Must be familiar with atomic properties, definitions, variations across the periods and down groups and be able to rationalise these trends.

Topics: physical properties, covalent/ionic bonding, oxidation states, allotropy, conductivity, bond energies, catenation, partial and full multiple bonding, electron deficient bonding, colour, simple compounds (oxides, halides, complexes).

Effective Nuclear Charge (Zeff)

­ across, generally no change down (except from 2s/2p to 3s/3p)

d- and f-Block Contraction

Pairing of 3p/4p and 5p/6p properties (radii, IE etc.)

e.g. covalent radii (pm):

B Al Ga In Tl

0.88 1.25 1.25 1.50 1.55

Atomic Radii

¯ across, ­ down (related to Z_eff, etc.)

Ionisation Energy (IE)

­ across, ¯ down

Anomalies across table due to stability of half-filled shell.

e.g. group 16 lower than expected: np⁴ ® np³

Electron Affinity

­ across, ¯ down

Anomalies as above due to half-filled shells.

Electronegativity (c )

­ across, ¯ down (Pauling scale used)

Very useful property in predicting the nature of elements and their compounds.

Anomalies due to d- and f-block contractions (e.g. Ga>Al).

­ ox. state, ­ electronegativity, (e.g. Tl(III) 2.04, Tl(I) 1.62)

Large decrease in electronegativity between period 2 and period 3, then much smaller decrease down each group.

Also, electronegativity increase in the sequence sp³<sp²<sp. Electrons are more tightly held in s orbitals.

Ionic and Covalent Bonding

Bonding not usually clear cut, ionic/covalent are extreme cases:

Molecular solids/liquids ® polymeric solids ® lattices

Ionic bonding is due to the attraction of oppositely charged ions e.g. M⁺X^-. Properties include: high mp, conduct electricity when molten or in solution, solids have regular crystalline lattice structures to maximise the ion-ion contacts (close packing of spheres), hard, brittle and dissolve only in polar solvents.

Covalent bonding is the sharing of valence electrons which are attracted to the positively charged nuclei of the atoms. Covalent compounds are molecular, gases, liquids or low melting solids which act as electrical insulators and dissolve in non-polar solvents (e.g. hydrocarbons).

Polar covalent bonding is intermediate (and most common) case. Compounds normally adopt 3D networks, layers, chains. Generally high mp, softer and less brittle than ionic solids. Dissolve in polar solvents. Only some are conductors in solution.

Three numbers give a guide to nature of the bond.

(1) Electronegativity

A rough guide is: if the difference in electronegativity between two elements is >1.7 it has predominately ionic character.

e.g. Be(1.6)-----F(4.0) difference 4.0-1.6=2.4 \ ionic

Be(1.6)----Cl(3.2) difference 3.2-1.6=1.6 \ Polar covalent

Pure covalent are homonuclear like O₂ or F₂ (difference is zero). No such thing as pure ionic, only a degree of character.

i.e. Na(0.9)------F(4.0) difference 3.1, but X-ray and theoretical data indicate some electron density between the atoms (not all on F) \ some covalent character.

A useful way to visualise the relationship between electronegativity and the nature of the bonding between elements is the van Arkel-Ketalaar triangle.
 
 

Large D c - ionic - lattices.

Medium D c - polar covalent - polymeric or macromolecular.

Small D c - If low S c then metallic. Small attraction between the valence electrons and the nuclei leads to delocalisation (e.g. Cs, Ca). If high S c then non-polar covalent. Large attraction between the nuclei and the valence electrons leads to localised bonding (e.g. F₂, O₂).

(2) Ionisation Energy

Low values indicate a tendency to form ionic bonding.

i.e. releases electrons readily.

S IE₁₊₂

Be 27.5 eV unlikely to lose 2 e-s totally so some covalent aspect.

Ca 18.0 eV more likely to give ionic compounds.

When two elements combine, one with a low IE (e.g. Cs, K, Ca etc.) and the second with a high electron affinity (e.g. O, F, Cl etc.) very polar bonds are formed i.e. ionic compounds.

Trend: compounds of the electropositive elements become more ionic as group is descended. Also, lower ox. states gives more ionic character (less electrons to share).

(3) Lattice Energy

Small ions pack well in a lattice. Energetically favourable since this maximises the number of positive ion-negative ion contacts. High lattice energies promote formation of ionic solids.

e.g. MgF₂ - Mg has high S IE₁₊₂ but it and F are small, pack well \ ionic.
 
 

• Remove s electron from Na (requires energy).
• Add electron to Cl atom (releases energy). However, the energy released is less than IE₁ for Na. Formation of the ions alone is not energetically favourable!
• Forming a bond (attraction of the charges) between Na and Cl makes the process favourable.
• Formation of the ionic lattice releases more energy due to extra ion-ion contacts.

Orbital and Electron Promotion Energies

Useful in rationalising oxidation states, conductivity etc.

In general, orbitals with the same principal quantum number (e.g. 2 for 2s/2p, 3 for 3s/3p) have similar energies. Note that the d-orbitals for the 3p elements and below are similar in energy to the s and p orbitals: available for promotion of electrons (increases in ox. states) and coordination of ligands (increase in coordination number).
 
