Metals in the first group of the periodic table, called alkali metals, are very reactive metals (see Activity Series of Metals).
The alkali metals are (in increasing atomic number order): lithium (Li), sodium (Na), potassium (K), rubidium (Rb), caesium (Cs) and francium (Fr).
They are monovalent with an \(ns^1\) outer shell strcuture and they can't have any other valence.
With electropositivity increasing down in the group, francium is believed to have the highest electropositive character. However, it was only obtained in very small quantities as it's very unstable and radioactive, so, until a way to stabilize it is found, caesium is the most reactive metal.
Their high reactivity, makes them violently react with water even at room temperature, according to the equation \(\ce{M + H2O \rightarrow MOH + \frac{1}{2}H2}\), forming MOH hydroxides that are very basic (almost fully ionized). This reaction is very exothermic and sometimes the heat generated is enough to ignite the released hydrogen, resulting in an explosion.
Alkali metals are also very reactive towards atmospheric oxygen, quickly reacting to form oxides when exposed to air. That's why, when exposed to air, their metalic shine quickly disappears. To prevent this, alkali metals are usually stored in oil.
They can form a number of different oxides: normal oxides (\(\ce{M2O}\)), peroxides (\(\ce{M2O2})\) and superoxides (\(\ce{MO2})\).
Because of their very highy electropositive character, alkali metals form very polarized bonds with a very predominant ionic character. This makes their salts almost always soluble (see Solubility Equilibra). Alkali metal cations can usually only be precipitated by complex organic reagents. One exception is slighly soluble potassium perchlorate \(\ce{KClO4}\)
In analytical contexts, a lot of anions are found in the form of an alkali metal salt, because the metal cation won't dirrectly affect analytical processes, as it doesn't form precipitates.
Metals in this second group, also called alkaline earth metals, are also very reactive, but not as reactive as alkali metals. Towards lower periods, higher in the group, the reactivity is pretty low, with berylium and magnesium being able to form partially covalent bonds. Nevertheless, they are still more reactiv than hydrogen and react with acids to displace hydrogen from them.
The alkaline earth metals are (in increasing atomic number order): berylium (Be), magnesium (Mg), calcium (Ca), strontium, (Sr), barium (Ba) and radium (Ra). As in the case of francium, radium is not very well studied because its high radioactivity.
They are divalent only, with an \(ns^{2}\) outer shell structure.
They also react with water to form hydroxides, according to the reaction \(\ce{M + 2H2O \rightarrow M(OH)2 + H2}\). In the case of higher period more reactive metals, like calcium and barium, the reaction is as violent as it is for alkali metals, but in the case of less reactive ones, like magnesium, it requires heating.
Alkaline earth metals do not always form soluble compounds, with their hydroxides being slightly soluble in low concentrations and their carbonates and sulfates (apart from magnesium sulfate) insoluble.
Lower period elements only form normal oxides with the general formula MO. However, more reactiv barium is able to form the peroxide \(\ce{BaO2}\)
Metals in the p block are at the limit between metals and nonmetals. They have a very low electropositivity, showing a number of nonmetal characterstics like multiple valencies.
They have an \(ns^{2}np^{x}\) outer shell structure, 2+x being the main valence, but not the only one.
A lot of them form amphoteric oxides and hydroxides that are able to react with both acids and bases.
They have a low reactivity, however some of them are more reactive than hydrogen and are able to displace it from acids. Most of them don't react with water, only some forming an oxide layer upon heating (not a hydroxide).
A lot of their compounds are insoluble or only slightly soluble in water due to the partial covalent character of the bonds they form.
Some hydroxides, like aluminium or lead hydroxides, react with strong bases to form complex species. For example:
\[\ce{Al(OH)3 + NaOH \rightarrow Na[Al(OH)4]}\]
\[\ce{Pb(OH)2 + 2NaOH \rightarrow Na2[Pb(OH)4]}\]
This turns white insoluble hydroxides into soluble complex salts, esentially dissolving the hydroxides.
Despite their relatively low electropositivity (compared to other metal), p-block metals still form cations, ionic bonds and predominantly act as metals in chemical reactions, so they are still considered metals.
Written by Alex Jicu