Give reason for the following: (a) Group 1 elements are strong reducing agents but poor complexing agents. (b) Caesium is used in photoelectric cells. (c) Sodium metal is kept under kerosene oil. (d) Lithium shows anomalous behaviour. (e) LiF is insoluble in water.

Group 1 elements, often known as alkali metals, play a pivotal role as reducing agents in various chemical reactions. Reducing agents are substances that donate electrons to other materials during reactions, leading to the reduction – or gain of electrons – of these materials. Alkali metals are particularly adept at this because they possess a single electron in their outermost shell, which they can readily lose to achieve a more stable, noble gas electronic configuration.

When an alkali metal atom loses an electron, it becomes a positively charged ion, effectively reducing another compound by transferring its electron. This electron donation is what makes alkali metals potent reducing agents. Notably, this ability to lose an electron easily also contributes to their inability to act as complexing agents, given that complexation often requires elements to share or accept electrons, something that alkali metals, with their eager electron shedding, are not predisposed to.

Caesium's utilization in photoelectric cells owes to its distinctive atomic properties. Photoelectric cells are devices that convert light energy into electrical energy through the photoelectric effect. In this process, when photons strike the surface of a material like Caesium, they can confer enough energy to dislodge electrons from the material's surface.

Caesium is particularly suitable for this because it has one of the lowest ionization energies among the elements, which means it takes very little energy to remove an electron from a Caesium atom. As a result, when exposed to light, Caesium easily releases electrons, thereby allowing for efficient conversion of light to electricity in photoelectric cells. This makes Caesium an ideal material for applications that rely on the photoelectric effect, such as light meters, photovoltaic cells, and various types of sensors.

Storing alkali metals, like Sodium, requires careful measures due to their high reactivity. Sodium, for example, reacts quite vigorously with both oxygen and water, forming sodium oxide and generating heat and potentially flammable hydrogen gas. To prevent these dangerous reactions, sodium is stored under kerosene oil.

The kerosene layer acts as a barrier, protecting the sodium metal from coming into contact with air and moisture. This precaution is critical not just for safety, but also to preserve the integrity of the metal for use in future applications. The principles of safe storage and handling of reactive alkali metals are foundational knowledge for anyone studying or working with these elements.

Lithium, the lightest of the alkali metals, exhibits anomalous behaviour not entirely in line with its Group 1 counterparts. This divergence stems from its comparatively small atomic and ionic sizes, high electronegativity, and high ionization energy. These features contribute to Lithium's tendency to form covalent bonds rather than ionic ones, which is the norm among alkali metals.

Lithium's distinctive properties also result in unique reactions and applications that differ from other group members. For instance, its covalent nature means it interacts differently with molecules and forms distinct types of compounds. Understanding Lithium's idiosyncrasies is essential in grasping the broader reactivity and bonding trends within the periodic table's first group of elements.

The solubility of compounds in water is a key concept in chemistry, impacting everything from biological processes to industrial applications. For instance, LiF's insolubility in water contrasts with the general trend of increasing solubility of group 1 metal salts. The explanation behind LiF's behaviour lies in its high lattice enthalpy, which is the energy required to break the bonds between its ions.

Due to the strong ionic bonds between Lithium ions and Fluoride ions, a significant amount of energy is needed to disrupt these interactions and dissolve LiF in water—an amount of energy that the solvation process does not provide. Understanding the factors that influence solubility, such as lattice enthalpy and hydration energy, is crucial when predicting the behaviour of salts in various environments and designing processes where dissolution or precipitation of materials is involved.