The P-Block Elements (Compounds Of Group 15 Elements)

Dinitrogen

---

Dinitrogen ($N_2$) is the most abundant component of the Earth's atmosphere and plays a vital role in biological systems.

Preparation

1. Laboratory Preparation:

• From Ammonium Nitrite: Gentle heating of ammonium nitrite ($NH_4NO_2$) produces dinitrogen and water. Ammonium nitrite is unstable and is usually prepared in situ by mixing solutions of ammonium chloride ($NH_4Cl$) and sodium nitrite ($NaNO_2$).

$NH_4Cl(aq) + NaNO_2(aq) \rightarrow NH_4NO_2(aq) + NaCl(aq)$

$NH_4NO_2(s) \xrightarrow{heat} N_2(g) + 2H_2O(l)$

• From Ammonium Dichromate: Gentle heating of ammonium dichromate ($ (NH_4)_2Cr_2O_7 $) also produces dinitrogen.

$(NH_4)_2Cr_2O_7(s) \xrightarrow{heat} N_2(g) + Cr_2O_3(s) + 4H_2O(g)$

• From Metal Azides: Heating of metal azides, especially barium azide ($Ba(N_3)_2$), produces pure dinitrogen.

$2Ba(N_3)_2(s) \xrightarrow{heat} 2Ba(s) + 3N_2(g)$

2. Commercial Production:

• Fractional Distillation of Liquid Air: This is the most important industrial method. Liquid air is obtained by liquefying air and then fractionally distilling it. Since nitrogen has a lower boiling point (-196°C) than oxygen (-183°C), it distills off first.

Properties

Physical Properties:

• Colorless, odorless, and tasteless gas.
• Lighter than air (molecular weight $\approx$ 28 g/mol).
• Very low solubility in water.
• Does not support combustion.
• Liquefies at 77 K (-196°C) and solidifies at 63 K (-210°C).

Chemical Properties:

• Inertness: Dinitrogen ($N_2$) is remarkably inert at room temperature due to the presence of a strong triple bond ($N \equiv N$) between the nitrogen atoms, which requires a large amount of energy to break.
• Reactions at High Temperatures: At high temperatures, it reacts with other elements.
• With Hydrogen: Forms ammonia ($NH_3$) in the Haber process.

$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g)$ (catalyzed by Fe at high T and P)

• With Oxygen: Forms oxides of nitrogen at very high temperatures (e.g., in electric discharge or lightning).

$N_2(g) + O_2(g) \rightleftharpoons 2NO(g)$

$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$

• With Metals: Reacts with highly electropositive metals like Lithium and Magnesium at high temperatures to form nitrides.

$6Li + N_2 \rightarrow 2Li_3N$

$3Mg + N_2 \rightarrow Mg_3N_2$

• Reducing Agent: $N_2$ can act as a reducing agent in some reactions, but it is generally less reactive than hydrogen.

Ammonia

---

Ammonia ($NH_3$) is a colorless gas with a characteristic pungent odor. It is a fundamental compound of nitrogen with wide-ranging industrial and biological importance.

Preparation

1. Laboratory Preparation:

• Heating Ammonium Salts with Strong Bases: Ammonia is typically prepared in the lab by heating an ammonium salt with a strong base like sodium hydroxide ($NaOH$) or calcium hydroxide ($Ca(OH)_2$).

$NH_4Cl(s) + NaOH(s) \xrightarrow{heat} NaCl(s) + H_2O(l) + NH_3(g)$

$2NH_4Cl(s) + Ca(OH)_2(s) \xrightarrow{heat} CaCl_2(s) + 2H_2O(l) + 2NH_3(g)$

• Note: Ammonia gas is dried by passing it through quicklime ($CaO$), as $P_4O_{10}$ and conc. $H_2SO_4$ react with it.

2. Commercial Production (Haber Process):

This is the primary industrial method for producing ammonia.

Reaction:

$$N_2(g) + 3H_2(g) \rightleftharpoons 2NH_3(g) \quad (\Delta H = -46.1 \text{ kJ/mol})$$

Conditions for Haber Process:

• Temperature: Around 400-750 K (700-1000 K is also mentioned, but optimum is around 400-450°C).
• Pressure: High pressure (150-350 atm or 15-35 MPa).
• Catalyst: Finely divided iron (Fe) with promoters (like $K_2O$, $Al_2O_3$).
• Le Chatelier's Principle: The conditions are chosen based on Le Chatelier's principle. The reaction is exothermic, so lower temperatures favor product formation, but lower temperatures also slow down the reaction rate. High pressure favors the formation of ammonia as there are fewer moles of gas on the product side.

Properties

Physical Properties:

• Colorless gas.
• Characteristic pungent smell.
• Has a sharp, irritating odor.
• Lighter than air (molecular weight $\approx$ 17 g/mol).
• Can be easily liquefied to form a colorless liquid at 240 K (-33°C) under pressure, and solidified at 198.4 K (-74.7°C).
• Highly soluble in water, forming an alkaline solution due to the formation of $NH_4OH$.