 

Note: promotional energy is important for covalent bonding (sharing of electrons) while ionisation energy is important for ionic bonding.

Trend: the energy required to promote an electron from an s to a p orbital increases down a group. Important for stability of ox. states.

Oxidation and Valence States

Octet Rule: tendency for main group atoms in covalent molecules to achieve a full octet, 8 electrons (noble gas configuration). Generally most important for 2^nd row elements.

Oxidation state is the formal charge on an atom in a molecule if the electrons are assigned to the more electronegative atom in a bond.

e.g. PF₅: P⁵⁺ F^-, NH₃: N^3- H⁺, NF₃: N³⁺ F^- (only a formalism!!)

Ox. states written as 4+, 3- or IV, -III (interchangeable)

No oxidation state for element-element bonds.

e.g. Cl₂: Cl⁰ Also, N₂F₄: N²⁺ F-

But for N three electrons are used in bonding i.e. the valency of N is three. Sometimes valency is more useful than ox. state!
 
 

Common Ox. States
 

Group	1	2	13	14	15	16	17	18
Ox. State	1	2	3, 1	4, 2	5, 3,-3	6, 4, 2, -2	-1	0

Group 1

+1 most common.

-1 available in the anions M^- from the disproportion reaction:

2M(s)  M⁺ + M^-

Very strong ligands such as crown ethers (see later) stabilise the cations.

Group 2

+2 most common.

Why no M+1?

Theoretically D H_form of M⁺¹ compounds is favoured over the elements!

i.e. Mg_(s) + 1/2Cl_2(g) ® MgCl_(s) D H_form -125 kJmol^-1

BUT

Mg_(s) + Cl₂ ® MgCl_2(s) D H_form -642 kJmol^-1

MgCl₂ is far more stable (mainly due to higher lattice energy since smaller, more charged M⁺² cf M⁺¹) so disproportionation takes place:

2MgCl_(s) ® Mg_(s) + MgCl_2(s) D H_form -392 kJmol^-1

\ Don’t get M⁺¹ compounds.

Group 13

Most common +3, +1

Trend: lower oxidation state becomes more stable going down the group.

Inert pair effect

Two factors combine:

• Increase in atomic size as descend group makes weaker bonds (decrease in bond energy).

• Higher promotion energies (PE) required to involve the s electrons in bonding (trivalent state, ns²np¹ ® ns¹np¹np¹).

For Boron, B-X bonds are short and strong, and PE for s® p is relatively small \ 2 extra E(B-X)>PE \ trivalent.

But for Thallium, Tl-X bonds are long and weak, and PE for s® p is relatively large \ 2 extra E(B-X)<PE \ monovalent

BCl₃ is very stable but:

TlCl₃ ® TiCl + Cl₂ at 40 ° C
 
 

Ox. state 3 favourable Ox. state 1 favourable

Group 14

Common ox. states: +4, +2, e.g. SnCl₄, CO₂, PbCl₂, SnO.

Again the inert pair effect leads to the lower ox. state becoming progressively more stable down the group.

Group 15

Common ox. states: +5, +3, -3 e.g. PCl₃, PCl₅, Li₃N

Inert pair effect again!

High positive ox. states only available in combination with very electronegative elements e.g. O, F, i.e. high energy from formation of strong bonds offsets promotional energy.

e.g. PF₅ is known but PI₅ does not exist. Impossible to fit five large iodines round a single phosphorus!

-3 state available (due to increasing electronegativity as we go along a period) in combination with the electropositive metals. e.g. Li₃N

Hypervalency

Where more than eight electrons fill the valence shell of a main group atom. Octet rule no longer obeyed!

e.g. ns²p¹p¹p¹ ® ns¹p¹p¹p¹d¹d¹

valency: 3 5

e.g. PF₅ where No. electrons at P is 10!

Low lying d orbitals are used for bonding in hypervalent molecules (non available for period 2 elements).

Group 16

Common ox. states: +6, +4, +2, -2 e.g. SF₆, SeCl_4, SCl₂, H₂O, Li₂S

ns²p²p¹p¹ ® ns²p¹p¹p¹d¹ ® ns¹p¹p¹p¹d¹d¹

Valency 2 4 6

Again high positive ox. states only in combination with O and F.

Group 17

Most common ox. state is -1 (by far) e.g. LiCl, MgCl₂

However, in combination with electronegative atoms ox. states of 1, 3, 5 and 7 can be formed.

IF IF₃ IF₅ IF₇

ns²p²p²p¹ ® ns²p²p¹p¹d¹ ® ns²p¹p¹p¹d¹d¹ ® ns¹p¹p¹p¹d¹d¹d¹

Valency: 1 3 5 7

Trend: The higher ox. states can only be formed by the elements at the bottom of the group!

The higher ox. states are relatively unstable but can be formed by the heavier elements. This is mainly due the larger size of atoms allowing higher coordination numbers.

Group 18

Common ox. state 0 e.g. Ne, Xe

As with halogens, can form others (2, 4, 6, 8 by promotion of the electrons into empty d-orbitals) with very electronegative groups (XeF₂, XeO₄).

