Group 15 Elements: The Nitrogen Family

Group 15 elements are nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi), all with the valence configuration $n{s}^{2}n{p}^3$. Down the group they change from non-metals (N, P) through metalloids (As, Sb) to a metal (Bi), and the +3 oxidation state becomes more stable than +5. This page covers the trends of group 15 elements and their key compounds: ammonia, nitric acid, oxides of nitrogen, phosphorus allotropes, phosphine, halides, oxoacids and fertilizers. JEE Main and NEET test the trends and the unique behaviour of nitrogen; JEE Advanced also asks about the compounds.

1. The Nitrogen Family at a Glance

Group 15 of the periodic table contains nitrogen (N), phosphorus (P), arsenic (As), antimony (Sb) and bismuth (Bi); the synthetic element moscovium (Mc) also sits here. N and P are non-metals, As and Sb are metalloids and Bi is a metal. Every member has five valence electrons, $n{s}^{2}n{p}^3$, with a half-filled $p$ subshell. The group is also called the pnictogens.

Occurrence. Nitrogen makes up 78% of air by volume and is found in the soil as sodium nitrate (Chile saltpetre) and potassium nitrate (Indian saltpetre); it is part of every protein. Phosphorus occurs in the apatite minerals, $\mathrm{C}{\mathrm{a}}_9(\mathrm{P}{\mathrm{O}}_4{)}_6\cdot \mathrm{Ca}{\mathrm{X}}_2$ (X = F, Cl or OH), and in bones, teeth, DNA and ATP. Arsenic, antimony and bismuth occur mainly as sulphides.

1.1 Profiles of the Five Elements

2. Trends in Physical Properties

Key data for the five elements:

• Atomic radius increases down the group because a new shell is added each time. The rise from As to Bi is small because the filled $d$ and $f$ orbitals shield the nucleus poorly.
• Ionisation enthalpy decreases down the group as size increases. It is much higher than for group 14 because of the larger nuclear charge and the extra-stable half-filled $n{p}^3$ subshell.
• Electronegativity decreases down the group, but only slightly after P (2.1 to 1.9).
• Metallic character increases: N and P are non-metals, As and Sb metalloids, Bi a metal.
• Physical state: nitrogen is a diatomic gas; the rest are solids. All except nitrogen show allotropy.
• Melting and boiling points: boiling points generally rise down the group. Melting points rise up to arsenic and then fall to bismuth, because Sb and Bi are held by weak metallic bonding in which the $n{s}^2$ pair takes little part.

3. Oxidation States and the Inert Pair Effect

Covalent character falls down the group, $\mathrm{N}>\mathrm{P}>\mathrm{As}>\mathrm{Sb}>\mathrm{Bi}$, as the atoms become larger and less electronegative. The common oxidation states are $-3$, +3 and +5.

• The $-3$ state becomes less common down the group as size grows and electronegativity falls; bismuth hardly forms any compound in the $-3$ state.
• The +5 state becomes less stable, and +3 more stable, down the group. This is the inert pair effect: the $n{s}^2$ electrons of heavy atoms are poorly shielded and reluctant to bond. The only well-characterised Bi(V) compound is $\mathrm{Bi}{\mathrm{F}}_5$, and Bi(V) is a strong oxidising agent.
• Nitrogen shows every state from $-3$ to +5. With oxygen it shows +1, +2 and +4 too. In acid solution the +1 to +4 states tend to disproportionate, for example $3HN{\mathrm{O}}_2\to \mathrm{HN}{\mathrm{O}}_3+{\mathrm{H}}_2\mathrm{O}+2NO$.
• Phosphorus shows +1 and +4 in some oxoacids; its +3 and +1 states disproportionate on heating, for example $4{\mathrm{H}}_3\mathrm{P}{\mathrm{O}}_3\to 3{\mathrm{H}}_3\mathrm{P}{\mathrm{O}}_4+\mathrm{P}{\mathrm{H}}_3$.
• Maximum covalency: nitrogen has no $d$ orbitals, so it can form at most four bonds (as in $\mathrm{N}{\mathrm{H}}_4^{+}$). Phosphorus and the heavier elements can use $d$ orbitals and reach covalency six ($[\mathrm{P}{\mathrm{F}}_6{]}^{-}$).

