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6.6 Group 15 (Pnictogens): N, P, As, Sb, Bi

Textbook: Chapter 15; Chapter 8, Section 8.5

• Trivalent or pentavalent. Bismuth prefers oxidation state +3. Nitrogen can achieve a (-)ve

oxidation state (e.g., N3–)

• Nitrogen and phosphorus are non-metals, with phosphorus existing in at least three different

allotropes (white, red and black). • Solid white phosphorus has an interesting structure shown

at right. It consists of highly strained P4 tetrahedra that explain its high reactivity – it is pyrophoric!

• Arsenic and antimony both exist in non-metallic and semi-

metallic modifications, while bismuth only exists in a semi-metallic form.

• The semi-metallic forms all conduct electricity, but are not ductile.

Typical compounds: • Hydrides: EH3 (Ammonia, Phosphine, Arsine, Stibine, Bismutane); N2H4 (hydrazine, a

rocket propellant). N2H4 + O2 → N2 + 2H2O ΔH° = -622 kJ mol-1

Liquid ammonia as a non-aqueous solvent Liquid NH3 is a colourless fluid; 0.70 density of water and only 0.25 as viscous; b.p. = - 33o C and ΔHvap ≈ 5 kcal/mol, which is very high. Liquid can be handled in lab by cooling with dry-ice. NH3 vs. H2O: NH4

+ is a weaker hydrogen-bonding acid than H30+, but NH3 is a stronger H-bonding base than H2O (N has lower EN than O). µ(H2O) = 1.82 Db µ(NH3) = 1.47 Db

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Solubilities:

H2O: Dominated by dipole interactions; van der Waals & polarizability less important

→ non-specific solvation of species with high polarity (large dipole moment)

NH3: Dipole interactions as important as polarizability and van der Waals interactions

→ better solvation of polarizable solutes (i.e., large electron “cloud”)

Some data:

Source: Purcell + Kotz, Inorganic Chemistry, Holt-Saunders, 1977.

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• Oxides: N2O (nitrous oxide – laughing gas); NO (nitric oxide); NO2 (nitrogen dioxide – brown smog); P2O5 (or P4O10; drying agent); P2O7

4- (diphosphate ion; e.g., phosphate linkage in ADP3-); and corresponding acids HNO2, HNO3, H3PO4.

The structure of P2O5 (left) is derived from P4 tetrahedra by oxidizing each P-P bond and each corner.

• Halides: most nitrogen halides are unstable (e.g., NI3·NH3 is explosive); ECl3 (E = P, As, Sb)

PCl5(g) is actually a salt in the solid state [PCl4+][PCl6

-]; SbF5 (strong Lewis acid used to produce some of the strongest Brønsted acids → superacids)

SbF5 itself is a tetrameric liquid, which reflects its Lewis acidity:

Forms a superacid in HF:

SbF5(l) + HF(l) → H2F+(sol) + SbF6-(sol)

Also forms a superacid in HSO3F (see next page…)

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Lit.: R.J. Gillespie, Acc. Chem. Res., 1, (1968), p. 202-209 Fluorosulfuric acid as a solvent…

m.p. = -89 oC b.p. = 163 oC ρ = 1.73 g/mL (25 oC) Kion = 10-8

Note the presence of only highly electronegative atoms on the sulfur. Their electron

withdrawing power makes the molecule so acidic.

Can be handled in glass if HF free.

Useful liquid temperature range.

Stronger than oleum (= fuming sulfuric acid) - protonates just about everything …

HF is not an acid in HSO3F.

When SbF5 is dissolved in HSO3F a superacid is formed:

Structures of the anion components:

Superacids are powerful enough to generate carbocations the foremost example of which is

the t-butyl cation, first generated in the 1950’s by George Olah in a Dow Chemical Lab in Sarnia, Ontario (!!!) - in 1994 George Olah received the Nobel Prize in Chemistry for this work.

The structure of Me3C+ and CH5+ (generated analogously):

Both cations have been characterized by NMR at low temperature. Carbocations had been postulated as transient intermediates in SN1 reactions - Olah's work proved their existence.

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• Nitrides: Saline nitrides (e.g., Li3N, Mg3N2); Covalent nitrides (e.g., boron nitride BN; sulfur

nitrides such as S2N2, S4N4 and (SN)x; cyanogen (CN)2; phosphazines based on R2PN unit).

