Arrange the following elements in the order of the increasing electropositive character. Li, \(\mathrm{Na}, \mathrm{K}_{1} \mathrm{Rb}, \mathrm{Cs}\) (a) \(\mathrm{Li}>\mathrm{Na}>\mathrm{K}>\mathrm{Rb}>\mathrm{Cs}\) (b) \(\quad \mathrm{Li}<\mathrm{Na}<\mathrm{K}<\mathrm{Rb}<\mathrm{Cs}\) (c) \(\mathrm{Li}>\mathrm{Na}<\mathbb{K}<\mathrm{Rb}<\mathrm{Cs}\) (d) \(\mathrm{Na}>\mathrm{Li}>\mathrm{K}<\mathrm{Rb}<\mathrm{Cs}\)

Alkali metals hold a unique position on the periodic table. They are located in Group 1, which is the column on the far left side. This group includes lithium (Li), sodium (Na), potassium (K), rubidium (Rb), and cesium (Cs), among others. These elements are known for being highly reactive, particularly with water, where they form alkaline hydroxides - hence the name 'alkali'.

What distinguishes alkali metals from other elements is their single valence electron in the outermost shell. This lone electron is responsible for their chemical properties, including their high electropositivity. Electropositivity refers to the tendency of an element to lose its valence electron and form a positive ion, or cation. In aqueous solutions, alkali metals readily ionize, releasing this solitary electron.

Due to their high reactivity, alkali metals are never found in their elemental state in nature. Instead, they're commonly found in compounds, such as salts. A notorious characteristic of alkali metals is their density, which is low enough for some members of this group to float on water. This feature, coupled with their softness, distinctive flame colors when burned, and varying melting points, reflects the diversity within this group despite their shared chemical properties.

Understanding periodic table trends is crucial for predicting the chemical behavior of elements. As we move from left to right across a period (row), the atomic size generally decreases while the electronegativity and ionization energy tend to increase. In contrast, as we move down a group (column), the atomic size increases, and both electronegativity and ionization energy typically decrease.

These trends are a direct result of the structure of the atom and the forces at play within it. Electrons are added to the same valence shell as we move across a period, increasing the nuclear charge. This increased charge pulls the electrons in closer, reducing the atomic size. Down a group, a new electron shell is added with each subsequent element, which increases the distance between the outermost electrons and the nucleus, making it easier for the outer electrons to be lost and thus increasing electropositivity.

The trend in electropositivity is especially pronounced in the alkali metals. As mentioned earlier, the ease of losing that single valence electron makes them particularly electropositive. This trend is useful when predicting the reactivity of elements, especially within the same group, and is why alkali metals are more reactive as you move down the group.

The atomic size pertains to the distance from an atom's nucleus to the outer boundary of its electrons. This size can be measured in different ways, including atomic radius, ionic radius, and covalent radius. For our purposes, considering the atomic radius - the distance from the center of the nucleus to the outer edge of the electron cloud - is most relevant.

As we look down the Group 1 elements on the periodic table, the increase in atomic size is evident. Starting with lithium at the top and moving down to cesium, each successive element adds a new electron shell, making the radius of the atom larger. This increment in size corresponds to weaker forces holding the electrons to the nucleus and thus contributes to the ease with which these electrons can be donated to others – a key reason for the increased electropositivity down the group.

It is also worth noting that the atomic size affects other properties such as density, melting and boiling points, and ionization energy. For example, the increased size typically correlates with lower melting and boiling points in the alkali metals, which is why cesium has a significantly lower melting point than lithium.