 3HI}}\]

\[{\text{8N}}{{\text{I}}_{\text{3}}}{\text{. N}}{{\text{H}}_{\text{3}}} \to {\text{5}}{{\text{N}}_{\text{2}}}{\text{ +  9 }}{{\text{I}}_{\text{2}}}{\text{ +  6N}}{{\text{H}}_{\text{4}}}{\text{I }}\]

(v) Reaction with \[{\text{KCl}}{{\text{O}}_{\text{3}}}\] or \[{\text{KBr}}{{\text{O}}_{\text{3}}}\]

\[{\text{2 KCl}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{I}}_{\text{2}}}\xrightarrow{\Delta }{\text{2KI}}{{\text{O}}_{\text{3}}}{\text{ +  C}}{{\text{l}}_{\text{2}}}{\text{ ; 2KBr}}{{\text{O}}_{\text{3}}}{\text{ +  }}{{\text{I}}_{\text{2}}}\xrightarrow{\Delta }{\text{2KI}}{{\text{O}}_{\text{3}}}{\text{ +  B}}{{\text{r}}_{\text{2}}}\]

(vi) Reaction with ozone (dry)

\[{\text{2}}{{\text{I}}_{\text{2}}}\left( {\text{s}} \right){\text{ + 3}}{{\text{O}}_{\text{3}}}\left( {\text{g}} \right) \to {\text{ }}{{\text{I}}_{\text{4}}}{{\text{O}}_{\text{9}}}\left( {\text{s}} \right)\]

4O9 is an ionic compound made up of 3+ and O3–)3 that has a metallic character (low.E, low E.N.) As with \[{\text{C}}{{\text{l}}_{\text{2}}}\] and \[{\text{B}}{{\text{r}}_{\text{2}}}\], there is a frequent reaction with \[{\text{NaOH}}\]. There is a reversible reaction with \[{{\text{H}}_{\text{2}}}\]

Oxides of Chlorine

Chlorine dioxide (ClO2 ) 

Preparation

\[\left( {\text{i}} \right){\text{ 2 Cl}}{{\text{O}}_{\text{3}}}^{\text{--}}\left( {{\text{aq}}} \right){\text{  +  S}}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right)\xrightarrow{{{{\text{H}}^ + }}}{\text{2Cl}}{{\text{O}}_{\text{2}}}\left( {\text{g}} \right){\text{  +  S}}{{\text{O}}_{\text{4}}}^{{\text{2--}}}\left( {{\text{aq}}} \right)\]

Sodium and potassium chlorates can be used.

$ \left( {{\text{ii}}} \right){\text{ 2KCl}}{{\text{O}}_{\text{3}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{{\text{C}}_{\text{2}}}{{\text{O}}_{\text{4}}}\xrightarrow{{{\text{9}}{{\text{0}}^0}}}{\text{2Cl}}{{\text{O}}_2}{\text{(g) + 2C}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{K}}_{\text{2}}}{{\text{C}}_{\text{2}}}{{\text{O}}_{\text{4}}}{\text{ + 2}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $

$ \left( {{\text{iii}}} \right){\text{ 2AgCl}}{{\text{O}}_{\text{3}}}{\text{ + C}}{{\text{l}}_{\text{2}}}\xrightarrow{{{\text{9}}{{\text{0}}^0}}}{\text{2AgCl}} \downarrow \left( {{\text{white}}} \right){\text{ + 2Cl}}{{\text{O}}_{\text{2}}}{\text{ + }}{{\text{O}}_{\text{2}}}{\text{ }} \ \\ $ $ \left( {{\text{iv}}} \right){\text{ C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{6}}}{\text{ + }}{{\text{N}}_{\text{2}}}{{\text{O}}_{\text{4}}} \rightleftharpoons {\text{ Cl}}{{\text{O}}_{\text{2}}}{\text{ + }}\left[ {{\text{N}}{{\text{O}}_{\text{2}}}^{\text{ + }}{\text{ }}} \right]{\text{ }}\left[ {{\text{Cl}}{{\text{O}}_{\text{4}}}^{\text{--}}} \right] \ \\ $ 

Properties

(i) A yellow gas that is flammable and soluble in water at room temperature. It also has oxidising properties. (It kills bacteria more effectively than \[{\text{C}}{{\text{l}}_{\text{2}}}\].)