Polyphosphazines can be rubbery like silicones and can be used as biodegradable supports for

medical applications such as bone regeneration.

31P NMR (and 19F NMR)

• Many nuclei other than 1H and 13C can be used for NMR analysis. • Typical NMR-active nuclei have a nuclear spin I = ½.

(If I = 0 there will be no signal. If I > ½, the nuclear quadrupole moment, a non-uniform distribution of electric charge, serves to broaden the signal.)

• 31P has I = ½ and a natural abundance of 100%. • 19F has I = ½ and a natural abundance of 100%. • As with 1H NMR, chemical shifts give information

about chemical environment, homonuclear coupling (coupling between two nuclei of the same element) and heteronuclear coupling (coupling between two nuclei of two element types, e.g., 1H – 31P coupling) can occur and can give information regarding structure.

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Source: C. E. Housecroft and A. G. Sharpe “Inorganic Chemistry, 2nd Ed.”, Pearson Prentice-Hall, 2005.

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Source: H. Günther “NMR Spectroscopy, 2nd Ed.”, Wiley, 1992. Related Textbook Exercises 8.9-8.11, 15.1, 15.5 – 15.9, 15.14, 15.16, 15.17

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6.7 Group 16 (Chalcogens): O, S, Se, Te, Po* (from the Greek chalcos = ore; genna = to generate)

Textbook: Chapter 16, Intro & Sections 16.1 – 16.9, 16.12 – 16.16

• Di-, tetra- and hexa-valent. • Oxygen is almost always in oxidation state –2, unless bonded to fluorine or to itself…e.g.,

• Sulfur exists in many different allotropes – the most stable is the familiar yellow solid

consisting of S8 crowns: • Catenation is common for S, Se, and Te. Typical Compounds: • Hydrides: H2O (water), H2O2 (hydrogen peroxide; liquid); H2S, H2Se, H2Te (foul-smelling,

toxic gases). • Halides: HOCl (hypochlorous acid), NaOCl (sodium hypochlorite; household bleach), OF2;

S2Cl2 (sulfur monochloride), SCl2 (sulfur dichloride), SF4, SF6; Se2Cl2, SeCl4, SeF6; Te2I, etc. • Oxides: SO2 (can be used as a solvent under pressure, like ammonia), SO3 and their oxoacids

H2SO3 and H2SO4; SOCl2 (thionyl chloride; precursor to mustard gas, a chemical warfare agent); TeO2 (tellurite, naturally occuring mineral).

SOCl2 is used for the manufacture of mustard gas – last used by Saddam Hussein in Kurdistan.

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In general the chalcogen oxides form strong acids, when reacting with water, e.g.: SO2 + n H2O → H2SO3 + n-1 H2 O → H3O+ + HSO3

- + n-2 H2O … recall our previous discussion of oxoacids and oxobases! Related Textbook Exercises 16.1 – 16.5, 16.7 – 16.9; Problem 16.1 6.8 Group 17 (Halogens): F, Cl, Br, I, At*

(from the Greek halos = salt; genna = to generate) Textbook: Chapter 17 • Monovalent, (for X ≠ F also tri-, tetra-, penta- and hepta-valent.) Usually oxidation state –1.

Other oxidation states occur in the oxygen acids of Cl, Br and I, in NX3 (X ≠ F) and in interhalogen compounds. (F is always oxidation state –1, by definition.) This is a consequence of the electronegativity order F > O > Cl > Br > I

• Elemental halogens X2 are very reactive, oxidizing and toxic – with decreasing reactivity

going down the group (why this trend?):

Halogen F-F Cl-Cl Br-Br I-I At-At

d [Å] 1.42 1.98 2.28 2.66 2.80

BE [kJ/mol] 158.1 243.52 192.99 151.34 95.0

As with the other groups, the lightest element (fluorine) shows anomalous behaviour.

Fluorine anomaly Given the trend (above table), we would predict F-F bond enthalpy to be ca. 300 kJ/mol. From covalent radii with other compounds we’d predict F-F bond distance to be ca. 1.28 Å. BUT due to the small atomic radius of F, repulsion between non-bonding electron pairs

decreases the BE and increases the d of the F-F bond. An MO picture shows poor overlap in the formation of bonding σ-type MO’s but reasonable overlap to form antibonding π* MO’s.