Xe XeO XeO₂ XeO₃ XeO₄

ns²p²p²p² ® ns²p²p²p¹d¹ ® ns²p²p¹p¹d¹d¹ ® ns²p¹p¹p¹d¹d¹d¹ ® ns¹p¹p¹p¹d¹d¹d¹d¹

Valency: 0 2 4 6 8

Second Row Anomalies

Marked difference in properties between 2nd row (Li-Ne) and 3rd row (Na-Ar) elements.

For 2^nd row:

1. Greater stability for element-element bonds (increased catenation and allotropy e.g. C vs Si).
1. Greater stability of multiple bonds (strong Nº N vs weak Pº P).
1. Octet rule generally obeyed (CF₄ but no CF₆^2- vs both SiF₄ and SiF₆^2- are stable).
1. Generally maximum coordination number of four (BF₃.NH₃ but no BF₃.2NH₃ vs AlF₃.2NH₃ stable).
1. Lower reactivity of compounds (CCl₄ vs SiCl₄).

(1) 2nd row elements have only a 1s² core shell shielding the outer electrons. This leads to high Z_eff, IE, c , small radii and contracted atomic orbitals. Also, the 2s and 2p orbitals are closer in energy and size compared to the 3s and 3p orbitals.

Effects: very efficient overlap of orbitals between 2nd row elements - strong bonds (allotropes, multiple bonding stable).

(2) Also, no low lying d orbitals for 2nd row elements.

Effects: limits oxidation number and coordination numbers to maximum of 4 (1s + 3p orbitals). Limits reactivity since no coordination sites available in compounds.
 
 

Allotropy

Different structural forms of the same element leads to very different chemical and physical properties.

General trend: Bottom left - metals, top right - non-metals, diagonal in p-block from B to Te - metalloids. Follows electronegativities.

Groups 1 and 2

All metals with either body-centred cubic, face-centred cubic or hexagonal close-packed structures.

Group 13

Boron is a metalloid (semiconductor). High electronegativity results in strong covalent B-B bonds. Iscosahedral B₁₂ polyhedra linked in 3D network. Forms multicentre bonds which are electron deficient (see later).

12 corners, 20 triangular faces. 2nd hardest element.
 
 

Gallium forms strong Ga-Ga pairs which interact weakly with six other Ga atoms (metalloid, structure related to I₂).

Cross-section through Ga.

Al, In and Tl all form metallic lattices.

The metalloid nature of Ga but not Al is due to the d-block contraction, electronegativity Ga>Al and so some extra covalent character - localised bonding.

Group 14

Shows clear trend for covalent to metallic character as ¯ group.

Carbon - non metal.

Diamond: 3D network, tetrahedral carbons, strong sp³ bonds, insulator, 2-centre 2-electron bonds.
 
 

Graphite: layers of stacked hexagonal rings. sp² bonds allowing p -bonding perpendicular to sheets – conductor (unique for a non-metal)!
 
 

Fullerenes: molecular carbon - cages. sp² carbons in hexagonal and pentagonal rings (needed for cage closure). Localised double and single bonds.

Si and Ge: metalloids with diamond structure. Semiconductor properties (see later).

Sn: Two forms.

Grey: diamond structure - mainly covalent.

White: interactions with more than four nearest neighbours - more metallic.

Pb: metallic.

Group 15

N: covalent molecular gas with a strong triple bond Nº N.

P: Three main forms, all contain three coordinate P with single P-P bonds. P₂ (Pº P) only at high temps and low pressure.

White: discrete P₄ tetrahedra.
 
 

Red: Polymeric linked P₄ units. Denser, higher mp, less reactive. Less strain in structure.

Black: most thermodynamically stable, layers of linked six-membered rings in a chair conformation.

As, Sb, Bi: all adopt hexagonal sheet structures similar to black P. At high pressures can convert Sb and Bi into metallic lattices.

In the layer structures of P, As, Sb and Bi each atom has three close neighbours and also three longer contacts to atoms in the next layer.

Trend: as the group is descended the ratio of long to short bonds decreases (become almost equal in length). Transition from localised covalent to more delocalised metallic bonding (consistent with lower c ).
 
 

Group 16

O: molecular, gaseous dioxygen with a double bond O=O. Also, ozone, O₃ with multiple bonding.

S, Se: all single bonds as rings (S₈, Se₈) or polymeric chains S_n, Se_n.

Te: Only chains as above.

Po: metallic (cubic with six nearest neighbours).

Trend: again the chain structures the interchain bond lengths decrease as the table is descended - more metallic.

Group17

All diatomic gases with a single element-element bond.

Under normal conditions: F₂, Cl₂ are gases, Br₂ is a liquid and I₂ is a solid. Trend is due to increasing van der Waals (or dipole-dipole ) interactions as the atoms become more easier to polarise.

Low temp solid state structures of the elements are similar - sheets.

Again trend towards smaller interlayer bond lengths down the group - more metallic.

Consistent with high pressure form of I₂ which becomes a conductor - forcing atoms together induces overlap of orbitals.
 
 

Group 18

Monoatomic, non-metallic.

