Doc Brown's Chemistry - Advanced Level Inorganic Chemistry Periodic Table Revision Notes

 Part 9. Group 7/17 The Halogens

9.11 Chemical Bonding in Halogen Compounds (elements, halides, halo–compounds etc.)

Examples of the chemical bonding in covalent and ionic halogen compounds are described with electron configurations and dot & cross diagrams and structural formulae and molecular shapes.

9.11 Chemical Bonding in Halogen Compounds

(1) Ionic bonding * (2) Covalent bonding * (3) Oxyanions and ligands

(1)  Ionic bonding in halide salts = Factors affecting bond formation

The large differences in electronegativity between metals and non–metal halogens often results ionic bonding

Ex 1 ONE  combines with ONE to form

electron configurations [1s²2s²2p⁶3s¹] + [1s²2s²2p⁶3s²3p⁵] ==> [1s²2s²2p⁶]=[Ne] + [1s²2s²2p⁶3s²3p⁶]=[Ar]

showing the electron transfer from the electropositive metal to the electronegative non–metal and the ensuing ion formation

Ex 2 ONE  combines with TWO to form

electron configurations [1s²2s²2p⁶3s²] + 2[1s²2s²2p⁶3s²3p⁵] ==> [1s²2s²2p⁶]=[Ne] + 2[1s²2s²2p⁶3s²3p⁶]=[Ar]

Ex 3 ONE  combines with THREE to form

electron configurations [1s²2s²2p⁶3s²3p¹] + 3[1s²2s²2p⁵] ==> [1s²2s²2p⁶]=[Ne] + 3[1s²2s²2p⁶]=[Ne]

• The greater the electronegativity between two bonded atoms, the greater the likelihood of an ionic bond being formed.
• However, the electronegativity of an element itself is not the only factor in deciding bond character!
  • The different results of the 'tug of war' between two positive nuclei acting on the intermediate bonding electrons produces a range of bond character from complete electron transfer in ionic bond formation (e.g. M⁺ and X^–), to a highly polar covalent bond (M^δ+–X^δ–) of partially charged atoms and at the other extreme a virtually non–polar bond (X–Y) of two atoms, neither of which carries a significant partial charge.
  • Electronegativity, the power of an element to attract bonding electrons towards it in a bonding situation, is just one, albeit important, factor in deciding the outcome of the character of an individual bond.
  • The stronger the polarising power of the cation and the higher the polarizability of the anion the more covalent character is expected in a bond.
  • If a cation has appreciable polarising power to draw bonding electron clouds towards it OR the bonding electron clouds of an anion are attracted towards the cation, then covalent bonding character is more likely.
  • Polarising power for cations is very much a case of increasing with increased ion charge/ionic radius.
    • So, the smaller the ionic radius or the bigger the positive charge, the greater the polarising power of the cation.
    • e.g. in terms of polarising power Al³⁺ > Mg²⁺ >Na⁺ for the series of Period 3 positive ions where you have both coincident decreasing radii and increasing charge.
    • For Group ions, polarising power will decrease down the group with increasing ionic radius and constant charge.
    • therefore in polarising power Li⁺ > Na⁺ >K⁺ for Group 1 Alkali Metals etc.
    • or Be²⁺ > Mg²⁺ > Ca²⁺ for Group 2 Alkaline Earth Metals etc. (see also s–block notes)
    • Examples of the outcome of this factor are
      • increasing ionic character NaCl > MgCl₂ > AlCl₃
        • The latter is just ionic in the lattice, but vaporises to a covalent dimer.
      • increasing ionic character KCl > NaCl > LiCl
        • As the cation gets larger for same charge its polarising power diminishes.
      • Change in oxidation state can also change the bond character significantly. Iron(II) chloride is essentially ionic in character and iron(III) chloride is basically covalent because the polarising power of the smaller and more highly charged iron(III) ion.
  • For anions, the larger the ionic radius and the greater its charge, the more polarisable it is.
    • So in terms of polarizability I^– > Br^– > Cl^– > F^– (for halide ions for constant charge and decreasing radius)
    • or polarizability of Si^4– > P^3– > S^2– > Cl^– (for a series of Period 3 anions of decreasing charge and decreasing ionic radius)
    • therefore you expect for ...
    • A series of Group 2 halides the ionic character CaCl₂ > MgCl₂ > BeCl₂
      • Calcium chloride is essentially ionic and beryllium chloride is essentially covalent.
    • The series of Period 3 chlorides the ionic character be NaCl > Na₂S > Na₃P > SiCl₄
      • In fact sodium chloride is very ionic high melting lattice and silicon(IV) chloride is very covalent low boiling liquid!
    • A series of Group 1 halide salts the ionic character trend is KF > KCl > KBr > KI
      • Potassium iodide is essentially ionic, but its 'partial' covalent character is shown by the fact that it dissolves in polar solvents like propanone (acetone) whereas highly ionic potassium chloride is ~insoluble in polar organic solvents.
  • What this set of paragraphs illustrates is a much wider and deeper approach to electronegativity than e.g. the electronegativity number quoted on the Pauling scale.