$NH_3(g) + H_2O(l) \rightleftharpoons NH_4^+(aq) + OH^-(aq)$

Chemical Properties:

1. Basic Nature: Ammonia is a weak base.

• It forms ammonium ($NH_4^+$) and hydroxide ($OH^-$) ions in water.
• It reacts with acids to form ammonium salts.

$NH_3(g) + HCl(g) \rightarrow NH_4Cl(s)$ (forms white fumes)

$2NH_3(aq) + H_2SO_4(aq) \rightarrow (NH_4)_2SO_4(aq)$

• Reaction with metallic salts: Forms precipitates of metal hydroxides if the metal hydroxide is insoluble.

$CuSO_4(aq) + 2NH_3(aq) + 2H_2O(l) \rightarrow Cu(OH)_2(s) + (NH_4)_2SO_4(aq)$

With excess $NH_3$: Forms a deep blue complex (tetramminecopper(II) ion).

$Cu(NH_3)_4^{2+}(aq) + 4H_2O(l)$

2. Reducing Nature: Ammonia acts as a reducing agent, especially at high temperatures, where it is oxidized to dinitrogen.

• With Metal Oxides: Reduces oxides of less reactive metals to metals.

$2NH_3 + 3CuO(s) \xrightarrow{heat} 3Cu(s) + N_2(g) + 3H_2O(g)$

• With $O_2$ (Catalytic Oxidation): In the presence of platinum or platinum-rhodium gauze at 500 K, ammonia is oxidized to nitric oxide ($NO$).

$4NH_3(g) + 5O_2(g) \xrightarrow{Pt, 500K} 4NO(g) + 6H_2O(g)$

3. Formation of Amides and Nitrides:

• Reacts with highly electropositive metals (like alkali metals) to form amides ($MNH_2$) and nitrides ($M_3N_2$).

$2Na(l) + 2NH_3(l) \rightarrow 2NaNH_2(l) + H_2(g)$

4. Complex Formation: Acts as a Lewis base and forms complex compounds with many metal ions.

Uses:

• Manufacture of nitric acid.
• Used in the production of ammonium salts (fertilizers).
• Used in refrigeration plants.
• Used as a cleaning agent.
• Used in the Haber process for ammonia synthesis.

Oxides Of Nitrogen

---

Nitrogen forms several oxides, which exhibit a range of oxidation states for nitrogen and have varying properties.

Common Oxides and Their Properties:

1. Dinitrogen Oxide ($N_2O$) (Nitrous Oxide):
• Preparation: Gentle heating of ammonium nitrate ($NH_4NO_3$).

$NH_4NO_3(s) \xrightarrow{heat} N_2O(g) + 2H_2O(g)$

• Properties: Colorless gas, sweetish odor, neutral, does not support combustion but relights a glowing splint.
• Uses: Used as an anesthetic ("laughing gas") and as a propellant in aerosol cans.
2. Nitric Oxide ($NO$):
• Preparation: By direct combination of $N_2$ and $O_2$ at very high temperatures (e.g., electric arc) or by reduction of nitric acid with copper.

$N_2(g) + O_2(g) \xrightarrow{high \ T} 2NO(g)$

$3Cu(s) + 8HNO_3(dilute) \rightarrow 3Cu(NO_3)_2(aq) + 2NO(g) + 4H_2O(l)$

• Properties: Colorless gas, readily oxidizes in air to reddish-brown $NO_2$.

$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$

• Uses: Industrial production of nitric acid.
3. Dinitrogen Trioxide ($N_2O_3$):
• Preparation: By mixing equal volumes of $NO$ and $NO_2$ at low temperatures.
• Properties: Colorless liquid (at low temp), acidic, decomposes into $NO$ and $NO_2$ on warming.
4. Dinitrogen Tetroxide ($N_2O_4$):
• Preparation: Dimerization of $NO_2$ at low temperatures.

$2NO_2(g) \rightleftharpoons N_2O_4(g)$

• Properties: Colorless gas, decomposes into reddish-brown $NO_2$ gas upon heating. It is an equilibrium mixture of $NO_2$ and $N_2O_4$.
5. Dinitrogen Pentoxide ($N_2O_5$):
• Preparation: By the dehydration of nitric acid with $P_4O_{10}$.

$4HNO_3 + P_4O_{10} \rightarrow 4HPO_3 + 2N_2O_5$

• Properties: White crystalline solid, acidic, decomposes easily. It is the anhydride of nitric acid.

Acidic Nature: Oxides of nitrogen like $N_2O_3$, $NO_2$, $N_2O_4$, and $N_2O_5$ are acidic and react with water to form corresponding acids (nitrous acid or nitric acid).

$N_2O_3 + H_2O \rightarrow 2HNO_2$

$N_2O_4 + H_2O \rightarrow HNO_3 + HNO_2$

$N_2O_5 + H_2O \rightarrow 2HNO_3$

NOx Pollution: Oxides of nitrogen ($NO$ and $NO_2$) are major air pollutants formed during combustion processes in engines and industrial furnaces at high temperatures. They contribute to acid rain and smog formation.