Structural relationship between Allotropes

Starting from diamond (group 14) (a): add an electron to each atom to produce a lone pair and break one bond - structure of black P, As, Sb, Bi (group 15) (b): add another electron to form a second lone pair and break another bond - structure of S, Se, Te (group 16) (c): add another electron to form a third lone pair and break another bond - structure of the halogens (group 17).

Physical / chemical differences between white and red P

Use this as an example to illustrate consequences of structure on properties.
 

White	Red
P4 tetrahedra, soft and wax-like,  translucent, density 1.82 gcm-3,  ignites in air, soluble in hydrocarbons,  poisonous, non-conductor.	polymeric chains, powder, red, density 2.61 gcm-3,  stable in air, insoluble, non-poisonous,  semi-conductor.

These points highlight the differences between molecular and polymeric materials:
 

Molecular compounds are more volatile with lower melting points. Weaker van der Waals forces holding the molecules together.

Conductivity

Metals: conductors, non-metals: insulators, metalloids: semiconductors.

Simple picture for conductivity in metals is the sea of electrons model, where the positive metal ions are surrounded by the valence electron which are free to move - hence conduct electricity. The electrons act as electrostatic glue holding the cations together. Malleability comes from delocalised bonding in all directions. Also, metallic strength increases with number of bonding electrons, Al>Na. Lustre of metal comes from mobile electrons interacting with light.
 
 

Band Theory

Molecular orbitals form between two atoms when their atomic orbitals overlap. When the orbitals of three or more atoms combine the electrons within them are said to be delocalised. In systems with many overlapping orbitals such as metals they crowd together into bands.

e.g. For Na
 
 

Build up metal - two atoms combine their 3s orbitals to form one bonding and one antibonding orbital. This process continues to form a bonding (valence) band and an antibonding (conductance) band which are separated by an energy gap (band gap). Cannot move electrons when the valence band if filled. However, in a metal the band gap is small and it requirenly a small amount of energy to promote an electron into the unoccupied band. This leaves ‘holes’ in the valence band and allows electrical conductivity.

For metals increasing temperature results in decreasing conductivity due to increased scattering of electrons (i.e. flow not only in direction of applied charge). However, for semi-conductors where the band gap is relatively large, temperature rise allows excitation of electrons between the levels and gives a dramatic increase in conductivity!
 
 

Doping Semiconductors. Adding impurities can greatly affect conductivity of semiconductors.

e.g. n-type doping: Adding a small amount of a group 15 element such as P into a group 14 semiconductor such as Si results in P occupying the normal tetrahedral sites but with an additional electron available in the conduction band – higher conductivity.

p-type doping: Adding a group 13 element such as B to Si again results in B in tetrahedral positions but with only three electrons available for bonding – holes in valence band – higher conductivity.
 
 

Trends

For metals both thermal and electrical conductivities generally increase along a period which is directly related to the number of valence electrons available (Al>Mg>Na). Little variation down a group.

For non-metals (group 14 onwards) conductivity increases down a group due to larger, more diffuse orbitals (less directional) used to bond the elements together and the band gap becomes smaller (more delocalised bonding, more metallic character down group).

Another way to view this is through the electronegativity of the elements.

High electronegativity – insulators (electrons are stabilised in strong bonds and unable to move).

Low electronegativity – conductors (electrons are only loosely bond and may move).

Intermediate electronegativity – semiconductors.

Illustrate with group 14:

C as diamond is highly crystalline and an insulator and inert due to very strong C-C bonding. In graphite the delocalised electrons within the sheets allows a degree of electrical conductivity (unique conducting non-metal).

Si and Ge have a grey/blue metallic colour. Both adopt the diamond structure but are more volatile due to weaker, longer bonds. Both are semi-conductors, lower conductivity than a metal but increases with temperature.

Sn and Pb are low mp metals. Sn has two allotropes: a -Sn (metallic) and grey b -Sn (diamond, with semi-metal properties).

Resistivity: Diamond 10¹⁴ Si 48 Ge 47 Sn 11x10^-6 Pb 20x10^-6 (ohm cm at 20 ° C)
 
 

Superconductivity

A material which has zero electrical resistance when cooled to a definite temperature. It conducts electricity without heat loss and conducts infinitely! It is also perfectly diamagnetic which leads to repulsion of any magnetic field (levitation of magnets - superfast trains!). Usually made of complex ceramics e.g. YBa₂Cu₃O₇. Also, found for metal doped fullerenes e.g. K₃C₆₀. Also for many metals at very low temperatures e.g. Hg at 4K.
 
 

Bonding in Compounds

So far only considered a bond where an electron pair is shared between two atoms. However, many instances where this is not the case.
 
 

Electron Deficient Bonding (multicentre bonding)

An electron deficient molecule has fewer valence electrons involved in bonding than the number of valence orbitals available.

Compounds of groups 1, 2 and 13 form many electron deficient structures.

e.g. Diborane, B₂H₆

14 valence orbitals (2x4 from B, and 1x6 from H) but only 12 valence electrons (2x3 from B, and 1x6 from H). So not simple 2-centre 2-electron bonds (No. valence orbitals = No. valence electrons).

Solution: Terminal bonds are 2c-2e (4x2) leaves six electrons to cover four bridging bonds. Take the bridging bonds to contain one orbital over three atoms B-H-B and fill with two electrons each (2x2) i.e. three-centre two electron bonds.