Properties of ionic compounds

• The diagram above shows a typical of the giant ionic crystal structure of ionic compounds like sodium chloride and magnesium oxide.
• The alternate positive and negative ions in an ionic solid are arranged in an orderly way in a giant ionic lattice structure.
• The ionic bond is the strong electrical attraction between the positive and negative ions next to each other in the lattice.
• The bonding extends throughout the crystal in all directions.
• Halide salts are typical ionic compounds.
• This strong bonding force makes the structure hard (if brittle) and have high melting and boiling points, so they are not very volatile!
• A relatively large amount of energy is needed to melt or boil ionic compounds. Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.
• The bigger the charges on the ions the stronger the bonding attraction e.g. magnesium oxide Mg²⁺O^2– has a higher melting point than sodium chloride Na⁺Cl^–.
• Unlike covalent molecules, ALL ionic compounds are crystalline solids at room temperature.
• They are hard but brittle, when stressed the bonds are broken along planes of ions which shear away. They are NOT malleable like metals.
• Many ionic compounds are soluble in water, but not all, so don't make assumptions.
  • Salts can dissolve in water because the ions can separate and become surrounded by water molecules which weakly bond to the ions.
  • This reduces the attractive forces between the ions, preventing the crystal structure to exist. Evaporating the water from a salt solution will eventually allow the ionic crystal lattice to reform.
• The solid crystals DO NOT conduct electricity because the ions are not free to move to carry an electric current. However, if the ionic compound is melted or dissolved in water, the liquid will now conduct electricity, as the ion particles are now free.
• See also Notes on ionic bonds and ionic compounds

(2)  Covalent bonding in small molecules of the halogens and their compounds

The smaller differences in electronegativity between two non–metallic elements usually results in covalent bonding

Two chlorine atoms (2.8.7) form the molecule of the element chlorine Cl₂

Ex 4  and combine to form where both atoms have a pseudo argon structure of 8 outer electrons around each atom. All the other halogens would be similar e.g. F₂, Br₂ and I_{2 etc}. Valency of halogens like chlorine is 1 here.

Ex 5 One atom of hydrogen (1) combines with one atom of chlorine (2.8.7) to form the molecule of the compound hydrogen chloride HCl

 and combine to form where hydrogen is electronically like helium and chlorine like argon. All the other hydrogen halides will be similar e.g. hydrogen fluoride HF, hydrogen bromide HBr and hydrogen iodide HI. These are polar molecules due to the greater electronegativity of the halogen compared to hydrogen i.e. H^δ+–X^δ–

Note: Hydrogen chloride gas is a true covalent substance consisting of small HCl molecules. If the gas is dissolved in a hydrocarbon solvent like hexane or methylbenzene it remains as HCl molecules and because there are no ions present, the solution does not conduct electricity. However, if hydrogen chloride gas is dissolved in water, things are very different and the HCl molecules split into ions. Hydrochloric acid is formed which consists of a solution of hydrogen ions (H⁺) and chloride ions (Cl^–). The solution then conducts electricity and passage of a d.c. current causes electrolysis to take place forming hydrogen and chlorine.

Ex 6 gaseous beryllium halides BeCl₂ (X = Be, Q = F, Cl, X = Be)

Linear molecules with a Q–X–Q bond angle of 180^o.

Ex 7 Q = Group 7 F, Cl etc. X = Group 5 N, P etc.

