The first element of alkali and alkaline earth metals differs in many respect from the other members of the group.
After studying this topic , you will be able to :
- Describe the general characteristics of the alkali metals and their compounds.
- Explain the general characteristics of the alkaline earth metals and their compounds.
- The biological significance of sodium , potassium , magnesium and calcium
The s-block is one of four blocks of element in the periodic table. The element of s-groups have a common property. The electrons in their most outwards electron shell are in the s-orbitals.
Elements in the s are in the first two periodic table groups. The elements in group one are called the alkali metals. The elements in group two are called the alkaline earth metals.
The modern periodic law says that " The properties of elements are periodic function of their atomic number " This means that some properties of elements are repeated as the atomic number of the elements gets larger.
These repeating properties have been used to separate the elements into four S . These S are S,P,D and F.
The members of this block lie on the extreme left of the periodic table.
The electrons present in an atom occupy various sub-orbitals of available energy levels in the order of increasing energy. The last electrons of an atom may find itself in either of the s,p,d,f subshells.
Accordingly, the elements of the atom having their last valence electron present in the sub shell.
What are S-Block Elements?
The s-block elements of the periodic table are those in which the last elections enters the outermost s-orbitals.
As, the s-orbitals can accomadate only two electrons (1&2) belong to the s-block of the periodic table.
Group I : group I of the periodic table consists of the elements Li, Na, K, Rb , Cs and Fr. They are so called because they form hydroxide on reaction with water which are strongly alkaline in nature.
Whereas , the s-block element having two electrons filling thier s-orbitals are called Group II or alkaline earth metal.
The elements of Group II includes Be, Mg , Ca, Sr, Ba, and Ra.
These elements with the exception of Be are commonly known as the alkaline earth metals.
These are so called because their oxides and hydroxide are alkaline in nature and these metal oxides are found in the earths crust .
Among the alkali metals sodium and potassium are abundant and Li , Rb and Cs have much lower abundance.
Fr is highly radioactive, its longest lived isotope. ²²³Fr has a half life of only 21 min , of the alkaline earth metals.
Ca and Mg rank fifth and sixth in abundance reoectively in the earths crust.
Sr and Ba have much lower abundance . Be is rare and Ra is the rarest of all comprising only 10^-10% of ignious rocks.
The general electronic configuration of s-block elements is (noble gas)ns1 for alkali metals and (noble gas)ns2 for alkaline earth metals.
Li and Be, the first elements of Group I and Group II respectively exhibit.
Some properties which are different from those of the other members of the respective group .
In these anamalous properties they resemble the second element of the following group. Thus, Li shows similarities to Mg and Be to Al in many of their properties.
This type diagonal similarity is commonly referred to as diagonal relationship in the periodic table.
The diagonal relationship is due to the similarity in ionic sizes and charge or radius ratio of the elements.
Monovalent Na and K ions and divalent Mg and Ca ions are found in large proportions in biological functions such as maintainance of ions balance and nerve impulse conduction.
Electronic Configuration of S-Block Elements :
Gruop I ; Alkali metals
The alkali metals show regular trends in physical and chemical properties with the increasing atomic Number.
The atomic and chemical and physical properties of alkali metals are discussed below :
The alkali elements in s-block consist of a single valence electron in their outermost shell. The outermost electron electron is loosely held which makes these metals highly electro positive.
They readily lose electrons to give monovalent M+ ions . Due to which they are not available in the free state in nature.
The general electronic configuration of s-blovk elements Group I are as shown below ;
Lithium (Li) — 1s2 2s1
Sodium (Na) — 1s2 2s2 2p6 3s1
Potassium (K) — 1s2 2s2 2p6 3s2 3p6 4s1
Rubidium (Rb) — 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s1
Cisium (Cs) — 1s2 2s2 2,p6 3s2 3p6 4s2 4p6 5s2 5p6 6s1
Francium(Fr) — 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s2 5p6 6s2 6p6 7s1
Group II : Alkaline Earth Metal ;
They follow alkali metals in the periodic table. These (except Be) are known as alkaline earth metals. The first element Be differs from the restvof the members and show diagonal relationship to Al.
