Chemistry Alkali metals and their properties

Topics to be covered

`color{green}(=>)` General introduction
`color{green}(=>)` The periodic table
`color{green}(=>)` Group 1 :alkali metals
`color{green}(=>)` Electronic configuration
`color{green}(=>)` Atomic and ionic radii
`color{green}(=>)` Ionization enthalpy
`color{green}(=>)` Hydration enthalpy
`color{green}(=>)` Physical properties


`color{green}( ★)` The s-block elements of the Periodic Table are those in which the last electron enters the outermost s-orbital.

`color{green}( ★)` As the s-orbital can accommodate only two electrons, two groups (1 & 2) belong to the s-block of the Periodic Table.

`color{green}( ★)` Group 1 of the Periodic Table consists of the elements: lithium, sodium, potassium, rubidium, caesium and francium.

`color{green}( •)` They are collectively known as the alkali metals.

`color{green}( •)` These are so called because they form hydroxides on reaction with water which are strongly alkaline in nature.

`color{green}( ★)` The elements of Group 2 include beryllium, magnesium, calcium, strontium, barium and radium.

`color{green}( •)` These elements with the exception of beryllium are commonly known as the alkaline earth metals.

`color{green}( •)`These are so called because their oxides and hydroxides are alkaline in nature and these metal oxides are found in the earth’s crust*.

`color{green}( ★)` Among the alkali metals sodium and potassium are abundant and lithium, rubidium and caesium have much lower abundances (Table 10.1).

`color{green}( ★)` Francium is highly radioactive; its longest-lived isotope `color{red}(text()^(223)Fr)` has a half-life of only 21 minutes.

`color{green}( ★)` Of the alkaline earth metals calcium and magnesium rank fifth and sixth in abundance respectively in the earth’s crust.

`color{green}( ★)` Strontium and barium have much lower abundances. Beryllium is rare and radium is the rarest of all comprising only 10^–10 per cent of igneous rocks.

`color{green}( ★)` The general electronic configuration of `color{red}(s)`-block elements is [noble gas] `color{red}(ns^(1))` for alkali metals and [noble gas] `color{red}(ns^(2))` for alkaline earth metals.

`color{green}( ★)` Lithium and beryllium, the first elements of Group 1 and Group 2 respectively exhibit some properties which are different from those of the other members of the respective group.

`color{green}( •)` In these anomalous properties they resemble the second element of the following group.

`color{green}( •)` Thus, lithium shows similarities to magnesium and beryllium to aluminium in many of their properties.

`color{green}( •)` This type of diagonal similarity is commonly referred to as diagonal relationship in the periodic table.

`color{green}( •)` The diagonal relationship is due to the similarity in ionic sizes and /or charge/radius ratio of the elements.

`color{green}( •)` Monovalent sodium and potassium ions and divalent magnesium and calcium ions are found in large proportions in biological fluids. These ions perform important biological functions such as maintenance of ion balance and nerve impulse conduction.

Periodic table


`color{green}( ★)` The alkali metals show regular trends in their physical and chemical properties with the increasing atomic number. The atomic, physical and chemical properties of alkali metals are discussed below.

Electronic Configuration

`color{green}( •)` All the alkali metals have one valence electron, `color{red}(ns^1)` (Table 10.1) outside the noble gas core.

`color{green}( •)` The loosely held `color{red}(s)`-electron in the outermost valence shell of these elements makes them the most electropositive metals.

`color{green}( •)` They readily lose electron to give monovalent `color{red}(M^+)` ions. Hence they are never found in free state in nature.

Atomic and Ionic Radii

`color{green}( ★)` The alkali metal atoms have the largest sizes in a particular period of the periodic table.

`color{green}( •)` With increase in atomic number, the atom becomes larger.

`color{green}( •)` The monovalent ions `color{red}((M^+))` are smaller than the parent atom.

`color{green}( •)` The atomic and ionic radii of alkali metals increase on moving down the group i.e., they increase in size while going from `color{red}(Li "to" Cs)`.

Ionization Enthalpy

`color{green}( ★)` The ionization enthalpies of the alkali metals are considerably low and decrease down the group from `color{red}(Li "to" Cs)`.

`color{green}( •)` This is because the effect of increasing size outweighs the increasing nuclear charge, and the outermost electron is very well screened from the nuclear charge.

Hydration Enthalpy

`color{green}( ★)` The hydration enthalpies of alkali metal ions decrease with increase in ionic sizes.

`color{red}(Li^+ > Na^+ > K^+ > Rb^+ > Cs^+)`

`color{red}(Li^+)` has maximum degree of hydration and for this reason lithium salts are mostly hydrated, e.g., `color{red}(LiCl . 2H_2O)`

Physical Properties

`color{green}( ★)` All the alkali metals are silvery white, soft and light metals.

`color{green}( •)` Because of the large size, these elements have low density which increases down the group from `color{red}(Li "to" Cs)`.

`color{green}( ★)` However, potassium is lighter than sodium.

`color{green}( ★)` The melting and boiling points of the alkali metals are low indicating weak metallic bonding due to the presence of only a single valence electron in them.

`color{green}( ★)` The alkali metals and their salts impart characteristic colour to an oxidizing flame.

`color{green}( •)` This is because the heat from the flame excites the outermost orbital electron to a higher energy level. When the excited electron comes back to the ground state, there is emission of radiation in the visible region as given below:

`color{green}( ★)` Alkali metals can therefore, be detected by the respective flame tests and can be determined by flame photometry or atomic absorption spectroscopy. These elements when irradiated with light, the light energy absorbed may be sufficient to make an atom lose electron. This property makes caesium and potassium useful as electrodes in photoelectric cells.