Pyramidal or trigonal pyramid shaped molecules with a Q–X–Q bond angle of ~180^o

Ex 8

COCl₂ is a good example to practice your dot and cross diagram skills. A trigonal planar molecule with a Cl–C=O bond angle of ~120^o

Ex 9  gaseous Group 3 halides Q = F, Cl and X = B, Al for F

Trigonal planar molecules with a Q–X–Q bond angle of 120^o

Ex 10   chloromethane, tetrahedral shape, bond angles of H–C–H and H–C–Cl are all ~109^o

Ex 11 Q = Cl, X = C, Si, Ge etc. in group 4, tetrahedral, bond angles 109o

Ex 12 e.g. PF₅ X=P, Q=F, trigonal bipyramid, bond angles, 180, 120 and 90^o

Ex 13 e.g. SF₆ X=S, Q=F, octahedral shape, bond angles 90 and 180^o

Ex 13 OR ? !

Aluminium chloride is a curious substance in its behaviour. The solid, AlCl₃, consists of an ionic lattice of Al³⁺ aluminium ions, each surrounded by six Cl^– chloride ions, BUT on heating, at about 180^oC, the thermal kinetic energy of vibration of the ions in the lattice is sufficient to cause it break down and sublimation takes place (s ==> g). In doing so the co–ordination number of the aluminium changes from six to four to form the readily vapourised covalent dimer molecule, Al₂Cl₆, shown above.

Read the discussion on electronegativity and cation polarising power in the ionic bonding section.

See also Covalent Bonding – small simple molecules and properties

Properties of small covalent molecules

• The electrical forces of attraction, that is the chemical bond, between atoms in a molecule are usually very strong, so,  most covalent molecules do not change chemically on moderate heating.

  • e.g. although a covalent molecule like iodine, I₂, is readily vapourised on heating, it does NOT break up into iodine atoms I. The I–I covalent bond is strong enough to withstand the heating and the purple vapour still consists of the same I₂ molecules as the dark coloured solid is made up of.

• So why the ease of vaporisation on heating?

  • The electrical attractive forces between individual molecules are weak, so the bulk material is not very strong physically and there are also consequences for the melting and boiling points.

• These weak electrical attractions are known as intermolecular forces and in particular instantaneous dipole – induced dipole forces.

• These intermolecular forces are readily weakened further on heating.

  • The effect of absorbing heat energy results in increased the thermal vibration of the molecules in the solid which weakens the intermolecular forces.

  • In liquids the increase in the average particle kinetic energy makes it easier for molecules to overcome the intermolecular forces and change into a gas or vapour.

  • Consequently, small covalent molecules tend to be volatile liquids with low boiling points, so easily vapourised, or low melting point solids.

  • On heating the inter–molecular forces are easily overcome with the increased kinetic energy of the particles giving the material a low melting or boiling point and a relatively small amount of energy is needed to effect these state changes.

  • Energy changes for the physical changes of state of melting and boiling for a range of differently bonded substances are compared in a section of the Energetics Notes.

  • This contrasts with the high melting points of giant covalent structures with their strong 3D network.

  • Note: The weak electrical attractive forces between molecules, the so called intermolecular forces should be clearly distinguished between the strong covalent bonding between atoms in molecules (small or giant), and these are sometimes referred to as intramolecular forces (i.e. internal to the molecule).

• Covalent structures are usually poor conductors of electricity because there are no free electrons or ions in any state to carry electric charge.

• Most small molecules will dissolve in some solvent to form a solution.

  • 'Like dissolves like' is a useful, if limited, 'rule of thumb'. Chloromethane and other halogenoalkanes will dissolve in hydrocarbon solvents but not highly polar water.

  • This again contrasts with giant covalent structures where the strong bond network stops solvent molecules interacting with the particles making up the material.

• Covalent Bonding – small simple molecules and properties

(3) Other examples of bonding in halogen compounds – oxyanions and ligands

Ex 15

Sketches of the dot & cross diagrams for the following anions: chlorate(III) – bent shape, chlorate(V) – pyramidal (trigonal pyramid) shape and chlorate(VII) – tetrahedral shape. Although ions, all the internal bonds are covalent and the some of the electrons are delocalised in orbitals between the Cl and O atoms constituting the bond.

Ex 16

The chloride ion acting as a electron pair donor ligand in transition metal complexes forming dative covalent bonds with the central transition metal ion

Left: The octahedral complex dichlorotetraaquachromium(III) ion.

Right: The geometric (or E/Z) isomers of the square planar complex platinum(II) complex 'platin'.

In both cases all the bond angles are 90 or 180^o

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