These elements have two electrons in the s-orbitals of the valence shell. Their general electronic configuration may be represented as (noble gas)ns2 . Like alkali metal the compounds of these elrments are also predominantly ionic.
Berrylium (Be) — 1s2 2s2
Magnesium (Mg) — 1s2 2s2 2p6 3s2
Calcium (Ca) — 1s2 2s2 2p6 3s2 3p6 4s2
Strontium (Sr) — 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s2
Barium(Ba) — 1s2 2s2 2p6 3s2 3p6 4s2 4p6. 5s2 5p6 6s2
Radium(Ra) — 1s2 2s2 2p6 3s2 3p6 4s2 4p6 5s2 5p6 6s2 6p6 7s2
# Physical Properties of S-Block Elements :
Both alkali and alkaline earth metal elements show a regular graduation in their properties among their respective group elements.
All of the s-elements are metal (except H2) . In general, they are shiny, silvery, good conductor of heat and electricity.
They lose their valence electrons easily. In fact they lose their trademark s-orbitals valence electrons so easily that the s-elements are some of the most reactive elements on the periodic table.
Group I : The elements in GroupI , known collectively as the alkali metals (except H2) , always lose their one valence electron to make a +1 ion.
These metals are characterized by being silvery, very soft, not very dense and having low melting point.
These metals react extremely vigorously with water and even oxygen to produce energy and flammable H2 gas.
They are kept in mineral oil to reduce the chance of an unwanted reaction or H2 gas. They are kept in mineral oil to reduce the chance of an unwanted reaction or worse, answer explaination.
Physical Properties of Group I:
All the alkali metals are silvery white, soft and light metals. Because of the large size and these elements have low density which increases down the group from Li to Cs.
Potassium is lighter than sodium because sodium has high density than potassium.
The melting and boiling points of the alkali metals are low and they have weak metallic bonding due to the presence of only single valence electrons.
The alkali metals and their salts have characteristic colour flame because the heat excites the outermost orbital electron to a higher energy levels. When excited electron comes back to the ground state, there is emission of radiation in the visible region of the spectrum.
Li — Crimson red colour
Na — yellow colour
K — Violet colour
Rb — Red violet colour
Cs — Blue colour
Alkali metals colour detect by flame test and flame photometry or atomic absorption spectroscopy.
This property makes Caesium and Potassium useful as electrodes in photoelectric cells.
Group II :
The elements in Group II, known as the alkaline earth metals (except helium) , always lose their two valence electrons to make a +2 ion.
like the alkali metals, the alkaline earth metals are silvery, shiny and relatively soft.
Some of the elements in the column also react vigorously with water and must be stored carefully.
S-elements are famous for being ingredients in fireworks. The ionic forms of K, Sr and Ba make appearance in fireworks display as the brilliant purples, red and greens .
Fr is considered to be the most rare naturally occurring element on earth. It is estimated that there is only ever one natural atom of Fr present on earth at s time.
Fr has a very unstable nucleus and undergoes nuclear display rapidly.
Physical Properties of Group II:
The alkaline earth metals are silvery white, lustrous and relatively soft but harder than the alkali metals.
Berrilium and magnesium show greyish colour.
The melting and boiling points of these metals higher than alkali metals due to smaller sizes.
The electropositive character increases down the group from Be to Ba. Because ionisation energy decreases down the group.
Alkaline earth metals also show colour in flame.
Ca — Brick Red colour
Sr — crimson red colour
Ba — apple green colour
Atomic and Ionic radii :
Group 1 : When the s block elements of the modern periodic table are observed it is seen that the size of the alkali metals is larger compared to other elements in a particular period.
As the atomic number increases the total number electrons increases along with the addition of shells.
