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S-Block Element (Alkali And Alkaline Earth Metals)

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s Block Element PDF Notes, Important Questions and Formulas

S - Block Elements (Alkali and Alkaline Earth Metals)

The elements in which the last electron enters the outermost s-orbital are called s-block elements. The group 1 and 2 of periodic table belong to the s-block.

The s-block elements of the periodic table are those in which the last electron enters the outermost sorbitol. 

  • Elements of group 1 are called alkali metals, and elements of group 2 are called alkaline earth metals.
  • The alkali and alkaline earth metals are the most violently active of all the metals.
  • These metals are not found in the Free State in nature. They are found in the combined form with halide, sulphate, carbonate, silicate ions etc.
  • Group 1 elements: Alkali metals
  • Group 1 elements are called alkali metals because they form hydroxides on reaction with water and are strongly alkaline in nature.
  • The general electronic configuration for alkali metals is [noble gas] ns1.

 Trends in properties of alkali metals: 

  • Atomic and ionic radii: Alkali metals have the largest atomic and ionic radii in their respective periods of the periodic table. On moving down the group, the atomic and ionic radii increase.
  • Ionisation enthalpies: Alkali metals have the lowest ionisation enthalpy in each period. Within the group, the ionisation enthalpies of alkali metals decrease down the group. The second ionisation enthalpies of alkali metals are high.
  • Melting and boiling points: Alkali metals are soft and have low melting and boiling points.
  • Density: Densities of alkali metals are quite low as compared to other metals. The densities increase on moving down the group. However, K is lighter than Na.
  • Electropositive or metallic character: All the alkali metals are strongly electropositive or metallic in character.
  • The electropositive character of an element is expressed in terms of the tendency of its atom to release electrons:
  • M M e + - → + CHEMISTRY THE s-BLOCK ELEMENTS
  • Oxidation states: All the alkali metals predominantly exhibit an oxidation state of +1 in their compounds.
  • Characteristic flame colouration: All the alkali metals and their salts impart characteristic flame colouration. Characteristic flame colouration by different alkali metals can be explained on the basis of difference in amount of energy absorbed for excitation of the valence electron.
  • Photoelectric effect: The phenomenon of ejection of electrons when electromagnetic radiation of suitable frequency strikes a metal surface is called photoelectric effect. Alkali metals exhibit photoelectric effect because they emit electrons when radiation strikes their surfaces
  • Chemical Properties:
    • Reactivity towards air or oxygen: When heated in excess of air or oxygen, they bum vigorously forming different types of oxides. Lithium forms monoxide, sodium forms peroxide and the other elements form superoxides.
    • Reactivity towards water: Alkali metals reacts with water to form metal hydroxides (MOH) and hydrogen gas is evolved.
    • Reactivity towards hydrogen: Alkali metals react with hydrogen to form hydrides which are ionic in nature (M+H-). The stability of hydrides decreases from Li to Cs.
    • Reactivity towards halogens: These metals readily combine with halogen to form ionic halides.
    • Hydration of alkali metal ions: Alkali metal ions are highly hydrated. The smaller the size of the ion, the greater is the degree of hydration. Thus, Li+ ion gets more hydrated than Na+ ion, which is more hydrated than K+ ion and so on. As a result of larger hydration of Li+ ion than Na+ ion, the effective size of hydrated Li+ ion is more than that of hydrated Naion. Hydrated ionic radii decrease in the order: Li+ > Na>K>Rb+ > Cs+
    • Due to extensive hydration, Li+ ion has the lowest mobility in water.
    • Reducing nature: Alkali metals are strong reducing agent. The tendency to act as a reducing agent lose electrons. Lithium is the strongest reducing agent. The tendency to act as a reducing agent depends on energy requirement involved in three processes, i.e. sublimation, ionisation and hydration.
    • M(s)→M(g)       sublimation enthalpy
    • M(g)→M+(g) +e-     ionisation enthalpy
    • M+(g)+H2O→M+(aq) hydration enthalpy
    • Solution in liquid ammonia: Alkali metals dissolve in liquid ammonia giving deep blue solutions which are conducting in nature. In solutions, the alkali metal atom readily loses the valence electron. Both the cation and the electron combine with ammonia to form ammoniated cation and ammoniated electron. The blue colour of the solution is due to the ammoniated electron which absorbs energy in the visible region of light.
    • M+(x + y)NH3→[M(NH3x]++[e(NH3)y]-
    • Basic strength of hydroxides: Hydroxides of alkali metals are strongly basic, and the basic strength increase down the group.

