Lecture 4. The Chemistry of Nitrogen

1. Electronic Configuration

[He]2s²2p³ => Accessible oxidation states -3 +5.

- More valence electrons than orbitals.

- Lone pairs, play important role in chemistry.

- Aqueous redox chemistry.

- High oxidation states oxidising.

2. Structure of the Element

- m.p. -210 deg.C, b.p. -196 deg.C

- ¹⁴N 99.54% I = 1; ¹⁵N 0.46% I = 1/2

- 78% by volume of Earth's atmosphere.

- Essential for life.

For Nitrogen:

- Compare weak N-N single bond, due to non-bonded electron repulsion.

- In contrast P^{[[equivalence]]}P weak, P-P stronger then N-N.

- Catenated N compounds rare, P-P probably OK.

- Very strong N[[equivalence]]N bond, N-E bonds commonly weak.

- Many nitrogen compounds are thermally unstable with respect to N2.

- With lone pair repulsion, nitrogen tries to use multiple bonds (NO, S4N4).

Table of some relevant bond energies (kJmol^-l)

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N-H 386 N-C 305 N-F 283

N-N 158 N-O 201 N-Cl 190

N-P -200

N^{[[equivalence]]}N 945 N=N 418 N-N 158

P^{[[equivalence]]}P 481 P-P 198

N=O 607 N^{[[equivalence]]}C 887 N=C 615

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3. Nitrogen as a Ligand

Like CO, N2 is a [[pi]]-acid

RuCl3 + N2H4 [Ru(NH3)5(N2)]²⁺ (contains Ru-N^{[[equivalence]]}N)

Nitrogen Fixation

N2 + 8H⁺ + 8e + 16 ATP 2NH3 + H2 + 16ADP + Pyrophosphate2

unknown mechanism

Enzyme nitrogenase contains the transition metals Mo, Fe, V; The N2 binding site uncertain. For many yearsthe molybdenum site thought to be the active centre, recent crystal structure shows coordinatively saturated Mo; now Fe likely active centre.

4. Nitrogen Halides

Structure

NX3: pyramidal

However: N(SiMe3)3 is planar due to d[[pi]]-p[[pi]] bonding

N2X4, NX2

Max coordination number = 4, eg. NF4⁺

5. Nitrides

- Many interstitial nitrides, chemistry not well developed.

- Nitrogen atoms in holes in metal structures.

- N N^3-

Requires lots of energy => need small cation to stablise structure.

(Madelung Energy)

Hence, the only Group 1 or 2 nitrides are Li3N, and Mg3N2

Figure 1: The layer structure of Li3N, more correctly formulated as Li[Li2N], has hexagonal Li6 nets. The nitrogen has hexagonal bipyramidal coordination, Li3N has high electrical conductivity.

6. Nitrogen Hydrides

NH3, N2H4, N2H2

- NH3 only thermally stable hydride.

- Prepared industrially by the Haber process (500deg.C, high pressure),

naturally by nitrogen fixation.

Uses of Ammonia

4NH3 + 5O2 4NO + 4H2O (Huge scale)

NH3 + NaOCl NH2Cl + NaOH

Chloramine

NH3 + NH2Cl N2H4 + NaCl + H2O (Raschig process)

Side reaction:

2NH2Cl + N2H4 N2 + HN4Cl

catalysed by traces (ppm) of metal ions, add gelatin.

Structure

Acid/Base properties

NH3 basic N2H4, weakly basic

NH4⁺ N2H5⁺, N2H6²⁺

Redox properties

- Both NH3 and N2H4 mild reducing agent (ox states -3 and -2).

- Reactions complex in aq solution

Summary of reactivity

7. Nitrogen Halides

NH4Cl + HOCl H2O + HCl + NH2Cl

NH2Cl + HOCl HNCl2 +H2O NCl3

Thermally unstable, NCl3, yellow oil, explosive

NH3 + F2 NF3 + N2F4 + N2F2

NF3 only thermaly stable NX3 compound

- Weaker base than NH3

- Weaker donor the NH3

NF3 + F2 + AsF5 NF4⁺AsF6^-

8. Nitrogen Oxides

- All oxides of nitrogen thermodynamically unstable with respect to N2 and O2:

=> Kinetic stability.