Consider each boron atom to be sp³ hybridised and have two electrons in each B-H-B bond.

Organometallic examples:

Group 1: MeLi is a colourless, crystalline solid with fairly low mp and can be sublimed (most covalent character in group). It is soluble in organic solvents such as ether and does not conduct electricity when molten. Organolithiums form molecular, discrete structures (not ionic lattices like lithium halides).

e.g. Methyl lithium is a tetramer (MeLi)₄
 

Tetrahedron with four faces. Each Me sits over one face and hence bonds to three lithiums. Carbon is formally six coordinate!

Has electron deficient bonding: 12 C-Li bonds \ should be 24 e-s but only 8 available (4 from C and 4 from Li) so bond order is 8/24 = 1/3 \ highly reactive (needs e-s).

Other aggregation states include hexamers such as (BuⁿLi)₆.

The rest of group 1 methyl derivatives form networks or lattices (more ionic).

Group 2: Me₂Be (again most covalent in group - not ionic). Both Me₂Be and Me₂Mg form polymeric chains with five coordinate carbon. Rest of group are networks or ionic lattices.
 

Group 13: Now all Me₃M compounds are predominately covalent. All are liquids or gases under normal conditions. Me₃Al undergoes a monomer dimer equilibrium in solution. Again 3c-2e bonding for methyls.
 

Groups 14-17: Generally molecular covalent compounds which are liquids or gases e.g. CMe₄, AsMe₅, Me₂Se.

Trend: Organometallics are generally more ionic down a group and less ionic going across table.
 
 

Coordinate Bonds (or dative bonds)

This is the linkage of two atoms by a pair of electrons, both electrons being provided by one of the atoms (the donor). Usually a lone pair from one atom fills an empty orbital on the second atom.
 

usually written as A ¬ D or A.D

Lewis acid Lewis base

(e- pair acceptor) (e- pair donor)

Examples of complex formation:

(a) neutral - H₃N:® BCl₃

where donors can be ethers R₂O, amines R₃N, phosphines R₃P etc. (all contain a lone pair). The boron has an empty p-orbital which makes it a Lewis acid.

Also, SiF₄.2Pyridine. Empty low-lying d orbitals available for coordination.
 
 

Dative bonds can also form bridges to give polymers.

e.g. BeCl₂

(b) anionic - usually donor anionic, acceptor neutral.

H^- ® AlH₃ giving AlH₄^- others BF₄^-, AlF₆^3-

(c) cationic - H₃N ® H⁺ to give NH₄⁺ others H₃O⁺

Important in aqueous chemistry get (H₂O)_n ® Mⁿ⁺ usually written Mⁿ⁺_(aq).
 
 

Multiple p Bonding

(a) Full p -bonds (double, triple) are common in period 2 (C, O, N) using 2p orbitals. e.g. C=C, C=O, Cº C, Nº N, N=O. Some of the strongest bonds in table! – CO, N₂.

2s/2p orbitals are similar in size and energy and therefore hybridise well. Also, mixing of 2s/2p orbitals on adjacent atoms is highly efficient (small and localised due to high Z_eff) and form strong bonds. Not for period 3 and below which have larger, more diffuse orbitals. So only very weak Si=Si, As=As etc.
 

Example: Clear difference between N and P is diatomic N₂ vs tetrahedral P₄. Due to difference in relative bond strengths.

Nº N 946 (kJmol^-1) Pº P 490

N=N 418 P=P 310

N-N 167 P-P 200

\ Triple bond preferred for N but not P (3x167 < 946 but 3x200 > 490). Hence P forms three single bonds in preference to one triple one.

(b) Partial p -bonds (strengthened single bonds)

These are ‘dative’ p -bonds. Can be pp -pp .
 
 

B-F in BF₃ is 130pm, expected 152pm for purely single bond.

Atom p donating can be hal, O, S, N, etc.

Can also have pp -dp but only in period 3 and below when d-orbitals are available.

e.g. In POCl₃ get P-O of 145pm, expected 176.
 

Radicals: weak single bonds (due to lone pair repulsions) and strong multiple bonding leads to radical formation.

e.g. NO

Homonuclear Bond Energies

Only considering element-element single bond energies.

Trend: Generally bonds become weaker as a group is descended. Atoms are larger and poorer overlap of diffuse orbitals.

Anomalies: Weaker bonds for N, O and F than P, S and Cl!

Rationalisation: The atoms N, O and F all have at least one lone pair and their small atomic radii leads to short bonds with one another and therefore large inter-electron repulsions - bonds weaker than expected! From row three down the atoms are larger and the lone pairs on each atom are further apart - stronger bonding compared to second row.
 

Trend: Across the periods there is an increase in bond energy from group 1 to 14, then a drop at group 15, followed by an increase.

Rationalisation: bonding strength increases with increasing electronegativity but a blip at group 15 when lone pairs first destabilise the bonds.

Catenation: the ability of an element to form element-element bonds. Related to homonuclear bond strength.

Trend: Generally decreases down a group.

Anomalies at N and O due to lone pair-lone pair repulsions limiting catenation.

Note: consider binding energy of the elements in their natural forms (not just a single bond).

