Points to Remember: Periodic Table Notes | EduRev

Inorganic Chemistry

Chemistry : Points to Remember: Periodic Table Notes | EduRev

The document Points to Remember: Periodic Table Notes | EduRev is a part of the Chemistry Course Inorganic Chemistry.
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Points to be Remembered

1. Name of the Pioneers: The name of some eminent scientists who are pioneer in this area are Dobereiner, (Law of triad) Newlands (Law of octaves). But, the most popular name is Mendeleev due to the discovery of periodic table. Lothar Meyer also did the same work. But, all credits goes to Mendeleev.
2. Periodic Law: According to Mendeleev the physical and chemical properties of elements are periodic functions of their atomic number.
3. Brief Description of Periodic Table: Mendeleev periodic table contains seven periods and eight groups.

  • First Period: It contains only two elements. That's why it is called very short period.
  • B. Second and Third Periods: Both 2nd and 3rd periods contains eight elements. That's why these are called first and second short periods.
  • C. Fourth and Fifth Periods: Both 4th and 5th periods contains eighteen elements. That's why these are called first and second large periods.
  • D. Sixth and Seventh Periods: Sixth period contains thirty two elements. That's why it is called very large period. As the seventh period is incomplete, it is called incomplete period.

4. Position of Hydrogen in Periodic Table: It shows resembles with both alkali metals as well as halogens.

  • Resembles with Alkali Metals:
    I. The valence shell electronic configuration of H is Is1 similar to the that of the valence shell electronic configuration of alkali metals ns1.
    II. The oxidation state is H atom is +1 similar to that exhibited by alkali metals.
    III. Hydrogen can form stable hydroxide water. Similarly, alkali metals can also form stable hydroxides MOH.
  • Resembles with Halogens:
    I. The valence shell electronic configuration of H atom is 1s1. Hence, it is one electron short from stable duplet configuration. The valence shell electronic configuration of halogens is ns2np5. Hence, it is also one electron short from the stable octet configuration.
    II. Hydrogen can convert to H-. Similarly, X can also get convert to X-.
    III. Hydrogen exist as diatomic molecule H2. Similarly, halogen also exist as the diatomic molecule X2. Due to the resembles with both alkali and halogens, hydrogen is called 'Naughty Element'.

2. Characteristics of s-block Elements: Some important characteristics of s-block elements are:

  • The valence shell electronic configuration of s- block elements is : ns1-2.
  • These are metallic in nature with low melting point.
  • Their oxides and hydroxides are alkaline in nature.
  • Generally these form colourless compounds (except if not the anions are not coloured).
  • Their ionization potential is comparatively low due to their larger size.
  • These elements does not show variable oxidation states. The oxidation states of alkali (gr. I) and alkaline earth metal (gr. II) are +1 and +2 respectively.
  • The lower membered of alkali (Na, K etc.) and alkaline earth metal (Ca, Sr, Ba etc.) shows colour on flame test.

3. Characteristics of p-block Elements: Some important characteristics of p-block elements are:

  • The valence shell electronic configuration of s- block elements is : np1-6.
  • These are both metallic and non-metallic in nature with comparatively high melting point.
  • Their oxides and hydroxides are alkaline to neutral to acidic in nature.
  • Generally these form both coloured and colourless compounds.
  • Their ionization potential is comparatively high due to their smaller size and higher effective nuclear charge.
  • These elements shows variable oxidation states.
  • These does not show colour on flame test.

4. Characteristics of d-block Elements: Some important characteristics of s-block elements are:

  • The valence shell electronic configuration of s- block elements is: (n-1)d1-10ns1-2
  • These are metallic in nature with moderate to high melting pointThese form complex compounds.
  • Generally these form coloured compounds.
  • Their ionization potential is comparatively high due to their smaller size and higher effective nuclear charge.
  • These elements shows variable oxidation states.