On moving down the group the atomic number increases. As a result, the atomic and ionic radius the alkali metals increases that is they increases in size while going from Li to Cs.
Group 2 : The atomic and ionic radii of the alkaline earth metals are smaller than those of the corresponding alkali metals in the same periods.
This is due to the increased nuclear charge
In these elements within the group, the atomic and ionic radii increases with increase in atomic number.
Ionisation Enthalpy :
As we go down the group the size of atoms inccreses due to which the attraction between the nucleus and the electrons in the outermost shell decreases.
As a result , the IE decreases. The IE of the alkali metals is comparatively lesser than other elements.
Group I : The IE of the alkali metals are considerably low and decrease down the group from Li to Cs.
This is because the effect of increasing size outweighs the increasing nuckear charge, and the outermost electron is very well screened from the nuclear charge.
Group II : The alkaline earth metals have low ionisation enthalpies due to fairly large size of the atoms.
Since the atomic size invreases down the group, their IE decreases . The first IE of the alkaline earth metals are higher than those of the corresponding Group I metals.
This is due to their small size as compared to the corresponding alkali metals. It is interesting to note that the second IE of the alkaline earth metals are smaller than those of the corresponding alkali metals.
Hydration Enthalpies :
As the ionic sizes of the elements increase , the hydration enthalpy decrease. Smaller the size of the hydration enthalpy is high as the atomic has the capacity to accommodate a larger no of water molecules around it due to high charge /radius and hence gets hydrated.
Group I :
The hydration enthalpies of alkali metals ions decrease with increase in ionic sizes.
Li+>Na+> K+> Rb+> Cs+
Li+ has maximum degree of hydration and for this reason lithium salts are mostly hydrated.
Example : LiCl. 2H2O
Group II : Like alkali metals ions, the hydration enthalpies of alkaline earth metal ions decrease with increase in ionic size down the group.
Be2+>Mg2+>Ca2+>Sr2+>Ba2+
The hydration enthalpy of alkaline metals ions are larger than those of alkali metal ions. Thus, compounds of alkaline earth metals are more extensively hydrated than those of alkali metals.
Example : MgCl2 and CaCl2, exist as MgCl2. 6H2O and CaCl2. 6H2O while NaCl and KCl do not form such hydrates.
Density :
GroupI : All are light metals. The densities are low Life, Na and K are lighter then water and for this very reason, they float on water. Density increases while moving from Li to Cs. K is lighter than Na.
Group II : These metals are denser than alkali metals in the same period because these can be packed more tightly due to their greayer nuclear charge and smaller size.
The densities decrease slightly up-to Ca and then increasez considerably up-to Ra. Irregular trend is due to the difference in the crystal structure of these elements.
Melting and boiling points :
GroupI : The energy binding the atoms in the crystal lattices of these metals is relatively low on account of a single electron in the valence shell.
Group II : The melting and boiling points of these elements are higher than the corresponding alkali metals.
This is due to the presence of two electrons in the valence shell and thus, they are strongly bonded in the solid state.
However, Melting and boiling points do not show any regular trends because atoms adopt different crystal structures.
Chemical Properties :
Group I : The alkali metals are highly reactive due to their large size and low IE.
The reactivity of these metals increases down the group.
(1) Reactivity towards air : The alkali metals tarnish in day due to the formation of their oxides which is turn react with moisture to form hydrates.
They burn vigorously in oxygen forming oxides. Like forms monoxide, Na forms peroxide, the other metals form superoxide.
The superoxide O2- ion is stable only in the presence of large cations such as K, R, Cs
4Li + O2 --------------> 2Li2O (oxide)
2Na + O2 ------------> Na2O2 (peroxide)
M + O2 ----------------> MO2 (superoxide)
(M = K, R, Cs)
In all these oxides the oxidation states of the alkali metals is +1. Like shows exceptional behaviour in reacting directly with Nitrogen of air to form the nitride, Li3N as well. Because of their high reactivity towards air and water, alkali metals are normally kept in kerosene oil.
