16.4: The Elements

Introduction

• Oxygen is another very ubiquitous (found everywhere) element. Compounds are known with every element except helium, neon and krypton.
• Oxygen forms ionic oxides containing O^2– with metals.
• Oxygen is frequently bi-covalent, X-O-Y or X=O.
• O^– will form one covalent bond, eg OH^– and O⁺ can form three, eg H₃O⁺. Vary rarely, tetracovalent O²⁺ is found. An example is Be₄O(ac)₆ where the O²⁺ is at the centre of a tetrahedron of Be atoms.

Ionic Oxides

• The formation of an ionic oxide requires a high lattice energy, and low ionization potential of the cation to overcome the highly endothermic O=O bond dissociation energy, 496 kJ mol^-1, and electron attachment enthalpy for 2 electrons, 752 kJ mol^-1 in total.

Covalent Oxides

• The compounds with non-metals are covalent and can be molecular, e.g. CO₂ or network structures, e.g. SiO₂.
• The same two examples demonstrate the ability of oxygen to be involved in pp - pp bonding or pp - dp bonding. (The Si-O-Si bonds in SiO₂ are nearly linear.)

Acid-Base Properties of Oxides

• The oxygen anions, oxide, O^2–, superoxide, O₂^–, and peroxide, O₂^2– are hydrolysed istantly by water:

  O^2– + H₂O 2OH^–

  O₂^– + H₂O O₂ + HO₂^– + OH^–

  O₂^2– + H₂O HO₂^– + OH^–

• Acidic Oxides: The non-metal oxides in this category are sometimes called "acid anyhdrides":

  N₂O₅ + H₂O 2HNO₃

  SO₃ + H₂O H₂SO₄

  CO₂ + H₂O H₂CO₃

• Sometime neutral water is not enough and a base is needed:

  Sb₂O₅ + 2OH^– + 5H₂O 2Sb(OH)₆^–

• Sometimes an acidic oxide is needed:

  Na₂O + SiO₂ Na₂SiO₃

  Some oxides are amphoteric and others are more inclined to undergo redox reactions than acid base reactions if they are not inert, e.g. MnO₂.

Occurence, Isolation and Allotropy

• There are three natural isotopes of oxygen of relative abundance: ¹⁶O = 99.759%, ¹⁷O = 0.0374% and ¹⁸O = 0.2039%

  The heavier ones can be enriched by fractionation of liquid oxygen and they are used in tracer studies.

• There are two molecular forms (allotropes), dioxygen, O₂ the common form, and trioxygen, ozone, O₃.
• Dioxygen is a paramagnetic molcule with 2 unpaired electrons. Molecular orbital theory provides a ready explanation for this: valence bond theory does not.
• Ozone is a paramagnetic bent molecule formed when dioxygen is passed through an electric discharge or irradiated with ultra-violet light. Its sharp odour can be smelled aroung electric motors and some older photcopiers. The liquid is dark blue. Liquid dioxygen is pale blue.

  Chemical Properties of Oxygen and Ozone

  • Ozone is the more oxidizing of the two.
  • Ozone can be measured by the reaction:

    O₃ + KI + H₂O I₂ + 2KOH + O₂

    The iodine is titrated with thiosulphate.

  • Ozone is used in place of chlorine for water treatment.
  • Oxygen is a common contaminent in organic and aqueous solvents, and it can be very difficult to remove.

Hydrogen Peroxide

• Hydrogen peroxide, H₂O₂, is a colorless liquid boiling at 152.1 ^oC.
• It is ~40% more dense than water, and like water it is extensively hydrogen bonded. When pure it can decompose explosively. The most stable rotational conformation has a tortion angle of 96.5^o - see Figure18-1.
• It is unstable with respect to water and dioxygen by -97 kJ mol^-1.
• It is somewhat more acidic than water:

  H₂O₂ H⁺ + HO₂^– k = 1.5x10^-12

• There are two principle methods of synthesis:
  1. 

  2. 2HSO₄^– HO(O)₂S-O-O-S(O)₂OH + 2e^– (by electrolysis)

     HO(O)₂S-O-O-S(O)₂OH + H₂O H-O-O-S(O)₂OH + H₂SO₄ (fast)

     H-O-O-S(O)₂OH + H₂O H₂SO₄ + H₂O₂ (slow)

     The H₂O₂ can be obtained 90-98% pure by distillation.

• Hydrogen peroxide is a good oxidizing agent, fast in base and slower in acid and often acts via a free radical mechanism.

The Peroxides and Superoxides

• Ionic peroxides are formed by the alkali and alkaline earth elements (not Be). They are also oxidizing agents.
• Paramagnetic ionic superoxides are formed by potassium, rubidium and cesium. They are very powerful oxidizing agents.
• Other covalent peroxides include the peroxo acids, e.g. peroxosulphuric and peroxodisulphuric acid, the intermediates in the production of H₂O₂ shown above, and the dangerously explosive organic peroxides ROOR which form by free radical oxidation of ethers. (They are the reason why ethers shold never be distilled to dryness unless the peroxides have been destroyed immediately beforehand by washing with with FeSO₄ or passing the ether over activated alumina. A solution of Fe²⁺ and SCN^– is used to test for their presence - look for the formation of the blood red [Fe^III(SCN)]²⁺ ion.)

The Dioxygenyl Radical (O₂⁺)

See notes on Chapter 21.

Dioxygen as a Ligand

Read this section - oxygen binding to transition metals is of interest because of its binding by haemoglobin and myoglobin. Note the structural types.

Oxygen Compounds as Ligands

More transition metal chemistry. Skip it for now.

Oxygen Florides

See notes on Chapter 20.