X_2(S) ® 2X_(g) (diatomic molecule)

X_n(S) ® nX_(g) (solid)

Trend: Increase from group 1 to group 14 then decrease to group 18.
 
 

Related to the number of valence electrons available for bonding. i.e. 1 to 4 for groups 1 to 14, then 3 to 0 for groups 15 to 18 (introduction of lone pairs). This energy is reflected in the elements properties: mp, hard or soft etc.

Heteronuclear Bond Energies
 

Trend: Generally much stronger then homonuclear bonds.

Rationalisation: heteronuclear bonds are polarised (d ⁺ d ^-) due to the difference in electronegativity between the atoms. Therefore these bonds contain a degree of ionic character (extra stabilisation through electrostatics).

Trends:

No clear trend down groups 1 and 2 but energies are higher in group 2 (higher lattice energies due to 2+ charge and smaller size).

Groups 13 - 17 decrease down group due to larger size of atoms (weaker bonding). Some anomalies at 2nd row (partial p -bonding, pp -dp for Si, P, S).

Across from 13-17 energies decrease due to less ionic contribution to bonding (electronegativities are becoming closer).

Origin of Colour

If a compound absorbs light in the visible region of the spectrum it appears coloured. Absorption of light results in the electronic transition of an electron in a molecule from one energy level to another.
 
 

The number and size of the transitions determines the colour of the compound.

Absorption in one region results in showing the transmitted light as the complementary colour.

e.g. absorb red: compound appears blue-green.

In ionic compounds the large electronegativity difference between the atoms results in a large separation between the highest energy occupied molecular orbital (HOMO) and the lowest energy unfilled molecular orbital (LUMO). Therefore the promotion of an electron requires a large amount of energy and the compounds appear colourless or white e.g. group 1 and 2 salts: NaCl, MgBr₂. Any colour present in such compounds is due to charge transfer since it involves transfer of an electron from an orbital localised on one ion to a second orbital localised on the other ion (most important for transition metal compounds).

In covalent compounds the electronegativity difference between the two atoms is small giving rise to molecular transitions of electrons and colour appears.

e.g.

SnF₄ - SnCl₄ - SnBr₄ - SnI₄

Colourless - Colourless -Yellow/Red - Orange

decreasing electronegativity difference between atoms.

Going down a group the HOMO and LUMO levels become closer resulting in colours.

e.g.

F₂ Cl₂ Br₂ I₂

Pale yellow Green-Yellow Red-Brown Violet

Flame Colours

Many metal salts are colourless unless anion is coloured (e.g. KMnO₄). This is because the energy jump to promote electrons is too high for visible light (but possible for uv). All have characteristic flame colours since ions are reduced to atoms at high temperature in flame and population of numerous excited states is increased.

Li crimson, Na yellow, K violet, Rb red, Cs blue.

e.g. Na⁺ 1s²2s²2p⁶ 3s⁰ Na 1s²2s²2p⁶3s¹ 3p⁰

Ion: large transition Atom: smaller transition

\ no colour \ colour

This is the basis for the analytical determination of many metals by atomic absorption spectroscopy.
 
 

Simple Compounds of the Main Group Elements

Simple compounds illustrate the trends outlined so far in the course.

1. Fluorides

General trend: ionic at bottom left and covalent molecular at top right separated by a diagonal band of polymeric (macromolecular). Consistent with differences in electronegativity between the elements.

Structures: ionic are lattices formed by the attraction of oppositely charged particles. e.g. NaF, CaF₂.

Covalent Molecular: little or no aggregation between molecules (non-polar) so gases, liquids or low mp solids.

Liquid (polymer) - Solid (tetramer)

Note: As ox. state increases for a given element in a series of binary compounds there is a trend from ionic to polymeric to covalent.

Example - GeF₂ is polymeric while GeF₄ is molecular. This is consistent with an increase in ox. state on Ge resulting in higher electronegativity (more positively charged), i.e. Ge(IV) is closer in electronegativity to F than Ge(II), so more covalent character in GeF₄ compared to GeF₂.

Size is also important! GeF₂ is only two coordinate and therefore space is available for further coordination (more likely to polymerise than GeF₄).

(2) Oxides

Trend from ionic to polymeric to molecular is as expected.
 

Groups 1 and 2 have lattice type structures and all of group 13 are polymeric.

In group 14 there is a stark contrast between CO/CO₂ and SiO₂ –gases versus hard polymeric solid. As mentioned previously, the strong multiple bonding between C and O leads to molecular species.

GeO₂ is similar to SiO₂ (as expected since they possess similar size and electronegativities).

SnO₂ and PbO₂ are polymeric but each metal has six nearest neighbours (larger atoms can accommodate more neighbours).

The lower oxidation states SnO and PbO show a movement towards more ionic. SnO and PbO both consist of sheets of oxygens, where a square of oxygen atoms is capped by metal atoms.
 
 

Group 15: Similar to group 14. All N oxides are molecular and contain strong p -bonding. Coordination number at N is three or less. At least eight molecular oxides, NO, NO₂, N₂O₃, N₂O, N₂O₄, N₂O₅, NO₃ and N₄O (only discovered in 1998).
 