So: N2 + O2 no reaction

Nitrous Oxide, Dinitrogen Monoxide (N2O)

Synthesis

NH4NO3 N2O + H2O

Structure

Nitric Oxide, Nitrogen Monoxide (NO)

Synthesis

NaNO2 + H2SO4 NO

Cu + HNO3 NO

Structure

In solid, dimer

[[pi]]*/[[pi]]* overlap, giving long and weak N-N bond, diamagnetic

Gas phase reaction: O2 + NO

3rd order reaction

Like CO, can act a [[pi]]-acid, many complexes known.

Reactivity

Dinitrogen Trioxide N2O3

Synthesis

NO + N2O4 N2O3

N2O3 readily dissociates:

N2O3 NO + NO2

Structure

The symmetric and asymmetric isomer can be readily interconverted

Like (NO)2 dimer, long and weak N-N bond

Nitrogen Dioxide NO2 Dinitrogen Tetroxide N2O4

Synthesis

Pb(NO3)2 PbO + 2NO2 + 1/2O2

Structure

- Liq NO2, can get anhydrous nitrates

Cu + 3N2O4 Cu(NO3)2.(N2O4) + 2NO

Dinitrogen Pentoxide N2O5

Synthesis

HNO3 + P2O5 N2O5

Structure

- NO2⁺NO3^- in solid or acids such as H2SO4

Aqueous redox chemistry: Oxidation State (Frost) diagrams.

1. For the reaction of element X in formal oxidation state n,

X(n) + ne^- X(0) Edeg. [[equivalence]] -[[Delta]]Gdeg. [[equivalence]] nFEdeg.

If -[[Delta]]Gdeg./F is plotted against the formal oxidation states, the higher the value the less stable, ie. more oxidizing, the species will be. For nitrogen at aH+ = 1 we have:

NO3^- + 6H⁺ + 5e^- 1/2N2 + 3H2O Edeg. = 1.25V

1/2N2O4 + 4H⁺ + 4e^- 1/2N2 + 2H2O Edeg. = 1.36V

HNO2 + 3H⁺ + 3e^- 1/2N2 + 2H2O Edeg. = 1.45V

NO + 2H⁺ + 2e^- 1/2N2 + H2O Edeg. = 1.68V

1/2N2O + H⁺ + e^- 1/2N2 + 1/2H2O Edeg. = 1.77V

1/2N2 + 2H⁺ H2O + e^- NH3OH⁺ Edeg. = -1.87V

1/2N2 + 5/2H⁺ + 2e^- 1/2N2H5⁺ Edeg. = -0.23V

1/2N2 + 4H⁺ + 3e^- NH4⁺ Edeg. = 0.27V

2. With -[[Delta]]Gdeg./F in units of volts and defining -[[Delta]]Gdeg. as per mol of nitrogen atoms, the coordinates will be:

N(V): NO3^- (5x1.25, 5)

N(IV): N2O4 (4x1.36, 4)

N(III): HNO2 (3x1.35, 3)

N(II): NO (2x1.68, 2)

N(1): N2O (1x1.77, 1)

N(-I): NH3OH⁺ [-1x(-1.87), -1]

N(-II): N2H5⁺ [-2x(-0.23), -2]

N(-III): NH4⁺ (-3x0.27, -3)

3. A plot of -[[Delta]]Gdeg./F against the formal oxidation state is called an oxidation state disgram -- very useful for comparative chemistry in aqueous solution.

Figure 2: The oxidation state diagram of nitrogen and phosphorus in acidic conditions.

Comparision of Nitrogen and Phosphorus

Three major features:

- the radius of phosphorus is 50% greater than that of nitrogen

- the ionization energies decrease from N to P

- the bond energy trends are different

1. The oxidation states for both elements vary from +5 to -3,

but the stability trends are different.