4. Anomalous Behaviour of Nitrogen

Nitrogen differs from the rest of the group because of its small size, high electronegativity, high ionisation enthalpy and lack of $d$ orbitals. The main differences are:

• Multiple bonds: nitrogen forms strong $p\pi$-$p\pi$ bonds with itself, carbon and oxygen ($\mathrm{N}\equiv \mathrm{N}$, $\mathrm{C}\equiv \mathrm{N}$, N=O). Phosphorus forms $d\pi$-$p\pi$ bonds instead, as in $\mathrm{PO}{\mathrm{X}}_3$ and ${\mathrm{R}}_3\mathrm{P}=\mathrm{O}$.
• Molecular form: nitrogen is a diatomic gas, ${\mathrm{N}}_2$; phosphorus is the tetratomic solid ${\mathrm{P}}_4$.
• Low reactivity: the $\mathrm{N}\equiv \mathrm{N}$ bond (one $\sigma$ + two $\pi$) has a bond enthalpy of 941.4 kJ mol${}^{-1}$, so ${\mathrm{N}}_2$ is nearly inert at room temperature.
• Weak catenation: the N-N single bond is weaker than the P-P single bond, because the lone pairs on the two small N atoms repel strongly at the short bond length.
• No $d\pi$-$p\pi$ bonding: ${\mathrm{R}}_3\mathrm{P}=\mathrm{O}$ exists, but nitrogen forms ${\mathrm{R}}_3\mathrm{N}\to \mathrm{O}$ (an N-oxide with a coordinate bond).
• Covalency limited to four; so $\mathrm{NC}{\mathrm{l}}_5$ does not exist while $\mathrm{PC}{\mathrm{l}}_5$ does.
• Hydrogen bonding: only $\mathrm{N}{\mathrm{H}}_3$ among the hydrides is hydrogen bonded, which raises its boiling point and solubility.

5. Chemical Reactivity of the Group

5.1 Hydrides ($\mathrm{E}{\mathrm{H}}_3$)

Every element forms a gaseous hydride $\mathrm{E}{\mathrm{H}}_3$: ammonia $\mathrm{N}{\mathrm{H}}_3$, phosphine $\mathrm{P}{\mathrm{H}}_3$, arsine $\mathrm{As}{\mathrm{H}}_3$, stibine $\mathrm{Sb}{\mathrm{H}}_3$ and bismuthine $\mathrm{Bi}{\mathrm{H}}_3$. They are made by hydrolysing a metal nitride, phosphide, arsenide or antimonide:

Nitrogen also forms hydrazine, ${\mathrm{N}}_2{\mathrm{H}}_4$ (from ammonia and sodium hypochlorite), and hydrazoic acid, $\mathrm{H}{\mathrm{N}}_3$:

In every hydride the central atom carries one lone pair, so the molecules are pyramidal. The trends are:

• Thermal stability decreases $\mathrm{N}{\mathrm{H}}_3>\mathrm{P}{\mathrm{H}}_3>\mathrm{As}{\mathrm{H}}_3>\mathrm{Sb}{\mathrm{H}}_3>\mathrm{Bi}{\mathrm{H}}_3$, because the E-H bond gets longer and weaker as the central atom grows.
• Reducing character increases down the group for the same reason: the weaker the E-H bond, the more easily H is given up. $\mathrm{Bi}{\mathrm{H}}_3$ is the strongest reducing agent. The heavier hydrides react with metal ions to give phosphides, arsenides or antimonides, and all except ammonia are highly poisonous.
• Basicity decreases $\mathrm{N}{\mathrm{H}}_3>\mathrm{P}{\mathrm{H}}_3>\mathrm{As}{\mathrm{H}}_3>\mathrm{Sb}{\mathrm{H}}_3\ge \mathrm{Bi}{\mathrm{H}}_3$: the lone pair is spread over a larger volume on the bigger atoms, so it is less available (Lewis base strength falls).
• Bond angle decreases from $107.8^{\circ }$ towards $90^{\circ }$ (Figure 3).
• Boiling point rises from $\mathrm{P}{\mathrm{H}}_3$ to $\mathrm{Bi}{\mathrm{H}}_3$ with molecular size, but $\mathrm{N}{\mathrm{H}}_3$ is anomalously high because of intermolecular hydrogen bonding.

