Choose the more metallic element from each pair. (a) Sr or Sb (b) As or Bi (c) \(\mathrm{Cl}\) or \(\mathrm{O}\) (d) \(S\) or \(\mathrm{As}\)

Understanding the trends of the periodic table is essential for grasping the concept of metallic character in chemistry.

Metals tend to lose electrons and form cations easily, a characteristic that is reflected in their position on the periodic table. As you look at a periodic table, you'll notice that metallic character increases from top to bottom within a group. This is because atoms get larger down a group, and the outer electrons are further from the nucleus; hence, they are held less tightly and can be lost more easily.

From left to right across a period (row), metallic character decreases. This happens because the atoms become smaller, their positive charge increases with each element, and electrons are added to the same energy level. The increased nuclear charge means that the nucleus holds onto the electrons more strongly, making it less likely for these elements to lose electrons and behave like metals.

Elements on the left-hand side of the periodic table are generally metallic, while those on the right-hand side exhibit non-metallic characteristics. Metals are usually found in Groups 1 (alkali metals) and 2 (alkaline earth metals) and in the transition metal blocks, while non-metals are housed to the right of the staircase line that distinguishes them from the metals.

The formation of cations is a process that is intricately linked to the metallic character of an element. Cations are positively charged ions that are created when an atom loses one or more of its electrons.

This electron loss is more readily observed in elements with a higher metallic character due to their lower electronegativity and larger atomic radii, which make it easier for the outermost electrons to be removed by chemical reactions. For instance, alkali metals in Group 1, which have a single valence electron, readily form cations with a +1 charge by losing that electron.

In contrast, non-metals, which are found on the right side of the periodic table, have higher electronegativity and much smaller atomic radii. This makes them more inclined to gain electrons, forming anions, not cations. The tendency of an element to form a cation is a helpful indicator of its metallic character and is a concept frequently employed by chemists to predict the behavior of elements in reactions.

When comparing elements across groups and periods to determine their metallic character, it's vital to draw on our understanding of periodic table trends.

In the periodic table, a group refers to a vertical column. As mentioned earlier, within a group, the metallic character increases as you move down. This is because each element down the group has an additional energy level compared to the one above, making it less efficient at holding onto its valence electrons. As a consequence, lower elements in a group are more likely to exhibit metallic properties; they are larger and less electronegative.

Conversely, a period refers to a horizontal row on the periodic table. Along a period, metallic character diminishes moving from left to right. This occurs because elements towards the right have a higher number of valence electrons and a stronger attraction between the nucleus and electron cloud, which decreases their ability to lose electrons and therefore their metallic character.

Utilizing this comparison of groups and periods helps us understand why elements like Cesium (Cs) in Group 1 are far more metallic than Fluorine (F) in Group 17, although they are in the same period. Likewise, Lead (Pb) is more metallic than Carbon (C) although they fall in the same group, simply because Pb is located lower in the group.