The chlorides and oxides of period 3 elements of the periodic table have been extensively studied and investigated and include many interesting and industrially important compounds. These compounds exhibit distinctive properties which vary widely across period 3, ranging from the use of sodium chloride in the chlor-alkali industry to produce chlorine and caustic soda (described by Cheresource, 2004); the use of inert silicon dioxide as a filler in building materials (see for example, Dorfner, no date (n.d.)), glass and in the electronics industry; phosphorus pentoxide (P4O10) which takes advantage of its strong affinity for water to be an excellent drying agent (Cotton FA, 1999); and the volatile gaseous oxides of sulphur used in the contact process for the production of sulphuric acid (as described by Encarta, 2007). The following sections will outline how these broad ranging physical and chemical properties can be explained in terms of the chemical bonding and structure of the oxides and chlorides, and the way in which the trends manifest themselves from left to right across group 3 of the periodic table.
Trends in the PHYSICAL properties of the OXIDES
The group 3 elements (sodium to chlorine) form a number of oxides of varying oxidation states as illustrated in table 1 below. Note that the noble gas argon does not form stable isolatable oxides due to its electron configuration with a full octet of outer shell electrons.
The most important oxides of group 3 elements, those in which the element is in its highest oxidation state, are highlighted in bold type. Graph 1, below, illustrates the boiling points of these highest oxides.
The differences in the boiling points in graph 1 can be explained in terms of the bonding and structures of each of the oxides. Sodium, magnesium and aluminium all form giant metallic oxides which contain strong ionic bonds. These bonds are difficult to break, resulting in high melting and boiling points. Western Oregon University a(n.d) explains the increase in boiling point from Na2O to MgO in terms of the increase in charge on the metal (i.e. Na+ vs. Mg2+), which increases the attraction between the positive metal ions and negative O2- ions. Based on this theory, it might be expected that the Al3+ cation in Al203 would produce an even greater coulombic force, resulting in a higher boiling point for aluminium oxide. However, on moving from left to right across the periodic table the electronegativity of the elements increases (on the Pauling scale described by Huheey JE, 1993), and hence, the difference in attractive forces between the metal cation and oxo anion is reduced as shown in table 2. This electronegativity difference has a greater influence on the boiling point of aluminium oxide than the increased charge on the cation, resulting in the boiling point of Al2O3 being lower than that of MgO.
As seen in table 2, the electronegativity difference between silicon and oxygen is only 1.54, which indicates that the silicon oxygen bonds in SiO2 are more covalent in nature than the ionic bonds observed in the metallic oxides. As described by Amethyst Galleries Inc (2006), there are two forms of pure SiO2, namely quartz and cristobalite, both of which have giant covalent structures with high melting and boiling points due to the strong silicon-oxygen covalent bond and regular tetrahedral arrangement of the atoms. The structure of cristobalite is very similar to that of diamond. The silicon atoms are each bonded to four oxygen atoms in a tetrahedral arrangement. The oxygen atoms form bridges between the silicon atoms to generate a vast repeating 3 dimensional structure.
The later oxides of period 3 (P4O10, SO3 and Cl2O7) have considerably lower melting and boiling points than the giant structures seen for sodium, magnesium, aluminium and silicon oxides. These elements form molecular oxides in which the molecules are held together by significantly weaker van der Waals dispersion forces or simple dipole-dipole interactions. The P4O10 molecule, which contains 6 bridging and 4 terminal oxygen atoms is shown in figure 1:
As described by Clark, J (2007) a sulphur trioxide has a number of different structures, the simplest being a trimeric ring structure. This explains why SO3 is a solid at room temperature (BP = 45°C) compared to gaseous SO2 (BP -10°C) which has a simple unimolecular structure.
The conductivity of the group 3 oxides can also be understood by considering the structure and bonding of the compounds. The giant metallic structures of Na2O, MgO and Al2O3 do not conduct electricity in the solid state because the ions are held rigidly in place. However, they all conduct via electrolysis in the molten state, in which the ions are free to move and discharge. This is particularly important in the smelting of aluminium oxide to form aluminium as described by the International Aluminium Institute (2000). The remaining oxides of silicon, phosphorus, sulphur and chlorine do not conduct due to their covalent/molecular structures.
Trends in the CHEMICAL properties of the OXIDES
As a general trend, the oxides vary from basic on the left hand side of period 3 (Na2O, MgO), through amphoteric (Al2O3), to acidic on the right hand side (SiO2, P4O10, SO3). This trend in chemical property is explained (by Western Oregon Univeristy b, (n.d.)) in terms of the electronegativity difference between the cation and oxygen, and how this affects the tendency of the oxide to donate or accept electron pairs. Due to the large difference in electronegativity between sodium and oxygen the Na2O molecule is highly polarised and genrates the highly basic oxide ion O2-. This oxide ion is readily protonated by water in a strongly exothermic reaction (reactions provided by Beavon, R (2007)):
Moving across period 3 the difference in electronegativity between the element and oxygen decreases, and hence the basic strength (the ability to form the oxide ion O2-) decreases too. This is exemplified by the fact that MgO is only partially soluble in water and Al2O3 is insoluble in water. However, magnesium oxide is still sufficiently basic to react with acids to form magnesium chloride (reactions provided by Bateman DW, (2007)):
As described in section 2, SiO2 has a stable, covalent structure, in which the low electronegativity difference between silicon and oxygen prevents it from forming basic oxide ions. Hence, it does not react with water or acid. Nevertheless, the SiO2 is weakly acidic, and will react with strong bases such as NaOH to form sodium silicate.