5.2 Oxides

The elements form two main series of oxides, ${\mathrm{E}}_2{\mathrm{O}}_3$ and ${\mathrm{E}}_2{\mathrm{O}}_5$; nitrogen also forms oxides in the +1, +2 and +4 states.

• Acidic character falls down the group. Oxides of N and P are strongly acidic (except the neutral NO and ${\mathrm{N}}_2\mathrm{O}$), oxides of As are weakly acidic, those of Sb amphoteric and those of Bi weakly basic.
• Reason: a small, electronegative central atom pulls the electron pair of the O-H bond towards itself (after the oxide dissolves in water), which helps release ${\mathrm{H}}^{+}$. As the central atom grows, this pull weakens and the oxide becomes basic.
• The oxide in the higher oxidation state is more acidic than the one in the lower state (for example, ${\mathrm{N}}_2{\mathrm{O}}_5>{\mathrm{N}}_2{\mathrm{O}}_3$).

5.3 Halides

Two series are known: trihalides $\mathrm{E}{\mathrm{X}}_3$ (all elements) and pentahalides $\mathrm{E}{\mathrm{X}}_5$ (all except nitrogen). Examples: $\mathrm{NC}{\mathrm{l}}_3,\mathrm{PC}{\mathrm{l}}_3,\mathrm{AsC}{\mathrm{l}}_3,\mathrm{SbC}{\mathrm{l}}_3,\mathrm{BiC}{\mathrm{l}}_3$ and $\mathrm{PC}{\mathrm{l}}_5,\mathrm{AsC}{\mathrm{l}}_5,\mathrm{SbC}{\mathrm{l}}_5$.

• Trihalides are covalent (except $\mathrm{Bi}{\mathrm{F}}_3$, which is ionic), pyramidal and $s{p}^3$ hybridised like the hydrides. Of the nitrogen trihalides, only $\mathrm{N}{\mathrm{F}}_3$ is stable.
• Pentahalides ($s{p}^{3}d$, trigonal bipyramidal) are more covalent than the trihalides, because the central atom in the +5 state has greater polarising power. They are thermally less stable than the trihalides.
• Nitrogen forms no pentahalide, because it has no $d$ orbitals to expand its octet.

Hydrolysis. All trihalides except $\mathrm{N}{\mathrm{F}}_3$ are hydrolysed by water. Those of P and As go all the way to the oxoacid or oxide; those of Sb and Bi stop at the oxychloride. $\mathrm{NC}{\mathrm{l}}_3$ hydrolyses by a different route and gives ammonia (see Solved Example 14):

Lewis acid behaviour. The trihalides of P, As and Sb and all the pentahalides accept electron pairs from Lewis bases:

Trihalides can also be oxidised to pentahalides, for example $\mathrm{P}{\mathrm{F}}_3+{\mathrm{F}}_2\to \mathrm{P}{\mathrm{F}}_5$. $\mathrm{PC}{\mathrm{l}}_5$ exists as molecules in the gas but is ionic, $[\mathrm{PC}{\mathrm{l}}_4{]}^{+}[\mathrm{PC}{\mathrm{l}}_6{]}^{-}$, in the solid; $\mathrm{PB}{\mathrm{r}}_5$ ($[\mathrm{PB}{\mathrm{r}}_4{]}^{+}\mathrm{B}{\mathrm{r}}^{-}$) and $\mathrm{P}{\mathrm{I}}_5$ are also ionic solids.

5.4 Reaction with Metals

All the elements combine with active metals to form binary compounds in the $-3$ state, such as calcium nitride $\mathrm{C}{\mathrm{a}}_3{\mathrm{N}}_2$, calcium phosphide $\mathrm{C}{\mathrm{a}}_3{\mathrm{P}}_2$, sodium arsenide $\mathrm{N}{\mathrm{a}}_3\mathrm{As}$, zinc antimonide $\mathrm{Z}{\mathrm{n}}_3\mathrm{S}{\mathrm{b}}_2$ and magnesium bismuthide $\mathrm{M}{\mathrm{g}}_3\mathrm{B}{\mathrm{i}}_2$.