5. Characteristics of f-block Elements: Some important characteristics of s-block elements are:

  • The valence shell electronic configuration of s- block elements is : (n-2)f1-14(n-1)d0-1ns1-2.
  • These has high melting point.
  • These form complex compounds.
  • Generally these form coloured compounds.
  • Their ionization potential is comparatively high due to their smaller size and higher effective nuclear charge.
  • These elements shows a few oxidation states.
  • Most of the elements are radioactive in nature.

6. Radius Order: The order of various radius is : vander Waals radius > covalent radius > metallic radius.

7. Ionization Energy:
Definition: The minimum amount of energy required to knock out the most loosely bound electron from the valence shell of an isolated gaseous neutral atom. This is called first ionization energy.
Similarly, the minimum amount of energy required to knock out the most loosely bound electron from the valence shell of an isolated gaseous mono-positive ion. This is called second ionization energy.

B. Factors Affecting Ionization Energy: The following factors affects the ionization energy of an element.

  • Nuclear Charge: Greater is the effective nuclear charge of an atom, higher will be the force of attraction between the valence shell electron and the nucleous. Hence, more energy will require to knock out the electron by breaking this attraction. As the nuclear charge increase from left to right along a period, I.E. will also increase from left to right along a period. That's why nitrogen has greater I.E. than carbon. Similarly, fluorine has greater I.E. than oxygen.
  • Size of an Atom: Smaller is the size of an atom, more will be the force of attraction between the valence shell electron and the nucleous. Hence, more energy will require to knock out the electron by breaking this attraction. As the size increase down a group, the I.E. will decrease down the group. That's why sodium has lower I.E. than lithium.
  • Penetration Power of Orbitals: The order of the penetration power of orbitals are : s > p > d > f. Hence, an electron present in a s-orbital is more stable than an electron present in p-orbital. Hence, more energy will require to knock out the electron from a s-orbital than that from a p-orbital.
  • Special Stability due to Half-filled and Fullfilled Configurations: According to the concept of exchange energy, half-filled and full-filled orbitals are more stable than their neighbours. Hence, to knock out an electron from a half-filled or, full-filled orbital will require more energy than normal expectation. Nitrogen and oxygen has valence shell electronic configurations : 2s22p3 and 2s22p4. From the concept of nuclear charge oxygen should have higher I.E. than nitrogen. But, in case of nitrogen the electron has to be removed from specially stabilized half-filled p-orbital. Hence, removal of electron will destroy the special stability making the electron removal process highly unfavourable. Hence, the greater amount of energy will require to knock out the electron from nitrogen atom. That's why nitrogen has greater I.E.1 than oxygen.

C. Successive Ionization Energy: In case of I.E.1 the electron has to be removed from neutral atom. For I.E.2,I.E.3,.... the electrons have to be removed from uni-positive, di-positive,... ions. As the order of size of the species of same element are: Neutral > Unipositive > Di-positive > .... and the order of nuclear charge of the species of same element are : Neutral < Uni-positive < Di-positive < ..... . the attraction between the nucleous and the valence shell electron will increase successively. Hence, the amount of energy required to knock out the most loosely bound electron will increase for the successive species. Hence, the successive ionization energies will increase i.e. I.E.1 < I.E.2 < I.E.3 < .....

8. Electron Affinity:
A. Definition: The amount of energy will be released when an electron is added to the valence shell of an isolated gaseous neutral atom. This is called first ionization energy.
B. Factors Affecting Ionization Energy: The following factors affects the ionization energy of an element.

  • Nuclear Charge: Greater is the effective nuclear charge of an atom, higher will be the force of attraction between the coming electron and the nucleous. Hence, more energy will be released due to the addition of electron. As the nuclear charge increase from left to right along a period, E.A. will also increase from left to right along a period. That’s why oxygen has greater E.A. than carbon.
  • Size of an Atom: Smaller is the size of an atom, more will be the force of attraction between the coming electron and the nucleous. Hence, more energy will be released due to the addition of electron. As the size increase down a group, the E.A. will decrease down the group. That's why iodine has lower E.A. than fluorine.
  • Penetration Power of Orbitals: The order of the penetration power of orbitals are : s > p > d > f. Hence, when an electron is added to the s-orbital become more favourable than when the electron is added to a p-orbital.
  • Special Stability due to Half-filled and Fullfilled Configurations: According to the concept of exchange energy, half-filled and full-filled orbitals are more stable than their neighbours. Hence, if the electron addition occurs to the half- filled or, full-filled orbitals, the process will become unfavourable. Hence, E.A. will decrease. Nitrogen and oxygen has valence shell electronic configurations : 2s22p3 and 2s22p4 Hence, the addition of an extra electron to nitrogen will destroy the special stability due to half-filled configuration. Hence, the electron addition process will become highly unfavourable. That's why E.A. value of nitrogen become positive.