Phosphorous pentoxide (P4O10) is acidic in nature and reacts violently with water to form a mixture of phosphoric acids (as described by Cotton FA, (1999)). Water attacks the electron rich bridging oxygen atoms shown in figure 1 earlier, resulting in the cleavage of the P-O-P bridge.
The acidity of the oxides continues to increase towards the right hand side of period 3, as the tendency to accept electron pairs from the oxygen atom of a water molecule increases. The following equations illustrate the reactions of SO3 and Cl2O7 with water to form the strong sulphuric and perchloric acids respectively:
The reaction of sulphur trioxide with water is particularly important, as described by Western Oregon University (n.d), since it is responsible for the formation of acid rain from the air pollutant SO3.
Trends in the PHYSICAL properties of the CHLORIDES
With the exception of the noble gas argon, the group 3 elements (sodium to sulphur) form a number of chlorides of varying oxidation states as illustrated in table 3 (modified from Clark, J (2007)b). Notice that, similarly to the oxides, the most important chlorides are those in which the period 3 element is in its highest oxidation state, (except in the case of sulphur, where S2Cl2 is the most stable and studied chloride).
The physical properties of the period 3 chlorides follow a similar trend to the oxides already discussed. Graph 2 shows that sodium and magnesium chlorides have high boiling points due to their giant ionic lattice structures, which require large amounts of energy to break the strong metal-chloride bonds.
The behaviour of aluminium trichloride during heating is more complex than that of sodium and magnesium and cannot necessarily be interpreted in terms of periodic trends. Initially, AlCl3 has a similar lattice structure to NaCl and MgCl2, however, the increased electronegativity of aluminium (vs. Na and Mg) means that the lattice has more covalent character. At around 180°C the aluminium chloride interchanges into a dimeric molecular form, Al2Cl6 (as shown in figure 2). This dimer is only held together by weak van der Waals and dipole-dipole interactions and, as a result, it sublimes rapidly after losing its lattice structure.
The remaining three elements all form simple molecular chlorides, (SiCl4, PCl5 and SCl4). All of these compounds have low boiling points because the molecules are only held together by weak van der Waals forces. SiCl4 and SCl4 are both gases at room temperature, whereas PCl5 has a slightly higher boiling point than the periodic trend may suggest due to slight differences in its structure. PCl5 actually exists in a dimeric form in the solid state. In this form, a chloride can transfer from one phosphorus atom to another to give oppositely charged complex ions (as described by Cotton FA, 1999).
When the dimer reaches 166.8°C the structure converts to molecular PCl5 and sublimes.
Trends in the CHEMICAL properties of the CHLORIDES
The chlorides of period 3, shown previously in table 3, undergo a range of reactions with water. The force of these reactions depends on the structure and bonding of the chloride. The rate of hydrolysis of the metal chlorides (see reaction schemes below) increases with increasing charge on the cation.
Since the Na+ cation is not sufficiently charged to attract water ligands or large enough to accommodate them, sodium chloride ionises and dissolves in water to form a neutral solution and does not form a hydrated species. However, the doubly charged cation of magnesium is sufficiently strong to afford hydrolysis and formation of the hexaaqua complex. This complex is mildly acidic due to the attraction of the electrons of the water molecules towards the Mg2+ cation. This weakens the O-H bond of water, allowing the hexaaqua complex to donate H+ to free water molecules as shown in the following equation:
The equilibrium above is significantly displaced to the left, since the Mg2+ cation is only sufficiently powerful to make the hexaaqua complex slightly acidic. By contrast, the Al3+ cation polarises the water molecules to a much greater extent than Mg2+, which results in a significantly more acidic hexaaqua complex. Hence, the analogous proton donating reaction for aluminium (shown below) is displaced to the right.
Finally, silicon tetrachloride and phosphorus pentachloride both react violently with water, generating fumes of hydrogen chloride:
The variation in the physical and chemical properties of the chlorides and oxides of the elements across period three can be explained by general trends in the bonding and structure of the compounds. The elements towards the left hand side of the period form large, stable ionic structures with high melting and boiling points and compounds which exhibit strongly basic chemical characteristics. Moving across the period, the bonding becomes increasingly covalent, with the compounds on the right hand side forming simple molecular structures with low boiling points, and increasing acidic nature. However, there are certain anomalies in these trends in physical and chemical properties, which can be explained by detailed consideration of the individual structures.