6. Dinitrogen (${\mathrm{N}}_2$)

6.1 Preparation

Commercially, nitrogen is obtained by liquefying air and distilling it fractionally: liquid ${\mathrm{N}}_2$ (b.p. 77.2 K) boils off first, leaving liquid oxygen (b.p. 90 K). In the laboratory, a solution of ammonium chloride is warmed with sodium nitrite. Ammonium nitrite forms and decomposes at once (it is too unstable to store):

Traces of NO and $\mathrm{HN}{\mathrm{O}}_3$ are removed by passing the gas through aqueous sulphuric acid containing potassium dichromate. Other methods:

Urea also gives nitrogen with nitrous acid (made from $\mathrm{NaN}{\mathrm{O}}_2$ + HCl):

6.2 Properties

• Colourless, odourless, tasteless and non-toxic gas; very slightly soluble in water. It has two stable isotopes, ${}^{14}\text{N}$ and ${}^{15}\text{N}$.
• Inert at room temperature because of the strong triple bond, but reactive at high temperature:

Fixing nitrogen as calcium cyanamide (the cyanamide process): heating calcium carbide in nitrogen at about 1273-1373 K gives calcium cyanamide. The mixture of $\mathrm{CaC}{\mathrm{N}}_2$ and carbon (trade name nitrolim) is a fertilizer, because it slowly releases ammonia in moist soil.

6.3 Uses

Making ammonia and other nitrogen compounds; an inert atmosphere in electric bulbs, metallurgy and food packaging; protecting historic documents; liquid nitrogen as a refrigerant for biological samples and food, and in cryosurgery.

7. Ammonia ($\mathrm{N}{\mathrm{H}}_3$)

Nitrogen forms three hydrides: ammonia $\mathrm{N}{\mathrm{H}}_3$, hydrazine $\mathrm{N}{\mathrm{H}}_2\mathrm{N}{\mathrm{H}}_2$ and hydrazoic acid ${\mathrm{N}}_3\mathrm{H}$. Ammonia is present in traces in air, formed by bacterial decay of nitrogenous plant and animal matter (for example, urea: $\mathrm{N}{\mathrm{H}}_2\mathrm{CON}{\mathrm{H}}_2+2{\mathrm{H}}_2\mathrm{O}\to (\mathrm{N}{\mathrm{H}}_4{)}_2\mathrm{C}{\mathrm{O}}_3$, which gives off $\mathrm{N}{\mathrm{H}}_3$). Priestley isolated it in 1774 by heating ammonium chloride with lime and called it 'alkaline air'. Berthollet (1785) showed it contains nitrogen and hydrogen, and Davy (1800) established the formula $\mathrm{N}{\mathrm{H}}_3$.

7.1 Laboratory Preparation

(i) Heating an ammonium salt with a base. Slaked lime is the usual laboratory method:

(ii) Heating ammonium chloride with litharge (PbO):

(iii) Hydrolysis of metal nitrides:

(iv) Reducing nitrates and nitrites with zinc and caustic soda (nascent hydrogen):

(v) Hydrolysis of calcium cyanamide:

(vi) Heating ammonium salts on their own:

(vii) Heating urea with caustic soda:

Drying. Ammonia is dried over quicklime (CaO). The usual drying agents cannot be used because each reacts with ammonia:

7.2 Manufacture: Haber's Process

Fritz Haber's process combines nitrogen and hydrogen directly. The reaction is reversible and exothermic, and the number of gas molecules falls from 4 to 2:

Older catalysts that were tried include finely divided iron with molybdenum as promoter, finely divided nickel with soda lime on pumice, and finely divided osmium or uranium. Nitrogen for the process comes from liquid air. Hydrogen was once made by electrolysis of water or from water gas; today it comes mostly from natural gas. Figure 4 shows the older coke-based route for making the feed gas.

7.3 Physical Properties and Structure

• Colourless gas with a characteristic pungent smell; lighter than air.
• Highly soluble in water because it hydrogen bonds with water. Solubility increases with pressure and decreases with temperature.
• Easily liquefied (b.p. 239.7 K, f.p. 198.4 K). Liquid ammonia has a high enthalpy of vaporisation, which is why it is used in refrigeration.
• Structure: N is $s{p}^3$ hybridised with one lone pair, so the molecule is trigonal pyramidal. The H-N-H angle is $107.8^{\circ }$, not $109^{\circ }28^{\prime }$, because lone pair-bond pair repulsion squeezes the bonds together; N-H = 101.7 pm. Molecules are hydrogen bonded in the solid and liquid.