Group-I Elements: (Alkali metals)

  1. The Elements: are Li, Na, K, Rb, Cs, Fr
    (Radioactive: t1/2 of Fr233=21 minute)
    Group-I elements are called alkali metals because they form hydroxides on reaction with water, which are alkaline in nature.
  2. Outer Electronic Configuration: ns1|
  3. Atomic and Ionic radii
    Li< Na< K< Rb< Cs.
  4. Density
    Li>Na>K>Rb>Cs.
  5. Ionization Energy
    Li>Na>K>Rb>Cs.
    As size increases, I.E. decreases down the group (so Cs have lowest I.P.)
  6. Hardness and melting points/boiling points
    These metals are very soft and can be cut with a knife. Lithium is harder than any other alkali metal.
    The hardness depends upon cohesive energy
    M. P.   Li>Na>K>Rb>Cs
    B. P.    Li>Na >K>Cs>Rb
  7. Electropositive character or metallic character
    Alkali metals are strongly electropositive and metallic. Down the group         
    Electropositive nature increase so metallic nature also increases.
    i.e., M → M+ e-
    Metallic Nature: Electropositive character ∝ begin mathsize 12px style fraction numerator 1 over denominator straight I. straight P end fraction end style
    Li < Na < K < Rb < Cs
  8. Oxidation  state
    Show +1 Oxidation state because by losing one electron they get stable noble gas configuration.
  9. Photoelectric effect
    The phenomenon of emission of electron when electromagnetic rays strikes against them is called photoelectric effect; Alkali metal have low I.P. so show photoelectric effect.
    *Cs and K are used in photoelectric cells.

 Group II Elements (Alkaline earth metals)

  1. The Elements are Be, Mg, Ca, Sr, Ba, Ra,
  2. Outermost Electronic configuration:- -ns2
  3. Atomic and ionic sizes
    *     The atomic and ionic radii of the alkali earth metal are smaller than corresponding alkali metals
    Reason: higher nuclear charge (Zeff)
    *     On moving down the group size increase, as value of n increases.
    Be>Mg>Ca>Sr>Ba
  4. Ionization Enthalpy
    Be>Mg>Ca>Sr>Ba
    Down the group IE decreases due to increase in size
    Q.   IEof AM <IE1 of AEM
    IE2 of AM > IE2 of AEM
    [Where Am= Alkali metal, AEM=Alkaline earth metal]
    Reason: IE1 to AEM is large due to increased nuclear charge in AEM as compared to AM but IE2 of AM is large because second electron in AM is to be removed from cation which has already acquired noble gas configuration.
  5. Melting and Boiling points
    They have low m.p and b.p but are higher than corresponding value of group I.
    Reason: They have two valency electron which may participate in metallic bonding compared with only one electron in AM. Consequently group II elements are harder and have higher cohesive energy and ∴ have much higher m.p/b.p than A.M
  6. Electropositive and Metallic character
    Due to low IE they are strong electropositive but not as strong as AM because of comparatively high IE.
    The electropositive character increase down the group.
    Be<Mg<Ca<Sr<Ba
  7. Oxidation state
    Show + 2 oxidation state.

 

Group –I & II

OXIDES

Sodium Oxide(Na2O):
Preparation:

  1. It is obtained by burning sodium at 180℃ in a limited supply of air or oxygen and distilling off the excess of sodium in vacuum.

    begin mathsize 12px style 2 Na plus 1 half straight O subscript 2 rightwards arrow with 180 to the power of degree on top Na subscript 2 straight O end style
  2. By heating sodium peroxide, nitrate or nitrite with sodium.
    Na2O2+2Na→2Na2O
    2NaNo2+10Na→6Na2O
    2NaNO2+6Na→4Na2O+N2

Properties:

  1. It is white amorphous mass.
  2. It decomposes at 400℃ into sodium peroxide and sodium

    begin mathsize 12px style 2 Na subscript 2 straight O rightwards arrow with 400 to the power of degree straight C on top Na subscript 2 straight O subscript 2 plus 2 Na end style
  3. It dissolve violently in water, yielding caustic soda.
    Na2O+H2O→2NaOH

Sodium Peroxides (Na2O2):

It is formed by heating the metal in excess of air or oxygen at 3000, which is free from moisture and CO2.

2Na+O2→Na2O2

Properties:

  1. It is pale yellow solid, becoming white in air from the formation of a film of NaOH and Na2CO3.
  2. In cold water (∼ 0℃) produces H2O but at room temperature produces
    O2. In ice-cold mineral acids also produces H2O2.

    begin mathsize 12px style Na subscript 2 straight O subscript 2 plus 2 straight H subscript 2 straight O rightwards arrow with tilde 0 to the power of degree straight C on top 2 text  NaOH+H end text subscript 2 straight O subscript 2 end style
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