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• Cl2 has been used as a chemical warfare agent - by Fritz Haber, German nationalist jew (no in 1916 that was not a contradiction!) in the Imperial German Army in WW I and of course the same guy who invented the Haber-Bosch process.

Typical compounds: • Salts: Combining any of the halogens with any metal gives ionogenic materials, i.e. salts –

consequence of their high electronegativity and electron affinity: NaCl, KBr, AgCl, … • Protic acids: HF, HCl, HBr, HI HF differs significantly from the other HX acids. HF is a liquid (the others are gases). This is due to strong H-F hydrogen bonding. HF is a weak acid in aqueous solution. HF will dissolve glass! HF will absorb through skin and cause necrosis of the skin and deep tissue as well as

decalcification of the bone via formation of CaF2. FIRST AID: An HF burn must be treated immediately by injecting a calcium gluconate solution or gel into the tissue of the affected area. • Covalent compounds with p-block elements:

Group 13: BF3, AlCl3, GaCl3, TlI

Group 14: CHCl3 (chloroform), (-CF2-)n (polytetrafluoroethene, PTFE), H3CI (methyl iodide,

alkylating agent), SiCl4, (SnFMe3)n

Group 15: PCl5, AgPF6, NOSbF6

Group 16: SF4, SF6, S2Cl2, SeCl4, TeCl2

Low polarizability of fluorine atoms leads to high volatility of fluorinated materials.

Teflon® is “non-stick” due to weak dispersion interactions of the fluorine “sheath” around the carbon polymer backbone. • Interhalogens: ClF, ICl, IBr, ClF3, IF7, BrF5,

ICl3 (a solid source of Cl2)

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• Polyhalides: I3+, I5

+, I5-, I4

2-, Cl3-, Br2

+, but most common is I3-

• Oxides: O2F2, OF2 (oxygen fluorides); extensive Cl (and Br) oxide chemistry is known

giving rise to compounds in which the oxidation state of chlorine ranges from +1 to +7: +1 +2 +3 +4 +5 +6 +7

neutral Cl2O

Chlorine dioxide

Cl2O3 ClO2 Cl2O6 Cl2O7

anionic ClO-

Hypochlorite (bleach)

ClO2-

Chlorite

ClO3-

Chlorate

ClO4-

Perchlorate

Related Textbook Exercises 17.1, 17.4 – 17.10, 17.12-17.22, 17.26

6.9 Group 18 (Noble Gases): He, Ne, Ar, Kr, Xe, Rn*

Textbook: Chapter 18

• All noble gases have a filled octet shell and are therefore practically inert – hence “noble”. • Preferred oxidation state is zero. Only xenon has a developed chemistry, mostly involving

oxygen and fluorine compounds. • A few krypton and radon fluorides are also known. • Helium was first identified by its solar emission spectrum. The sun burns hydrogen into

helium by nuclear fusion: 4 1H + 2 e --> 4He + 2 neutrinos + 6 photons + 26 MeV

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• Main use of noble gases (Ar, Ne, Kr) is in fluorescent lights (“neon tubes”): Current is passed through a mixture of Hg vapor and a noble gas, ionized by a strong electric field. This excites the atoms to higher energy states. Relaxation occurs with emission of UV light. A phosphorescent coating in the tube then “translates” the UV into visible light with some heat loss (cf. Yablonski Term Scheme, CHEM 207).

• Other uses: cryo-technology (He(l) to cool superconducting magnets; NMR!), deep-sea

diving (He), ballons (He), welding and inert gas technology (Ar; e.g., Preuss lab glove box). • Helium is the closest thing to an ideal gas and becomes superfluid below 4.2 K.

Typical compounds of Xe

• Fluorides: XeF2, XeF4, XeF6. These are strong oxidizing agents. They react with strong Lewis acids (e.g., XeF2 + SbF5 → [XeF+][SbF6

-]) and form complexes with Lewis base F-; XeF5

-, XeF7-, XeF8

2-.

• Oxides and oxofluorides: XeO3, XeOF4, H2XeO6. Xenon oxides are endoergic (ΔfGθ > 0) and usually highly explosive.

• Organoxenon compounds: [C6F5Xe]+, Xe(C6F5)2. First Xe-C bond reported in 1989.

Related Exercises 18.1, 18.2, 18.4 – 18.10