7.4 Chemical Properties

(i) Stability. Ammonia is quite stable, but decomposes at red heat or when electric sparks are passed:

(ii) Combustion. It burns in oxygen with a yellowish flame:

(iii) Basic nature. The lone pair on nitrogen accepts a proton (Lewis base), forming the ammonium ion through a coordinate bond. Its aqueous solution is a weak base, and it forms salts with acids (the white fumes of $\mathrm{N}{\mathrm{H}}_4\mathrm{Cl}$ are a test for $\mathrm{N}{\mathrm{H}}_3$ or HCl):

(iv) Reducing nature (oxidation of ammonia). Heated CuO or PbO oxidises ammonia to nitrogen:

Chlorine and bromine also oxidise ammonia. What forms depends on which reactant is in excess:

Nitrogen trichloride formed with excess chlorine is an explosive oily liquid. Iodine gives the explosive nitrogen triiodide ammoniate, which detonates when dry:

Hypochlorites, hypobromites and bleaching powder oxidise ammonia to nitrogen:

Restricted oxidation with air over heated platinum gauze (973-1073 K) gives nitric oxide. This is the first step of Ostwald's process for nitric acid (Section 9):

(v) Formation of amides. Heated sodium displaces hydrogen from dry ammonia:

(vi) Reactions of aqueous ammonia with metal salts. Ammonium hydroxide precipitates metal hydroxides. Some of these dissolve in excess ammonia because the metal ion forms an ammine complex; this is used in qualitative analysis.

(vii) Nessler's reagent. Potassium iodide added to mercuric chloride first precipitates red $\mathrm{Hg}{\mathrm{I}}_2$, which dissolves in excess KI to give ${\mathrm{K}}_2\mathrm{Hg}{\mathrm{I}}_4$. The alkaline solution of ${\mathrm{K}}_2\mathrm{Hg}{\mathrm{I}}_4$ is Nessler's reagent. It gives a brown precipitate (iodide of Millon's base) with even traces of ammonia or ammonium salts:

7.5 Uses of Ammonia

1. Liquid ammonia is used as a refrigerant.
2. Aqueous ammonia is a laboratory reagent in qualitative and quantitative analysis.
3. Manufacture of nitric acid (Ostwald's process), sodium bicarbonate and washing soda (Solvay process) and ammonium compounds.
4. Fertilizers such as ammonium sulphate, ammonium phosphate, calcium ammonium nitrate and urea; ammonium nitrate is also used in explosives.
5. A cleansing agent for removing grease.

8. Oxides of Nitrogen

Nitrogen forms six well-characterised oxides, in oxidation states +1 to +5. $\mathrm{N}{\mathrm{O}}_3$ and ${\mathrm{N}}_2{\mathrm{O}}_6$ are also reported but are unstable and poorly characterised.

8.1 Preparation

(a) ${\mathrm{N}}_2\mathrm{O}$ from ammonium nitrate. It is called laughing gas because it causes laughter; mixed with oxygen, it is used as a mild anaesthetic by dentists.

(b) NO by reducing nitrous acid or nitrites with $\mathrm{F}{\mathrm{e}}^{2+}$ or ${\mathrm{I}}^{-}$, or nitric acid with copper:

NO is also the intermediate in the manufacture of nitric acid from ammonia.

(c) ${\mathrm{N}}_2{\mathrm{O}}_3$ separates as an intense blue liquid or pale blue solid when an equimolar mixture of NO and $\mathrm{N}{\mathrm{O}}_2$ is cooled. On warming its colour fades as it dissociates back:

(d) $\mathrm{N}{\mathrm{O}}_2$ by heating a heavy-metal nitrate, or by reducing concentrated nitric acid with copper:

(e) ${\mathrm{N}}_2{\mathrm{O}}_4$ by cooling $\mathrm{N}{\mathrm{O}}_2$ (the brown colour fades):

(f) ${\mathrm{N}}_2{\mathrm{O}}_5$, the anhydride of nitric acid, by dehydrating nitric acid with phosphorus pentoxide at low temperature:

