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What are Alkali Metals? – Definition
The word “alkali has been derived from Arabic word ‘ Alquili ‘ meaning the ashes of plants from which certain compounds of the elements sodium and potassium were initially isolated.
Group 1 elements readily dissolve in water to form soluble hydroxides which are strongly alkaline in nature.
The alkali metals are also called s-block elements. Since the elements listed below have one electron each in the valence s-subshell of their atoms i.e., they have ns1 configuration (n represents the valence shell).
Alkali Metals Examples
All elements are silvery-white, soft and light metals. Group IA elements are metals because they have low ionization energies and have a few valence electrons as compared to available vacant orbitals. These are highly malleable (can be pressed out into sheets) and ductile (can be drawn out into wires). When freshly cut, they have a bright luster which quickly tarnishes on exposure to air.
The similarity in the electronic configuration results in similar alkali metals properties.The group 1 of the periodic table comprises six elements :
- lithium (Li)
- Sodium (Na)
- Potassium (K)
- Rubidium (Rb)
- Cesium (Cs) and
- francium (Fr).
Alkali Metals Periodic Table
- Besides hydrogen which is not a member of the alkali metal family. Hydrogen has been included because of the similarity in the electronic configuration with these elements. These are called alkali metals.
- Last element francium is of radioactive nature and is rather unstable with a very small life period of 21 minutes. Therefore, very little information is available about this element.
Alkali Metals Uses
- Alkali metals are highly malleable (can be pressed out into sheets) and ductile (can be drawn out into wires).
- Owing to its very low density, it floats to the surface of kerosene. Hence, uses for an application where low density required.
- They are diamagnetic and colorless.
- Due to their strong electropositive nature, they emit electrons even when exposed to light (photoelectric effect). This property is responsible for their use in photoelectric cells; cesium and potassium are used, in particular, for this purpose.
- Due to the presence of loosely held valence electrons which are free to move throughout the metal structure, the group 1 elements are good conductors of heat and electricity.
Characteristics of Alkali Metals – Properties
|Atomic Radius – Increases from Lithium (Li) to Francium (Fr). Electronegativity – Decreases down to Group. Electropositive- Increase Down to Group. Ionization energy decreases on moving down from Li to Cs. Melting and Boiling points decrease on moving down the group from Li (m.p. 186°C) to Cs (m.p. 28.5°C). The hydration energy of alkali metal ions decreases from Li+ to Cs+. Reaction with air – Alkali metals cannot be placed in air.|
Alkali Metals Reactivity
Alkali Metals reactivity increase with the increase of atomic radius.
Cs> Rb> K > Na> Li(Atomic radius)
Atoms of the group 1 elements are the largest in their corresponding periods. Atomic, as well as ionic size, increases from Li to Fr due to the presence of an extra shell of electrons. Atomic volume (At. wt/Density) also increases in moving down from Li to Cs.
Behaviour with water- Smaller the size of a cation, greater is its charge density and hence greater is its tendency to draw electrons from molecules which are thus polarized. Lithium-ion, being smallest in size among alkali metal ions, is the most extensively hydrated while Cs+ ion, the largest alkali metal ion, is the least hydrated.
Cs+ > Rb+ > K+ > Na+ > Li+ (Relative ionic radii)
Li+ > Na+ > K+ > Rb+ >Cs+ (Relative ionic radii in water) (Relative degree of hydration)
Lithium-ion, being heavily hydrated, moves very slowly under the effect of electric current and is thus the poorest conductor of electricity as compared to other alkali metal ions. Thus it is the degree of hydration of the ions rather than their size that determines the electrical conductivity of the alkali salt solutions. According to electrical conductivity measurements, the alkali metal ions conduct electric current in the following order.
Cs+ > Rb+ > K+ > Na+ > Li+ (Relative electrical conductivity)
Alkali Metals in Water
All alkali metal salts are ionic (except Li) and soluble in water, a solvent of high dielectric constant. The solubility in water is due to the fact that the cations get hydrated by water molecules. The degree of hydration depends on the size of the cation.