(ii) Reaction with ozone

\[{\text{2Cl}}{{\text{O}}_{\text{2}}}{\text{ +  2}}{{\text{O}}_{\text{3}}}\xrightarrow{{{{\text{H}}^ + }}}{\text{ C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{6}}}\left( {{\text{yellow solid}}} \right){\text{  +  2}}{{\text{O}}_{\text{2}}}\]

                                  dichlorine hexa oxide

\[{{\text{O}}_{\text{3}}}\] acts as an oxidising agent in this process.

\[{\text{C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{6}}}\] (s) is a mixed anhydride of \[{\text{HCl}}{{\text{O}}_{\text{3}}}\] and \[{\text{HCl}}{{\text{O}}_{\text{4}}}\] because it gives a combination of these two acids when dissolved in water.\[{\text{C}}{{\text{l}}_{\text{2}}}{{\text{O}}_{\text{6}}}\] occurs in solid form as \[{\text{Cl}}{{\text{O}}_{\text{2}}}^ + \] and \[{\text{Cl}}{{\text{O}}_4}^ - \].

(iii) Reaction with Alkaline \[{{\text{H}}_2}{{\text{O}}_2}\]

\[{{\text{H}}_2}{{\text{O}}_2}\] functions as a reducing agent in this process. \[{\text{Cl}}{{\text{O}}_2}\] is reduced to \[{\text{Cl}}{{\text{O}}_2}^ - \]

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}} \to {\text{2}}{{\text{H}}^{\text{ + }}}{\text{ +  }}{{\text{O}}_{\text{2}}}{\text{  +  2}}{{\text{e}}^{\text{--}}}{\text{ ; }}{{\text{e}}^{\text{--}}}{\text{  +  Cl}}{{\text{O}}_{\text{2}}} \to {\text{Cl}}{{\text{O}}_{\text{2}}}^{\text{--}}{\text{ ]  \times  2}} \ \\ $

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{ +  2Cl}}{{\text{O}}_{\text{2}}} \to {\text{2Cl}}{{\text{O}}_{\text{2}}}^{\text{--}}{\text{ +  2}}{{\text{H}}^{\text{ + }}}{\text{  +  }}{{\text{O}}^{\text{2}}} \ \\ $

$  {{\text{H}}_{\text{2}}}{{\text{O}}_{\text{2}}}{\text{  +  2Cl}}{{\text{O}}_{\text{2}}}{\text{  +  2O}}{{\text{H}}^{\text{--}}} \to {\text{2Cl}}{{\text{O}}_{\text{2}}}^{\text{--}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O  +  }}{{\text{O}}_{\text{2}}} \ \\ $ 

(iv) Reaction with HΙ

In this reaction, H acts as a reducing agent, converting \[{\text{ Cl}}{{\text{O}}_{\text{2}}}\] to \[{\text{C}}{{\text{l}}^{\text{--}}}\] before being oxidised to 2.

$\left[ {{\text{5}}{{\text{e}}^{\text{--}}}{\text{ +  4}}{{\text{H}}^{\text{ + }}}{\text{ +  Cl}}{{\text{O}}_{\text{2}}} \to {\text{C}}{{\text{l}}^{\text{--}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}} \right]{\text{  \times  2 }} \ \\ $

$  \left[ {{\text{2}}{{\text{I}}^{\text{--}}} \to {\text{ }}{{\text{I}}_{\text{2}}}{\text{ +  2}}{{\text{e}}^{\text{--}}}{\text{ }}} \right]{\text{  \times  5 }} \ \\ $

