Sodium reacts with water to yield hydrogen gas. Why is this reaction not used in the laboratory preparation of hydrogen?

When sodium, an alkali metal, is dropped into water, it reacts vigorously to produce sodium hydroxide and hydrogen gas. This type of reaction is a hallmark of the alkali metal group due to their keen ability to donate electrons. The reactivity of sodium can be pinned down to its electronic configuration, which has a single electron in its outermost shell. This makes sodium atoms unstable and highly eager to react with other substances, particularly water, in order to achieve a more stable electron configuration.

The balanced chemical equation for this reaction is: \[2Na (s) + 2H_2O (l) \rightarrow 2NaOH (aq) + H_2 (g)\].Here, \((s)\) stands for solid, \((l)\) for liquid, \((aq)\) for aqueous or dissolved in water, and \((g)\) for gas. The equation confirms that for every two moles of sodium reacting, one mole of hydrogen gas is produced.

However, the sodium water reaction is not merely a spectacle of fizz and bubbles; it is also quite dangerous. Sodium's eagerness to react makes the process highly vigorous, often leading to splattering of the sodium hydroxide produced, and the rapid release of hydrogen gas poses a risk of explosion if not conducted in a controlled environment.

Alkali metals are known for being some of the most reactive metals in the periodic table. What makes this group outstanding in terms of reactivity is their electron configuration; they all have one electron in their outermost shell, which they are inclined to lose in order to reach a noble gas configuration, resulting in a stable fully-filled shell or an empty shell after the electron has been lost. Among the alkali metals, reactivity increases further down the group, meaning lithium reacts less vigorously than sodium, and sodium less so than potassium.

The high reactivity is not just an academic curiosity—it has practical implications. It requires that alkali metals like sodium must be stored under oil to prevent reactions with water vapor in the air. In a laboratory setting, their high reactivity can make them difficult to handle safely and predictably. Safety protocols must be strictly adhered to when working with these metals to prevent potentially violent reactions, especially when they come into contact with water or other substances that can facilitate a rapid release of electrons.

An exothermic reaction is a chemical reaction that releases energy by light or heat. It is the opposite of an endothermic reaction, which consumes energy. Exothermic reactions are generally more spontaneous than endothermic reactions. In the case of sodium's interaction with water, the reaction is so exothermic that the generated heat is often enough to ignite the produced hydrogen gas.

Diverse exothermic reactions are everywhere in our daily lives, from the combustion of fuels for power to the heat released when rust forms on iron. In a controlled environment such as a chemistry lab, manipulating the conditions like temperature, concentration, or surface area can influence the rate and the output of exothermic reactions. Despite the usefulness of such reactions, the immense amount of energy released in the sodium water reaction, for example, can be hazardous, as it may lead to explosions if the hydrogen gas catches fire. This potential danger is a primary reason why less volatile methods are preferred for generating hydrogen gas in laboratory settings.