Reaction with water (Formation of hydroxides). Alkali metals, their oxides, peroxides and even superoxides dissolve in water to form hydroxides which are soluble and are called alkalies (water soluble hydroxides are known as alkalies).
- 2Na + H2O →2NaOH + H2
- Li2O + H2O→ 2LiOH
- Na2O2 + 2H2O→ 2NaOH+ H2O2
- 2KO2+ 2H2O →2KOH + H2O2 + O2
The reactions of metals with water are so highly exothermic that the hydrogen gas evolved catches fire accompanied by an explosion. Therefore, alkali metals are not kept in contact with water.
Physical Properties of Alkali Metals
Their densities are quite low and increase from lithium to cesium.However, potassium is lighter than sodium (anomaly) probably due to an unusual increase in an atomic size of potassium. Lithium, sodium, and potassium are lighter than water; Lithium is the lightest known metal (density 0.534). Owing to its very low density, it floats to the surface of kerosene, hence it can’t be stored in kerosene. It is kept wrapped in paraffin wax.
Since these metals are highly electropositive, their electronegativity (i.e. tendency to attract electrons) values are very low. Further, since electropositive character increases on moving down the group, the electronegativity decreases in the same order, i.e. from Li to Cs.
Due to their large size, the outermost solitary s-electron is at a large distance from the nucleus and, therefore, can be easily removed. Thus their ionization energies are low.
Further, as the atomic radius increases on moving down the group the outer electron gets farther and farther away from the nucleus and, therefore, ionization energy decreases on moving down from Li to Cs.
The alkali metal atoms show only +1 oxidation state. Because of their low ionization energies, they easily lose the outermost S-electron to form the unipositive ions. Since the unipositive ions have the stable noble gas configuration (s2 Or s2p6) in the valence shell, the energy required to pull out another electron from the valence shell is very high. Hence the second ionization energies of alkali metals are very high. Therefore, we can say that group 1 elements are univalent and form ionic compounds.
Further, since the alkali metal ions have a noble gas configuration with no unpaired electron, they are diamagnetic and colorless. In fact, all the compounds of alkali metal ions are colorless except those where the anion is colored, viz, para-manganates and dichromates.
Electropositive character of Alkali Metals
It is the tendency of the element to lose an electron. On account of their low ionization energies, these metals have a great tendency to lose the ns1 electron and form positive ions.
M → M+ + e–
Thus group 1 elements are strongly electropositive (or metallic in nature). Further, since ionization energy decreases from Li to Cs, the electropositive character increases in going down from Li to Cs.
Melting and boiling points.
The melting and boiling points are very low because of the weak bonding in the crystal lattice of the metals (weak bonding in the crystal lattice also explains the softness of alkali metals). The weak interatomic bonds (binding energies) is due to their large atomic radii and especially due to the presence of only one valence electron per metal atom as compared to a large number of available vacant orbitals.
With the increase in the size of the metal atoms, the repulsion of the non-bonding electrons increases and therefore melting and boiling points decrease on moving down the group from Li (m.p. 186°C) to Cs (m.p. 28.5°C).
Flame Test Showing Alkali Metals Properties
The alkali metals and their salts, when introduced into the flame, give characteristic color to the flame.
|Crimson red||Golden yellow||Pale violet||Violet||Violet|
This property of the group 1 elements offers a very sensitive and reliable test (flame test) for alkali metals which are difficult to be identified by chemical methods as they do not form many insoluble compounds.
The reason for flame colouration is that when an alkali metal or its any salt is introduced into the flame, the outermost electrons of the alkali atoms absorb energy and excited to the higher energy levels. When the excited electrons return to their original (ground) level, they release the absorbed (excited) energy as visible light. Now for the same excitation energy, the energy level to which the electron in Li will rise is lower than that to which the electron in Na will rise and this, in turn, is lower than the level to which the electron in K will rise and so on.
These differences are due to differences in their ionization energies. Consequently, when the electron returns to the ground state, the energy released will be lowest in Li+ and will increase in the order: Li+ , Na+ , K+ , Rb+ , and Cs+ . As a result of this, the frequency of the light emitted in the Bunsen flame is minimum in lithium and corresponds to the red region of spectra. In potassium, the frequency of the light emitted corresponds to the violet region of spectra.