Nitric Acid

---

Nitric acid ($HNO_3$) is a highly corrosive and strong mineral acid.

Preparation

1. Laboratory Preparation:

• By heating a nitrate salt (like $KNO_3$ or $NaNO_3$) with concentrated sulfuric acid.

$NaNO_3(s) + H_2SO_4(conc.) \xrightarrow{heat} NaHSO_4(s) + HNO_3(g)$

• The nitric acid vapor is condensed and collected.

2. Commercial Production (Ostwald Process): This is the most important industrial method.

Steps:

1. Catalytic Oxidation of Ammonia: Ammonia gas is catalytically oxidized with oxygen in the presence of a platinum-rhodium gauze catalyst at about 500 K.

$4NH_3(g) + 5O_2(g) \xrightarrow{Pt/Rh \ gauze, 500K} 4NO(g) + 6H_2O(g)$

2. Oxidation of Nitric Oxide: Nitric oxide ($NO$) formed is then cooled in air to form nitrogen dioxide ($NO_2$).

$2NO(g) + O_2(g) \rightarrow 2NO_2(g)$

3. Absorption in Water: Nitrogen dioxide ($NO_2$) is absorbed in water in the presence of oxygen to form nitric acid.

$3NO_2(g) + H_2O(l) \rightarrow 2HNO_3(aq) + NO(g)$

The $NO$ formed in this step is recycled back into step 2.

Concentration: The nitric acid obtained from the Ostwald process is about 68% pure. It can be concentrated further by dehydration with concentrated sulfuric acid or by distillation over boiling sulfuric acid to obtain anhydrous or fuming nitric acid.

Properties

Physical Properties:

• Pure nitric acid is a colorless liquid.
• Commercial samples are often yellowish due to the presence of dissolved $NO_2$.
• It has a pungent odor.
• It is a strong acid and corrosive.
• It is highly soluble in water.

Chemical Properties:

1. Acidic Nature: Nitric acid is a strong monoprotic acid, completely ionizing in water to produce hydronium and nitrate ions.

$HNO_3(aq) + H_2O(l) \rightarrow H_3O^+(aq) + NO_3^-(aq)$

2. Oxidizing Action: Nitric acid is a very strong oxidizing agent, especially when hot and concentrated. It gets reduced to various products depending on the reducing agent and reaction conditions.

• With Metals:
• Reacts with most metals (except noble metals like Au, Pt) to form nitrates and $NO$, $NO_2$, or $N_2O$, $N_2$, $NH_4^+$ depending on the concentration of the acid and the reactivity of the metal.

With $Cu$ (dilute acid): $3Cu(s) + 8HNO_3(dilute) \rightarrow 3Cu(NO_3)_2(aq) + 2NO(g) + 4H_2O(l)$

With $Cu$ (concentrated acid): $Cu(s) + 4HNO_3(conc.) \rightarrow Cu(NO_3)_2(aq) + 2NO_2(g) + 2H_2O(l)$

• With very dilute nitric acid, metals like Zn can reduce $NO_3^-$ to $NH_4^+$.

$4Zn(s) + 10HNO_3(very \ dilute) \rightarrow 4Zn(NO_3)_2(aq) + NH_4NO_3(aq) + 3H_2O(l)$

• Does not react with noble metals like Gold and Platinum.
• With Non-metals:
• Oxidizes non-metals like Iodine, Sulfur, Phosphorus.

$I_2(s) + 10HNO_3(conc.) \rightarrow 2HIO_3(aq) + 10NO_2(g) + 4H_2O(l)$

$S(s) + 6HNO_3(conc.) \rightarrow H_2SO_4(aq) + 6NO_2(g) + 2H_2O(l)$

$P_4(s) + 10HNO_3(conc.) \rightarrow 4H_3PO_4(aq) + 10NO_2(g) + 2H_2O(l)$

• With Oxidizing Agents: It is itself oxidized by very strong oxidizing agents like ozone.

3. Formation of Aqua Regia:

• A mixture of concentrated nitric acid and concentrated hydrochloric acid in the molar ratio 1:3.
• It is one of the few reagents that can dissolve noble metals like gold ($Au$) and platinum ($Pt$).

$Au + 3HNO_3 + 4HCl \rightarrow H[AuCl_4] + 3NO_2 + 2H_2O$

$Pt + 4HNO_3 + 6HCl \rightarrow H_2[PtCl_6] + 4NO_2 + 4H_2O$

• The oxidizing power comes from the formation of highly reactive species like nitrosyl chloride ($NOCl$) and free chlorine ($Cl_2$).

Uses:

• Manufacture of fertilizers (ammonium nitrate).
• Manufacture of explosives (like nitroglycerin, TNT).
• In metallurgy for dissolving noble metals and refining metals.
• In engraving, etching, and tanning industries.
• As an oxidizing agent in various chemical processes.