C. Successive Electron Affinity:

  • First E.A.: When an electron is going to add to the valence shell of an isolated gaseous atom, two types of interactions will occur:
    1. The attractive force between the positively charged nucleous and the negatively charged coming electron.
    2. The repulsive force between the negatively charged coming electron and the negatively charged electrons that are present in the atom. If the system is already not stabilised due to half and full-filled configurations, the attractive interaction will be dominates over the repulsive interaction making the electron addition process favourable. Hence, the E.A.i value will become negative.
  • Successive E..A.: For second, third,.... E.A.s the addition of electrons will occur to the uni-negative, di-negative,.... ions. Hence, the electron-electron repulsion factor will dominates over the electron- proton attraction factor. Hence, addition of successive electrons will become unfavourable making the successive positive. E.A.s values.

9. Electronegativity:
A. Definition: The property due to which an atom withdraw the bonding electron pair more towards itself is known as electronegativity.
B. Factors Affecting Electronegativity: The following factors affects the electronegativity of an element.

  • Nuclear Charge: Greater is the nuclear charge of an atom present in a bond, more it will have attraction for the negatively charged bonded electrons. Hence, its electron attraction power i.e. electronegativity will increase. As the nuclear charge increase from left to right in a period, the electronegativity will also increase along a period. That's why fluorine has greater electronegativity than oxygen.
  • Size of an Atom: Smaller is the size of an atom greater it will have attraction for the negatively charged bonded electrons. Hence, its electronegativity will increase. As the size increase down the group, electronegativity will decrease down the group. Hence, iodine has electronegativity than fluorine.
  • Oxidation State of an Atom: Higher is the oxidation states of an atom, smaller will be size and higher will be its nuclear charge. Hence, greater will be its electron withdrawing power i.e. greater will be its electronegativity. That's why Fe3+ has greater electronegativity than Fe2+.
  • Hybridization of the Central Atom: As s-orbital is closer to the nucleous, greater is the % of s character in a hybrid orbital, higher will be its electronegativity. The % of s character of the sp, sp2 and sp3 hybrid orbitals are 50 %, 33.3% and 25% respectively. Hence, the electronegativity of the hybrid orbitals will be: sp > sp2 > sp3.

10. Inert Pair Effect: The penetration power of various orbitals runs as : s > p > d > f. Due to lowest shielding power of f-orbital, when electrons enter to f- orbital, the effective nuclear charge at the periphery of the atom increase. This will cause the contraction of 6s orbital. This is called f-orbital contraction.
Due to the f-orbital contraction the most stable O.S. of these element like Pb is decreased by two unit than the group O.S. This is called inert-pair effect. Due to the inert pair effect the most stable oxidation state for these elements will become two unit less than their group oxidation state. Evidence : Pb is a gr-14 element. Hence, its group oxidation state is +4. But, due to the inert pair effect the most stable oxidation will be +2. That's why PbClis almost non-existent.
11. Diagonal Relationship: The physical and chemical properties of some diagonal elements such as Li and Mg, Be and Al, B and Si are similar. This is known as diagonal relationship. The physical and chemical properties is largely dependent on the ionic potential of the element. Ionic potential is defined as the : ϕ = charge of cation/size of cation. As the second element of each pair has larger size and larger charge, the ionic potential of both the elements in each pair will be same. That's why the physical and chemical properties will become similar.

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