(2) Reactivity towards water : Form hydroxide and dihydrogen.
2M. + 2H2O ----------------> 2M+ + 2OH- + H2
(M = an alkali metals)
Li has very high negative value of electrode potential, it's reaction with water is less vigorous than sodium because Na has less negative electrode potential value.
Other metals of the group react explosively with water.
They also react with proton sonata such as alcohol, gaesous NH3 and alkynes.
(3) Reactivity towards dihydrogen : The alkali metals react with dihydrogen at about 673K to form hydrides.
All the alkali metal hydrides are ionic solids with high Melting point.
(4) Reactivity towards halogens : The alkali metals readily react vigorously with halogens to form ionic halodes M+X-
Like halodes are covalent in nature beacuse it has high polarisation capability of Li ion (The distortion of electrons cloud of the anion by the cation is called polarisation).
The Li+ ion is very small in size and has high tendency to distort electron cloud around the negative halodes ion.
Since anion with large size can be easily distorted, among halodes LiI is the most covalent in nature.
(5) Reducing Nature : The alkali metals are strong reducing agent.
Li is the high reducing agent because it has highly negative electrode potential value.
(6) Solutions in Liquid ammonia : When alkali metals dissolve in liquid ammonia, it gives deep blue solution which flow electricity that is conducting in nature.
M + (x +y) NH3 ---------------> [M(NH3)x] ^+ or. [e(NH3)y] ^-
The blue colour of the solution is because of ammoniated electron which absorbs energy from visible light hence it emit radiation and show blue colour.
The solution is paramagnetic in nature because it release free electron. But in concentrated solution the blue colour changes to bronze colour and becomes diamagnetic.
M+ + e- + NH3 -------------> MNH3 + 1/2H2(g)
Uses of alkali metals :
- Li metal is used to make useful alloys.
- Li is used in thermonuclear reaction.
- Li is also used to make electrochemical cells.
- Sodium is used to make a Na/Pb alloy to make PbEt4 and PbMe4.
- Li Na metal is used as a coolant in fast breeder nuclear reactors.
- K has a vital role in biological system.
- KCl is used as a fertilizer.
- KOH is used in the manufacture of soft soap.
- KOH is also used as an excellent absorbent of CO2.
- Cs is used in devising photoelectric cells.
Chemical properties of Alkaline Earth metals :
Alkaline earth metals are less reactive than alkali metals.
Reactivity increases going down the group.
(1) Reactivity towards air and water: Be and Mg are kinetically inert to oxygen and water because of the formation of an oxide film on their surface.
(2) Reactivity towards the halogens : All the alkaline earth metals combine with halogens at elevated temperature forming their halodes.
M + X3 -------------> MX3 (X = F, Cl, Br, I)
(3) Reactivity towards hydrogen : All the elements except Be combine with hydrogen upon heating to form their hydrides MH2.
BeH2 can be prepared by the reaction of BeCl2 with LiAlH4.
2BeCl2 + LiAlH4 ------------> 2BeH2 + LiCl + AlCl3
(3) Reactivity towards acids :The alkaline earth metals readily react with acids liberating dihydrogen.
M + 2HCl -------------> MCl2 + H2
(4) Reducing Nature : Alkaline earth metal are strong reducing agent. Because alkaline earth metal has high negative potential value of their reduction potential.
(5) Solutions in liquid ammonia : Alkali metals, the alkaline earth metals dissolve in liquid ammonia to give deep blue black solutions forming ammoniated ions.
M + (x+y)NH3 ----------> [M(NH3)x]2+ or. 2[e(NH3)y]-
Uses of Alkaline earth metal :
- Be is used in the manufacture of alloys.
- Cu-Be alloys are used in the preparation of high strength springs.
- Metallic Be is used for making windows of X-Ray tubes.
- Me is used in flash powders and bulbs, incendiary bombs and signals.
- Radium salts are used in radiotherapy.
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