$  {\text{2Cl}}{{\text{O}}_{\text{2}}}{\text{  +  8}}{{\text{H}}^{\text{ + }}}{\text{ +  10 }}{{\text{I}}^{\text{--}}} \to {\text{5}}{{\text{I}}_{\text{2}}}{\text{ 2C}}{{\text{l}}^{\text{--}}}{\text{ +  4}}{{\text{H}}_{\text{2}}}{\text{O}} \ \\ $ 

Hydra Acids (Halogen Acids):

\[{\text{HCl}}\], \[{\text{HBr}}\] & \[{\text{HI}}\] : 

Preparation: 

(i) By direct combination of elements

$ {{\text{H}}_{\text{2}}}{\text{  +  C}}{{\text{l}}_{\text{2}}} \to {\text{2HCl }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{  +  B}}{{\text{r}}_{\text{2}}}\xrightarrow{{{\text{Pt}}}}{\text{ 2HBr }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{  +  }}{{\text{I}}_{\text{2}}}\xrightarrow{{{\text{Pt 45}}{{\text{0}}^0}{\text{C}}}}{\text{2HI}} \ \\ $ 

(ii) By heating a halide with acid

$ {\text{NaCl  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{NaHS}}{{\text{O}}_{\text{4}}}{\text{ +  HCl }} \ \\ $

$  {\text{NaHS}}{{\text{O}}_{\text{4}}}{\text{ +  NaCl}} \to {\text{N}}{{\text{a}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  HCl}} \ \\ $ 

\[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] is used as an acid for \[{\text{HCl}}\], while \[{{\text{H}}_3}{\text{P}}{{\text{O}}_4}\] is used for \[{\text{HBr}}\] and \[{\text{H}}\].

(a) Because \[{{\text{P}}_{\text{2}}}{{\text{O}}_{\text{5}}}{\text{ }}\left( {{{\text{P}}_{\text{4}}}{{\text{O}}_{{\text{10}}}}} \right)\]and quick lime chemically react with gas, \[{\text{HCl}}\] cannot be dried over them.

$ {\text{CaO  +  2HCl}} \to {\text{CaC}}{{\text{l}}_{\text{2}}}{\text{ +  }}{{\text{H}}_{\text{2}}}{\text{O }} \ \\ $

$  {{\text{P}}_{\text{4}}}{{\text{O}}_{{\text{10}}}}{\text{  +  3HCl}} \to {\text{POC}}{{\text{l}}_{\text{3}}}{\text{ +  3HP}}{{\text{O}}_{\text{3}}} \ \\ $ 

As a result, \[{\text{HCl}}\] is dried by passing it through a condenser of \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\].

(b) Heating bromide (iodide) with conc. \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] does not produce \[{\text{HBr}}\] (or \[{\text{H}}\]), because \[{\text{HBr}}\] and \[{\text{H}}\]are strong reducing agents that reduce \[{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}\] to \[{\text{S}}{{\text{O}}_{\text{4}}}\] and oxidise themselves to bromine and iodine, respectively.

$ {\text{KX  +  }}{{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}} \to {\text{KHS}}{{\text{O}}_{\text{4}}}{\text{ +  HX }} \ \\ $

$  {{\text{H}}_{\text{2}}}{\text{S}}{{\text{O}}_{\text{4}}}{\text{ +  2HX}} \to {\text{S}}{{\text{O}}_{\text{2}}}{\text{ +  }}{{\text{X}}_{\text{2}}}{\text{  +  2}}{{\text{H}}_{\text{2}}}{\text{O}}\left( {{\text{X  =  Br or n}}} \right) \ \\ $ 

As a result, bromides and iodides are heated with conc. \[{\text{HP}}{{\text{O}}_{\text{3}}}\] to produce \[{\text{HBr}}\] and \[{\text{H}}\], respectively.

\[{\text{3KBr}}\left( {{\text{KI}}} \right){\text{  +  }}{{\text{H}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}} \to {{\text{K}}_{\text{3}}}{\text{P}}{{\text{O}}_{\text{4}}}{\text{ +  3HBr }}\left( {{\text{HI}}} \right)\]

(iii) By reaction of \[{{\text{P}}_{\text{4}}}\] (Laboratory Method)