2. Valence shell expansion: the coordination number of nitorgen

is <= 4, but phosphorus is larger and the C.N. can be up to 6.

PF6^- and PCl6^- but NH4⁺ and NF4⁺.

3. Bond enthalpies:

N-H vs P-H chemistry

The N-H bond is stronger than the P-H bond (386 vs 321 kJmol^-l), so despite a lower [[Delta]]Hdeg.(atm) for P, PH3 is thermodynamically unstable. PH3 is pyrophoric (kinetics). NH3 also a stronger base.

Lone-pair repulsion

Nitrogen:

Weakens N-X bonds in many nitrogen compounds resulting in thermodynamic instability, e.g., the oxides and halides.

Phosphorus:

Much reduced in P-X bonds (also d[[pi]]-p[[pi]] bonding, see below)

hence stronger bonds and thermodynamically stable halides and oxides.

Compare:

Hydrolysis of NCl3 gives NH3 and Cl2O

whereas PCl3 gives H3PO3 and HCl.

p[[pi]]-p[[pi]] bonding

Nitrogen:

Very important feature of nitrogen chemistry. Coupled with weak N-X single bonds, nitrogen compounds tend to be monomeric or dimeric, e.g., elemental nitrogen (N2), oxides, nitrates and nitrites.

Phosphorus:

Quite strong multiple bonds but single bonds even more important. Hence many phosphorus compounds "polymerize", e.g., the oxides (N2O3 and N2O5 vs P4O6 and P4O10), oxy-anions, suphides [note: NO is monomeric but N4S4 and P4S3, P4S4 etc], elemental phosphorus (P4), and nitrogen compounds [e.g., (NPCl2)4].

d[[pi]]-p[[pi]] bondinq

Open only to phosphorus, and particularly in compounds with the more electronegative elements N, O, F. Strengthens both P-X and P=X bonds and hence chain or cyclic structures for the oxides and oxy-anion salts. Kinetic effects too:

X3N->O X3P=O

Normal N-O bond length; Short P=O bond;

high polarity; low energy; low polarity; high energy;

high reactivity inert

4. Oxidation states

- Negative oxidation state stability N > P

- Positive oxidation states P > N

- Oxidation states: NCl3 very unstable but PCl3 stable;

N2O5, NO3 , NF4 but P4O10, PF5, PCl5, PBr5; nitric acid

is very oxidizing but phosphoric acid is not.

- In aqueous solution: refer to oxidation state diagram.

Note: H3PO2 is in fact H2P(=O)OH, and phosphorous acid is

not P(OH)3 but HP(=O)(OH)2 - an indication of the P=O

bond strength

- Donor properties of NR3 and PR3 in complex formation.

Supplementary Material

Liquid ammonia: a non-aqueous solvent

1. Despite low boiling point (-33.4 deg.C), easy to handle

2. Solubilities:

- relatively high dielectric constant (ammonia, [[epsilon]]0 = 26.7 @ -60 deg.C;

water, [[epsilon]]o = 82 @ 18 deg.C).

=> ionic compounds can be soluble but the lower [[epsilon]]0 compared to water means that salt with highly charged, non-polarisable anions such as carbonates, sulphates, and phosphates are insoluble.

- NH3 is more polarisable than H2O, so salts with

more polarisable anions are more soluble, hence the

solubility trends.

F^- < Cl^- < Br^- < I^-

PO4^3- < SO4^2- < OAc^- < NO3

- specific solvation: NH3 is a better a-donor than H2O, and ammine complexes are formed, especially with the later transition (Ni²⁺, Cu²⁺) and B metals (Ag⁺, Zn²⁺). Hence higher solubilities for compounds of these metals than those of the A-metals.

3. Self-ionization of ammonia

2NH3 NH4⁺ + NH2^- K223K ca. 10^-30

is much "weaker" than water. Liquid ammonia will therefore tolerate very strong bases such as C5H5^- which would otherwise be hydrolysed in water.

4. Ammonia is kinetically stabilized to reduction (but easily oxidized) by many reagents, e.g., the reaction;

Na + NH3 NaNH2 + H2(g)