P oxides are all larger molecules with much weaker p -bonding influences. Polymers and also oligomers (discrete molecules containing three or more atoms of an element). Two common examples P₄O₆ and P₄O₁₀.

P₄O₆ P (III) P₄O₁₀ P (V)

Other forms of P₄O₁₀ are macromolecular based on above structure but linked together.

As₄O₆ and Sb₄O₆ adopt a similar discrete structure to P₄O₆ but also macromolecular 3D structures of edge sharing EO₃ pyramids.

Bi₂O₃ is exclusively polymeric where Bi has a coordination number of 5.

Group16: Excellent example of vertical trends.

O₂ and O₃ both are gases with strong multiple bonding.

S is relatively small and therefore good overlap of orbitals with O leads to strong multiple bonding. SO₂ is a molecular gas. However, SO₃ is molecular in the gas phase but solidifies as either a cyclotrimer or a polymeric helix.
 
 

SO₂:

SO₃:

Why is SO₃ polymeric while SO₂ is only molecular? This is different from GeF₂ and GeF₄ where the lower ox. state polymerised.

Rationalisation: (1) Three electron withdrawing O atoms surrounding S induce a significantly high +ve charge on S to promote bridging. Also, (2) Coordination number of S in SO₃ is three whereas Ge in GeF₄ is four – more space available for further coordination interactions in SO₃ (coordinatively unsaturated).
 
 

SeO₂ is exclusively polymeric (consistent with move to more polar/ionic as group is descended). Linear zig-zag structure.

SeO₃ is a cyclotetramer (cf SO₃).

TeO₂ is polymeric with single Te-O bonds and a Te coordination number of 4. TeO₃ is macromolecular containing TeO₆ octahedra.

PoO₂ is ionic with the fluorite structure. Po coordination number is eight.

Trend: Larger coordination number as group is descended: 1 for O₂, 2 for SO₂, 3 for SeO₂, 4 for TeO₂ and 8 for PoO₂. Now higher ox. states become more polymeric (contrast with group 14!).
 
 

Group 17: Oxides of F, Cl and Br are molecular and contain p -bonding.
 
 

Oxides of I are oligomers or polymers (consistent with larger size, greater difference in electronegativity).

I₂O₅ sheets containing bridged subunit below.

All oxides of groups 17 and 18 are thermodynamically unstable and decompose to their elements. However, they may be kinetically stable for some time – HAZARD may detonate!

Group 18: Only XeO₃ and XeO₄ which are both molecular are known.

General Tends for Oxides: Crossing table the heteronuclear bond energies decrease (loss of ionic component to bonding).
 
 

(3) Macrocycles:

(a) Crowns and Cryptands

Conventional belief was that group 1 and 2 metal ions M⁺ or M²⁺ have weak coordinating ability due to large size and low charge. However, macrocyclic polyethers coordinated strongly with M⁺ to form stable complexes in salts.
 
 

15-crown-5  and dibenzo-14-crown-4

The stability of the complex depends on the size and shape of the cavity of the ligand compared to the size of M⁺.
 

Cation	Ion Diameter (pm)	Crown	Hole-Size (pm)
Li+	152	14-crown-4	120-150
Na+	204	15-crown-5	170-220
K+	276	18-crown-6	260-320
Rb+	304	21-crown-7	340-430

These complexes have covalent properties, low mp, soluble in organic solvents etc. X-ray data tells us the metal sits inside the cavity of the crown and the anion outside. In effect the ionic lattice has been broken down and the ionic nature of the metal is hidden by being imbedded in an organic sheath.

Cryptands are marcobicyclic compounds which have similar complexing properties to crowns except contain N/O/S/P heteroatoms.

The cryptand above captures Rb to form a bicapped trigonal prismatic structure.

Uses for crowns and cryptands include solvent extraction (excellent selectivity for ions) and phase transfer catalysis. They are also used as model compounds to mimic biological systems i.e. the transfer of Na/K between cell membranes. Crowns and crypts also form strong complexes with group 2 metals.

(b) Porphyrins

Nitrogen containing macrocycles. Commonly found in nature (haemoglobin).
 
 

Chlorophyll is a substituted porphyrin with a Mg atom lying in the centre (used in photosynthesis). Chlorophyll absorbs red light and passes this energy onto other chemical intermediates in a cell. Magnesium bonds to all four nitrogens (Lewis bases) and its function includes keeping the macrocycle flat to allow stacking of many molecules.
 
 

Highlights of Group Trends

Group 1

• All metals are malleable and become softer down the group.

• Only one electron to bond with.

• Ox. state +1 most common.

• No higher ox. states due to huge energy to remove non-valance shell electrons.

• Reactivities increase down the group.

• Decreasing IE.

• All form ionic salts with halides and oxides.

• All electropositive metals.

• Only Li reacts with N₂ (to form Li₃N) due to huge lattice energy. Similarly only Li and Na form the carbides M₂C₂.
• Form strong complexes with crowns and cryptands.
• Lithium alkyls have covalent characteristics such as solubility in hydrocarbons forming oligomeric structures (other group 1 alkyls form 3D structures).

• Anomalous Li due to high charge/size ratio.