Chemistry 242 - Inorganic Chemistry II
Chapter 19 - Sulphur, Selenium, Tellurium and Polonium

Introduction

These elements differ considerably from oxygen:

• 1. Their electronegativity is lower.
  2. Their covalent bonding is generally weaker.
  3. Hydrogen bonding, where it is a possibility is very weak.
• The do not form compounds where pp - pp bonding is needed, but rather use dp - pp bonding especially with oxygen.
• A covalence exceeding 4 is possible by the use of empty d-orbitals. The 6-coordinate geometry is increasingly favoured down the group.
• There are sulphur compounds with very long chains (second only to carbon)
• Tellurium an dpolonium are fairly metallic in their properties.

Occurence an dReactions of the Elements

• Sulphur occurs "native" (i.e. as sulpur) in deposits from which it is extracted with high pressure hot water (Frasch process). It is also obtained from hydrogen sulphide in natural gas and petroleum - if it were left, it cause a pollution problem when the fuels are burned.

  2H₂S + SO₂ 3S + H₂O

• Selenium and tellurium come from silver and copper smelting flue gases. (These metals come from sulphide ores. Selenium and tellurium tend to be found with sulphur.)
• Sulphur forms a number of allotropes:
  • Behaviour on melting:

    Solid	~112 ^oC (just melted)	160 ^oC	444.6 ^oC (just boiling)	Vapour phase
    S8 Yellow	S13.8	Sn (n is maximum) Dark brown	Sn, (incl S3, S4 Dark red)	S8 going to S2 at higher temperatures

    A rubbery material called "plastic sulphur" can be obtaind by quickly cooling molten sulphur.

  • There are several crystalline modifications of S₈ stable at different temperatures.
  • It is possible to isolate other rings sizes from 6 to 20. Engel's sulpur contains S₆ rings.
• Sulphur is used extensively to harden synthetic and natural rubbers - "vulcanization". Bridging S₂ units are of the things that hold proteins in their correct shapes.

Hydrides

• The main one is hydrogen sulphide, H₂S, which smells of rotten eggs and is very much more poisonous than hydrogen cyanide.
• The compounds are all gases whose stability decreases down the group.
• The acidity of the hydrides increases down the group.
• The series of sulphanes, H₂S_n where n is 2 to 6 have been characterized. There are higher ones, but their separation is impossible since the chains tend to break. They are synthesized by the reaction:

  S_nCl₂ + 2H₂S H₂S_n+2 + 2HCl

Halides and Oxohalides of Sulphur

Sulphur Flourides

• Synthesis:

  S₈ + xs F₂ SF₆ (+ SF₄ + S₂F₁₀)

  SCl₂(l) + 4NaF(s) + SF₄(g) + Na₂S + 2NaF

• Sulphur tetrafluoride (bp = -30 ^oC) is quite reactive, for example it is easily hydrolysed:

  SF₄ + 2H₂O SO₂ + 4HF

  It is used as a selective fluorinating agent:

  >C=O is converted to >CF₂

  -C(O)OH is converted to -CF₃

• Sulphur hexafluoride (sublimes at -64 ^oC) is very inert kinetically (since the hydrolysis is thermodynamically very favorable). Presumably it cannot further expand its coordination sphere to a reaction intermediate.

  It is used as a gaseous electrical insulator, much better than air, because of its high dielectric constant and lack of reactivity.

Sulphur Chlorides

• Disulphur Dichloride and Sulphur Dichloride

  ¼S₈(s) + Cl₂(g) S₂Cl₂

  (SCl₂ is unstable, decomposing slowly to S₂Cl₂ and chlorine.)

  Sulphur dichloride or disulphur dichloride will dissolve more sulphur to form sulphanes upto around S₁₀₀Cl₂.

• Thionyl Chloride

  SO₂ + PCl₅ SOCl₂ + POCl₃

  Thionyl chloride is a liquid (bp = 80 ^oC) which is rapidly hydrolysed:

  SOCl₂ + H₂O SO₂(g) + 2HCl(g)

  Because the hydrolysis products are both gases, one of its great uses in inorganic chemistry is in the dehydration of hydrated metal chlorides to produce the anhydrous substance.
• Sulphuryl Chloride

  SO₂ + Cl₂ SO₂Cl₂ FeCl₃ catalyst

  It is used as a chlorinating agent in organic synthesis.

Oxides and Oxo Acids

Take note of Table 19-1 in the text - it contains an important summary of the sulphur oxoacids. Know the names and structures.

Sulphur Dioxide

• is a gas boiling at -10 ^oC which is a useful solvent when liquified and can act as a ligand towards transition metals.
• Although it is formally the anhydride of sulphurous acid, H₂SO₃ it does not react with water to any significant extent. However, the soluble salts of bisulphite, HSO₃^2–, and sulphite, SO₃^2– are well characterized.

Sulphur Trioxide

• This oxide is the anyhdride of sulphuric acid, and does indeed react with water to produce it.
• It is produced by the oxidation of sulphur dioxide with oxygen catalysed by a heterogeneous catalyst such as V₂O₅ (the "contact process") or a homogeneous catalyst such as nitric oxide (the "lead chamber process").

Thiosulphate Ion

• This ion is formed by the reaction of sulphur with sulphite:

  SO₃^2–(aq) + S(s) S₂O₃^2–(aq)

• It is used in iodometric titrations forming the tetrathionate ion:

  S₂O₃^2– + I₂ 2I^2– + S₄O₆^2–

• It is used as "fixer" in photography, specifically, it dissolves silver chloride by complexing the silver ion as [Ag(S₂O₃)₂]^3-:

  AgCl(s) + 2Na₂S₂O₃(aq) Na₃[Ag(S₂O₃)₂] + NaCl