Oxides, of Period 3 elements

Lesson 2 -Reaction of elements of Period 3 with oxygen. -Variation in melting points of oxides of Period 3 elements. -Use of a spider concept map (Fig. 5) to test students prior knowledge. -Use of a concept map (Fig. 6) for formative assessment. -Use of a spider concept map (Fig. 7) for summarizing lesson and as an evaluative tool. 

Fig. 6 Concept map used for assessing students understanding of the oxides of Period 3 elements... 

Table 3.22 summarizes the formulas and properties of the oxides of period 3 elements. 

As with the oxides of Period 3 elements, the chlorides also show characteristic behaviour when we add them to water. Once again, this is linked to their structure and bonding (Table 10.12). 

The oxides of period 3 change from ionic to covalent across the period. Oxides of the left-hand side metals are basic, whereas non-metallic oxides of the right-hand side elements are acidic. There is a gradual change from basic to acidic character of the oxides across the period ... 

Silicon [7440-21-3] Si, from the Latin silex, silicis for flint, is the fourteenth element of the Periodic Table, has atomic wt 28.083, and a room temperature density of 2.3 gm /cm. SiUcon is britde, has a gray, metallic luster, and melts at 1412°C. In 1787 Lavoisier suggested that siUca (qv), of which flint is one form, was the oxide of an unknown element. Gay-Lussac and Thenard apparently produced elemental siUcon in 1811 by reducing siUcon tetrafluoride with potassium but did not recognize it as an element. In 1817 BerzeHus reported evidence of siUcon occurring as a precipitate in cast iron. Elemental siUcon does not occur in nature. As a constituent of various minerals, eg, siUca and siUcates such as the feldspars and kaolins, however, siUcon comprises about 28% of the earth s cmst. There are three stable isotopes that occur naturally and several that can be prepared artificially and are radioactive (Table 1) (1). 

Figure 8.17 The trend in acid-base behavior of eiement oxides. The trend in acid-base behavior for some common oxides of Group 5A(15) and Period 3 elements is shown as a gradation in color (red = acidic blue = basic). Note that the metals form basic oxides and the non-metals form acidic oxides. Aluminum forms an oxide (purple) that can act as an acid or as a base. Thus, as atomic size increases, ionization energy decreases, and oxide basicity increases.

The fact that the elements become less metallic, from left to right across a period, is shown by the characters of their chlorides and oxides. The chlorides of metals on the left-hand side of a period are ionic and dissolve in water to form neutral solutions whereas the chlorides of non-metals on the right-hand side of the table are covalent and react with water. There is a gradual change between these two extremes across the period. Consider the chlorides of period 3 ... 

Table 7.6 lists the oxides of the elements of the s- and p-blocks of periods 3-6. There is a general pattern across each period. This is a transition from ionic basic oxides, through polymeric covalent oxides, some being amphoteric and the later ones being acidic, to the molecular acidic oxides of the later groups. Down each group there is a tendency for the oxides to be of similar stoichiometry and to be more basic, less acidic, towards the heavier members. 

Table 7.6 The oxides of the s- and p-block elements of Periods 3-6. The oxides shown in red are the most stable, those with the barker background are ionic, those with the lighter background are polymeric and the remainder, without background, are molecular. The oxides with double stars are basic, those with one star are amphoteric and the starless remainder are acidic ...

Q6 Describe the acid-base character of the oxides of the period 3 elements Na to Ar. [3]... 

Table 9.2 Formulas of the highest oxides and their oxidation numbers in the elements of period 3... 

We will now look at the chemistry of some of the elements of Period 3 and their compounds, focusing on the oxides and chlorides. 

Table 10.7 shows the formulae of some of the common oxides of the Period 3 elements. 

The maximum oxidation number of each element rises as we cross the period. This happens because the Period 3 element in each oxide can use all the electrons in its outermost shell in bonding to oxygen (ox. no. = -2). They all exist in positive oxidation states because oxygen has a higher electronegativity than any of the Period 3 elements. See page 60 for more about electronegativity. 

Table 10.7 Oxidation numbers of the Period 3 elements in some common oxides. Chlorine has other oxides, such as CijO, in which its oxidation number is +1, and CljOg, in which its oxidation number is +5.

The electronegativity of oxygen is 3.5. The greater the difference in electronegativity between the Period 3 element and oxygen, the more likely it is that the oxide will have ionic bonding. Electrons will be transferred... 

Oxidation numbers of the Period 3 elements in their chlorides... 

Mendeleef based his original table on the valencies of the elements. Listed in Tables 1.6 and 1.7 are the highest valency fluorides, oxides and hydrides formed by the typical elements in Periods 3 and 4. 

Iron is the second most abundant metal on earth. It is a group 8 and period 4 element with [Ar] 3cf4s as electronic configuration. Iron as a metal is rarely found because it oxidizes readily in the presence of oxygen and moisture. Hence, it forms salts in its preferred oxidation state +2 and +3. 

Not only molecules with LLPCN > 4, but all molecules of the elements in period 3 and beyond in their higher valence states, including most of their numerous oxides, oxoacids, and related molecules such as SO3 and (H0)2S04 should be regarded as hypervalent if AO bonds are described as double bonds (1). However, Lewis did not regard these molecules as exceptions to the octet rule because he wrote the Lewis structures of these molecules with single bonds and the appropriate formal charges (2). 

In Sec. 13.2 we will learn to determine oxidation numbers from the formulas of compounds and ions. We will learn how to assign oxidation numbers from electron dot diagrams and more quickly from a short set of rules. We use these oxidation numbers for naming the compounds or ions (Chap. 6 and Sec. 13.4) and to balance equations for oxidation-reduction reactions (Sec. 13.5). In Sec. 13.3 we will learn to predict oxidation numbers for the elements from their positions in the periodic table in order to be able to predict formulas for their compounds and ions. 

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