Hydration of ions is an exothermic process. The energy released in the hydration of ions is known as hydration energy. Since the degree of hydration of M+ ions decreases as we go down the group, the hydration energy of alkali metal ions decreases from Li+ to Cs+ .
Chemical Properties of Alkali Metals
Group 1 elements are highly reactive chemically because of their low ionization enthalpies and enthalpy of atomization. Some of the important chemical properties of the members of the family are discussed.
Reaction with air
When freshly cut, alkali metals have luster but their surfaces get tarnished when exposed to air due to the formation of a layer of oxide, hydroxide, and carbonate.
- 4M + O2 → 2M2O (M=Metal)
- M2O + H2O → 2MOH
- 2MOH + CO2 → M2CO3 + H2O
The alkali metals cannot be placed in air. Similarly, these are also not placed in water due to strong affinity. These are normally kept in chemically inert solvents such as kerosene.
Reaction with oxygen
Alkali metals combine with oxygen upon heating to form different oxides depending upon their nature. Lithium forms a normal oxide (Li2O), sodium forms peroxide (Na2O2) while potassium and rest of the metals form superoxides (MO2 where M = K, Rb and Cs) upon heating in oxygen.On Heating,
- 4Li + O2 →2Li2O (Lithium monoxide) Oxide ion: O2-
- 2Na + O2 →Na2O2 (Sodium Peroxide) Peroxide : O22-
- K + O2 → KO2 (Potassium superoxide) Superoxide ion: O2
Reactivity with oxygen increases from Li to Cs
Reaction with hydrogen
All alkali metals combine with hydrogen upon heating to form colorless crystalline hydrides which are of ionic nature.
- 2M + H2 →2M–H+
(M=Li, Na, K. Rb, Cs)
The ionic character of the hydrides increases from Li to Cs.
As the alkali metals have low ionization enthalpies, their atoms can easily lose valence electron to a hydrogen atom and form ionic hydrides (M– H+ ). Since the ionization enthalpy decreases down the group, the tendency to form positive ion increases accordingly. Therefore, the ionic character of hydrides also increases.
Reaction with halogens
Group 1 elements(M) combine with halogens(X) directly to form metal halides.
- 2M + X2 → 2MX
With the exception of certain lithium halides, the halides of rest of the metals are of ionic nature (M+X–). They have high melting and boiling points. The fused halides are good conductors of electricity and in the fused state, these are used for the preparation of the alkali metals.
Order of reactivity of M : Li < Na < K < Rb < Cs
Order of reactivity of X2 : F2>Cl2 > Br2 > I2
Reaction with Sulphur and Phosphorus
Alkali metals react with sulphur and phosphorus upon heating to form the corresponding sulphides and phosphides as follows:
- 16 Na + S8 → 8Na2S (Sod. sulphide)
- 12Na + P4→ Na3P (Sod. phosphide)
Both sulphides and phosphides are hydrolysed by water as follows:
- Na2S + H2O→ NaOH + NaHS
- Na3P + 3H2O → 3NaOH + PH3
Alkali Metals With Ammonia
Solubility in liquid ammonia. Group 1 elements dissolve in liquefied ammonia to give highly conducting solution with blue color.
Alkali metals due to their low ionization energies, ionize in the ammonia solution to form ammoniated cations and ammoniated electrons.
- M + (x + y) NH3 → (M(NH3)x) + (e(NH3)–
The blue color of the solution is attributed to the fact that when light falls on the ammoniated electrons, they absorb energy corresponding to the red color and the transmitted light has a blue color. The electrical conductivity of the solution is because of ammoniated cations as well as ammoniated electrons.
(i) In concentrated solution, the color changes from blue to bronze. The blue solutions are paramagnetic while the concentrated solutions are diamagnetic in nature.
(ii) The presence of free ammoniated electrons makes the blue solution reducing in nature.
(iii) When dry ammonia gas is passed over heated metals, amides are formed and hydrogen gas is evolved.
- 2M + 2NH3 → 2MNH2(metal amide) + H2
Alkali metal amides are powerful reducing agents due to the presence of amide (NH2) ions.
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