• Form electron deficient compounds e.g. MeLi (tetramer).

• Be forms covalent compounds generally with sp³ bonding. Rest of group have mainly ionic compounds.

• Be has highest IE, smallest atom in group.

• Ox. state of +2 common.

• No +1 due to disproportionation: 2MgCl ® MgCl₂ + Mg.

• Oxides and halides from Mg down are ionic.

• Again electropositive metals.

• Reactivities increase down the group. Be is very unreactive.
• Metals are harder than those of group 1.

• Two valence electrons available for bonding.

• Coordination numbers increase down a group. Be:4, Mg: 4-6, Ca-Ba:8.

• Strong Lewis acids due to empty valance orbitals.

• Again tend to form electron deficient compounds e.g. Me₂Mg (polymer).

• B is very different to rest of group:

• semiconductor
• allotropes are molecular networks of B₁₂
• very inert
• coordination number limited to 4
• important p -bonded compounds (e.g. BF₃ and BN)

• From Al down the elements are metallic.
• Again elements are harder than group 2.
• Still electron deficient compounds e.g. B₂H₆.

• Only three valence electrons but four orbitals available (1s + 3p)

• Oxides and halides from Al down are predominately ionic but some have covalent characteristics e.g. AlI₃ and GaI₃ are molecular dimers.

• Due to close electronegativity between Al/Ga and I.

• Ox. states of +3 and +1 are common. +1 becomes more stable down the group.

• Inert pair effect.

• d-block and f-block contractions effect properties down the group.

• Pairing of Al/Ga and In/Tl for IE, size etc.

• +4 and +2 oxidation states are common. +2 becomes more stable down the group.

• Inert pair effect.

• Lower ox. states are more ionic.

• S IE 1+2 <<< S IE 1+2+3+4

• Oxides and halides change from molecular to polymeric down the group.

• Elements are now becoming more electronegative (more covalent).

• Reactivity of compounds increases down the group.

• Due to weaker bond energies and larger size of atoms.

• Trend in hardness between elements is difficult to assess now due to differences in structure.

• For example compare soft graphite with hard diamond.

• Catenation and multiple bonding decreases down the group.

• Due to poorer overlap between the orbitals, weaker element-element bonding.

• More metallic character down the group.
• Conductivity increases down the group: insulators ® semiconductors ® conductors.

• Weaker bonding to electrons leads to smaller band gap.

• Higher coordination numbers down the group.

• Hypervalency due to low lying d-orbitals, e.g. SiF₆^2-.

• Most nitrogen compounds are covalent (except nitrides Li₃N).
• Common ox. states of +3, +5, -3.

• Forms –3 with electropositive metals e.g. Mg₃N₂.

• Weak N-N single bonds.

• Due to lone pair repulsions. This leads to catenation order N<P>As>Sb>Bi.

• Oxides and halides change from molecular to polymeric.
• Generally the reactivity of elements decreases down the group.

• Exception is for N₂ which has very strong Nº N.
• Weak covalent bonds formed by lower elements due to longer weaker bonds inhibits reactivity.
• As, Sb and Bi are stable in air.

• Strong multiple bonds only for N.

• Three single P-P bonds are stronger than one P-P triple bond.

• More metallic character down the group.
• Lone pair available for ox. state 3 compounds e.g. NH₃, so Lewis bases. Also for period 3 and below low lying d-orbitals allow acceptance of an electron pair.

• May also be Lewis acids or Lewis bases e.g. SO₂.

• More metallic down group.

• O-Se are non-metals, Te is a metalloid and Po is metallic.

• Lighter elements readily form anions e.g. O^2-, S^2-.

• Now have a significant degree of electronegativity and can stabilise a negative charge.

• Very strong multiple bonding for O.

• Strong pp - pp bonding due to efficient overlap of orbitals.

• Also, dp -pp bonding for S and Se.

• Empty low lying d-orbitals available.

• Increasing coordination numbers down the group.

• C can be 1-4 (CO, CO₂, H₂C=CH₂, CH₄), S can be up to 6 and Te can be 8.

• O tends to have ox. state of –2 but others form 2, 4 and 6.

• Oxygen is the most electronegative in the group.

• Catenation O<S>Se>Te>Po.

• Single bond energy for O is weak due to lone pair repulsions. Weaker bonds down the rest of the group from S due to larger atoms and less efficient overlap of orbitals.

• Reactivity again generally decreases down the group.

Group 17

• All elements are diatomic.

• Fits octet rule.

• Melting points increase down the group.

• Due to increasing van der Waals interactions.

• Although the elements are non-metals they form highly ionic compounds with metals, e.g. NaCl.

• Highly electronegative elements, so form ionic compounds with electropositive elements.

• F forms the most ionic compounds with metals. Decreasing ionic character of MX down the group due to decreasing electronegativity.
• Ox. state of –1 is most common. Other available for interhalogens e.g. IF₅, ICl₂^-.

• Primarily inert molecular gases.

• Filled octet.

• Can form compounds with very electronegative elements, e.g. XeO₂.

• Only for heavier elements Kr, Xe and Rn. Others too electronegative to share electrons.

• Oxidation states of 2, 4, 6, 8 are available.

• Use of low-lying d-orbitals.