8.2 Properties

• All the oxides are oxidising agents. ${\mathrm{N}}_2\mathrm{O}$ even supports the combustion of sulphur and phosphorus. NO, which is thermally more stable, supports the burning of Mg and P but not of S, because the sulphur flame is not hot enough to decompose NO.
• ${\mathrm{N}}_2\mathrm{O}$ and NO are neutral; the rest are acidic. ${\mathrm{N}}_2{\mathrm{O}}_3$ is the anhydride of $\mathrm{HN}{\mathrm{O}}_2$ and ${\mathrm{N}}_2{\mathrm{O}}_5$ of $\mathrm{HN}{\mathrm{O}}_3$; $\mathrm{N}{\mathrm{O}}_2$ is a mixed anhydride: $2N{\mathrm{O}}_2+{\mathrm{H}}_2\mathrm{O}\to \mathrm{HN}{\mathrm{O}}_3+\mathrm{HN}{\mathrm{O}}_2$.
• NO and $\mathrm{N}{\mathrm{O}}_2$ have an odd number of electrons, so they are paramagnetic. $\mathrm{N}{\mathrm{O}}_2$ dimerises to colourless, diamagnetic ${\mathrm{N}}_2{\mathrm{O}}_4$ on cooling. NO combines with air at once to give reddish-brown $\mathrm{N}{\mathrm{O}}_2$: $2NO+{\mathrm{O}}_2\to 2N{\mathrm{O}}_2$.
• Liquid ${\mathrm{N}}_2{\mathrm{O}}_4$ ionises itself, ${\mathrm{N}}_2{\mathrm{O}}_4\rightleftharpoons \mathrm{N}{\mathrm{O}}^{+}+\mathrm{N}{\mathrm{O}}_3^{-}$, and is used as a non-aqueous solvent.
• Solid ${\mathrm{N}}_2{\mathrm{O}}_5$ is ionic, $[\mathrm{N}{\mathrm{O}}_2{]}^{+}[\mathrm{N}{\mathrm{O}}_3{]}^{-}$. In the gas, the ${\mathrm{N}}_2{\mathrm{O}}_5$ molecule has a bent N-O-N bridge.

9. Nitric Acid ($\mathrm{HN}{\mathrm{O}}_3$)

Nitric acid is the most important oxoacid of nitrogen; the alchemists called it aqua fortis ('strong water').

9.1 Preparation

(a) Laboratory method: heating a nitrate with concentrated sulphuric acid in a glass retort; the volatile $\mathrm{HN}{\mathrm{O}}_3$ distils over.

(b) Birkeland-Eyde process (historical). Air is passed through an electric arc at about 3000 °C. The reaction is endothermic, so the high temperature favours NO; the gases are then cooled quickly to about 1000 °C so that NO does not decompose:

(c) Ostwald's process (modern). Ammonia and air are passed over platinum-rhodium gauze. Industrial converters run at about 1023-1173 K (750-900 °C); NCERT quotes 500 K and 9 bar. NO is oxidised to $\mathrm{N}{\mathrm{O}}_2$ by air (the gases are cooled to about 50 °C) and absorbed in water:

Distillation concentrates the acid to about 68% by mass; further dehydration with concentrated sulphuric acid gives 98% acid. Fuming nitric acid is concentrated nitric acid containing dissolved $\mathrm{N}{\mathrm{O}}_2$, which makes it brown. It is obtained by distilling concentrated $\mathrm{HN}{\mathrm{O}}_3$ with a little starch, which reduces part of the acid to $\mathrm{N}{\mathrm{O}}_2$.

9.2 Physical Properties and Structure

• Colourless, fuming liquid with a pungent smell; freezes at 231.4 K and boils at 355.6 K; miscible with water.
• It turns yellow on standing because light decomposes it into $\mathrm{N}{\mathrm{O}}_2$, which dissolves in the acid. It is therefore stored in brown bottles.

Structure: in the gas phase $\mathrm{HN}{\mathrm{O}}_3$ is planar. The two terminal N-O bonds (121 pm) are shorter than the N-OH bond (140 pm), because the N=O double-bond character is shared between them by resonance.

9.3 Chemical Properties

(a) Strong acid. In water it ionises almost completely and shows all the usual properties of acids:

(b) Oxidising agent. Concentrated acid is reduced to $\mathrm{N}{\mathrm{O}}_2$, dilute acid to NO, releasing 'available oxygen':

(i) Non-metals are oxidised to their highest oxoacids (or oxide, for carbon), and the acid is reduced to $\mathrm{N}{\mathrm{O}}_2$: