What are Alkali Metals?
Alkali metals take up the leftmost side of the periodic table. The group 1 elements consist of elements:
They belong to the s-block elements of the periodic table as their outermost electron enters the s orbital giving them the electronic configuration of ns1.
As the alkali metals have only 1 electron in their valence shell, they readily lose it, making them count among the most reactive elements on earth. Thus, they are highly electropositive metals. They are called alkali metals because they form strongly alkaline hydroxides with water.
In this article, we will talk about the electronic configurations, ionization enthalpy, hydration enthalpy and atomic, ionic radii and other physical and chemical properties of the group 1 alkali metals.
Electronic Configuration of Alkali Metals
- Alkali metals have one electron in their valence shell.
- The electronic configuration is given by ns1.
For example, the electronic configuration of lithium is given by 1s2 2s1.
- They tend to lose the outer shell electron to form cations with charge +1 (monovalent ions).
- This makes them the most electropositive elements; and due to the same reason, they are not found in the pure state.
Atomic and Ionic Radii
- In a particular period, the alkali metals have the largest atomic sizes.
- The atomic radius of the elements increases down the group.
- When these elements lose an electron to form an ion, they lose an entire electron shell (since there is only one electron in the valence shell). Hence, the ionic radii are much smaller than the atomic radii.
- The ionic radius also increases down the group.
- In general, it is low for the alkali metals.
- The ionization enthalpies decrease down the group since the distance from the nucleus increases. Hence, it outweighs the increasing nuclear charge.
- When an element is brought into contact with water, new bonds are formed between water molecules and the ions. The energy released in the process is called the hydration enthalpy.
- As the ionic size increases, the hydration enthalpies decrease.
- Lithium has the highest ionization enthalpy. Hence, most of the lithium salts are hydrated.
The physical property of alkali metals:
- They are silvery white in color.
- They are soft and light metals.
- The density of these elements is generally low and increases while moving down from Li to Cs.
- They have only one electron in their outermost shell. Hence, they form weak bonds. This results in low melting and boiling point.
- They transmit characteristic color to an oxidizing flame. This happens because the heat provided by the flame excites the electron and it moves to a higher energy. When the excited electron moves back to the ground state, it imparts radiation in the visible region. This property of alkali metals helps in detecting metals from their respective flame test.
- The property of emission of light makes potassium and cesium useful in the making of electrodes of photoelectric cells.
The chemical property of alkali metals:
Alkali metals are highly reactive metals because of their large size and low ionization enthalpy. Their reactivity increases on moving down the group.
- Reactivity towards air: When these metals are kept in dry air, they get tarnished due to the formation of oxides, which become hydroxides on reacting with moisture. The reaction of alkali metals with oxygen is an exothermic reaction. The oxidation state +1.
Example: 4 Li + O2 → 2Li2O (oxide)
- Reactivity towards dihydrogen: Hydroxide and dihydrogen are formed when an alkali metal reacts with water.
Example: 2M + 2H2O → 2M+ + 2OH– + H2 (M = alkali metal)
- Reactivity towards dihydrogen: At 673 K, alkali metals reacts with dihydrogen to form hydrides.
Example: 2M + H2 → 2M + H–
- Reactivity towards halogen: Alkali metals react vigorously with halogens to form ionic halides M+X–.
- Reducing nature: Alkali metals are a strong reducing agent. Among them, lithium is the strongest, sodium being the least reducing agent.
- Solutions in liquid ammonia: Alkali metals are soluble in liquid ammonia and form a blue solution which is conducting in nature.
Example: M + (x+y)NH3 → [M(NH3)x]+ + [e(NH3)y]
The blue color of the solution is due to the ammoniated electron which absorbs the color from the visible light and imparts a blue color to the solution. The solution is paramagnetic and releases hydrogen to form amides on standing still for a long time.
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