Three oxoacids of nitrogen are : (i) \(HN{ O }_{ 2 },\) nitrous acid (ii) \(H{ N{ O } }_{ 3 },\) nitric acid and \({ H }_{ 2 }{ N }_{ 2 }{ O }_{ 2 },\) hyponitrous acid. O.S. of N in \(H{ NO }_{ 2 }\) = (+ 1) + x + 2 (- 2) = 0 or X= + 3
 \(3H\overset { +3 }{ { NO }_{ 2 } } \xrightarrow { Disproportionation } H\overset { +3 }{ { NO }_{ 3 } } +2\overset { +2 }{ NO } +{ H }_{ 2 }O\) 

The chemical equations for the reactions involved are :
  
Now 124 g of white P give HCl = 12 X 36.5 g
\(\therefore \) 62 g of white P will give HCl \(=\frac { 12\times 36.5 }{ 124 } \times 62=219g\)

The chemical equations for the reactions involved are :

Now 220 g of \({ P }_{ 4 }{ O }_{ 6 }\) require NaOH for neutralization = 320 g
\(\therefore \) 1.1 g of \({ P }_{ 4 }{ O }_{ 6 }\) will require NaOH = \(\frac { 320 }{ 220 } \times 1.1=1.6g\)
Now 1000 mL of 0.1 MHCl contain NaOH  = 40 X 0.1 = 4 g
In other words, 4 g of NaOH are present in 0.1 M NaOH = 1000 mL
\(\therefore \) 1.6 g of NaOH will be present in 0.1 M NaOH =  \(\frac { 1000 }{ 4 } \times 1.6=400\quad mL=0.4L\)

Reason: All these acids on losing a proton give their corresponding conjugate bases, i.e.,\(Cl{ O }^{ - },\ Cl{ O }_{ 2 }^{ - },\ Cl{ O }_{ 3 }^{ - }\ and\ Cl{ O }_{ 4 }^{ - }.\) Their structures are :
         
Since oxygen is more electronegative than chlorine, therefore, the dispersal of the -ve charge present on oxygen atom (singly bonded to CI)increases as the number of oxygen atoms attached by a double bond to chlorine increases due to \(d\pi -p\pi \) back bonding. In other words, stability of the conjugate bases increases in the order: \({ ClO }^{ - }<{ ClO }_{ 2 }^{ - }<{ ClO }_{ 3 }^{ - }<{ ClO }_{ 4 }^{ - }.\)
Thus, due to increase in stability of the conjugate base, acidic strength increases in the same order :\( \ (HClO\) )

\({ NO }_{ 3 }^{ - }\left( aq \right) +3{ Fe }^{ 2+ }\left( aq \right) +4{ H }^{ + }\left( aq \right) \longrightarrow NO\left( g \right) +3{ Fe }^{ 3+ }\left( aq \right) +3{ H }_{ 2 }O\left( l \right) \)
\(\\ Nitrate\quad ion\quad \quad \quad \quad \quad \quad \quad \quad \quad \quad Nitric\quad oxide\)
 \({ \left[ Fe{ \left( { H }_{ 2 }O \right) }_{ 6 } \right] }^{ 2+ }+NO\left( g \right) \longrightarrow { \left[ Fe{ \left( { H }_{ 2 }O \right) }_{ 5 }NO \right] }^{ 2+ }+{ H }_{ 2 }O\)
\(\\ \quad \quad \quad \quad \quad \quad \quad \quad \quad \quad Pentaaquanitrosonium\quad iron\left( I \right) \\ \quad \quad \quad \quad \quad \quad \quad \quad \quad \quad \quad \quad \left( Brown\quad complex \right) \)