Classification Of Elements And Periodicity In Properties
Need Of ClasssificationWith passage of time,many more new elements were discovered.Study of these elements and compounds individually
became more difficult.So, it was felt that these elements should be classified into a few groups to make their study systematic and easier.
Dobereiner's Triads When the elements in a triad were arranged in the order of increasing atomic masses , the atomic mass of the middle element was found to be approximately equal to the arithmatic mean of the atomic mass of other two elements e.g. Li (7),Na (23), K(39) .
Limitation This idea of classification could not be applied to all the elements known at that time.
Newland's Law Of Octaves The properties of the elements were repeated at every eighth element like the eighth note of an octave in music.
LimitationIt couldn't work for elements having atomic number more than 20.
Mendeleef's Periodic Law "The Properties of the elements are a periodic function of their atomic masses".
Essential Features Of Mendeleef's Periodic Table
1. There were eight groups and seven periods.
2. The elements in each group resembled each other in many properties.
3. He left some gaps to accomodate yet-to-be discovered elements.
Limitations1. Certain pairs of elements were placed in reverse order of atomic masses e.g. Co was placed before Ni.
2. Isotopes have similar chemical properties and different atomic mass but, isotopes were not given seperate places.
3. Chemically dissimilar elements have been grouped together e.g. Cu and Ag have been grouped together in group I.
4. Position of hydrogen was not made clear.
Modern Periodic Table "The properties of elements are periofic function of their atomic number".
Features Of Modern Periodic Table
It has 7 horizontal rows called Periods and 18 vertical columns called Group.
* Group 1 elements are called Alkali metals.
* Group 2 elements are called Alkaline earth metals.
* Group 15 elements are called Pnicogens.
* Group 16 elements are called Chalcogens.
* Group 17 elements are called Halogens.
* Group 2 elements are called Noble gases.
3.* 1st period - 2 elements
* 2nd and 3rd period - 8 elements
* 4th and 5th period - 18 elements
* 6th period - 32 elements
* 7th period - incomplete (32 elements)
4.* s block elements - groups 1 and 2.
* p bock elements - groups 13 to 18 .
* d block elements - groups 3 to 12 .
* Two f-block series lanthanoids and actinoids are placed in the bottom of periodic table.
Advantages Of Modern Periodic Table
1. It made the study of chemistry systematic and easy.
2. It is based upon atomic number which is more fundamental property.
3. It removes all the drawbacks of Mendeleef's Table.
IUPAC Nomenclature Of Elements With Atomic No. > 100 The abbreviations are used for digits are
0 - nil
1 - un
2 - bi
3 - tri
4 - quad
5 - pent
6 - hex
7 - sept
8 - oct
9 - enn
s-block elementsElements in which the last electron enters the s-orbital of outermost shell.
General Electroic Configuration : ns1-2
1. The elements in which last electron enters in s-orbital ,are known as s-block elements.
2. Their general electronic configuration is ns1-2.
3. They are highly electropositive in nature.
4. They impart characteristic colour to the flame.
5. They behave as good reducing agent.
6. They are soft metals.
7. They have low melting and boiling points.
8. They form ionic compounds.
p-block elementsElements in which the last electron enters p-orbitals of outermost shell.
General Electroic Configuration : ns2np1-6
1. The elements in which last electron enters in p-orbital,are known as p-block elements.
2. Their general electronic configuration is ns1-2np1-6.
3. They include both metals and non-metals but the no. of non-metals is much higher than metals.
4. They have ionisation enthapy than s-block elements.
5. They mostly form covalent compounds.
6. They can show variable oxidation states.
d-block elements Elements in which the last electron enters d-orbitals of penultimate shell.
General Electroic Configuration :(n-1)d1-10ns0-2
1. The elements in which last electron enters in d-orbital, are known as d-block elements.
2. Their general electronic configuration is ns0-2(n-1)d1-10.
3. Ther are hard and high density metals.
4. They show variable oxidation states.
5. They form coloured complexes.
6. They are metallic in nature.
7. Ther have high melting and boiling points.
8. They are used as catalysts.
f-block elementsElements in which the last electron enters f-orbitals of ante-penutimate shell.
General Electroic Configuration :(n-2)f0-14(n-1)d0-2ns2
1. The elements in which last electron enters in f-orbital,are known as f-block elements.
2. Their general electronic configuration is (n-2)f1-14(n-1)d0-1ns2.
3. No. of radioactive elements is more than other blocks.
4. They show variable oxidation states.
5. They are paramagnetic in nature.
6. They show shielding effect.
7. They all are metals.
General Properties And Their Trends:
Atomic Radius The distance from the centre of the nucleus to the outermost shell containing electrons.
(a) Atomic Radius : It is one-half the distance between the nuclie of two covalently bonded atoms of the same element in a molecule.
(b)Vander Waals Radius : It is one- half the distance between the nuclie of two identical non-bonded isolated atom.
(c)Metallic Radius : It is one-half the intermolecular distance between the two adjacent metal ions in the metallic lattice.
Order : Vander Walls Radius > Metallic Radius > Covalent Radius
Variation Along A Period : The Atomic radius decreases with increase in atomic number.
Explanation : As we move from left to right in a period, nuclear charge increases by one unit in each succeeding element while the no. of shells remains the same.So, the electrons of all the shells are pulled closer to the nucleus. This results in a decrease of atomic radius.
Important Trends :
(a) N > O > F
Explanation :When we move from N to O , the nuclear charge increases by one. But at the same time one of p-orbital has now two electrons which repel each other. However, in case of O inter - electronic repulsions outweigh the effect of increased nuclear charge and hence the atomic size increases from N to O.
On further moving from O to F, the nuclear charge further increases by one and at the same time two of the p-orbitals now have two electrons each which repel each other. However, in case of F , the enhanced nuclear charge outweighs the effect of inter-electronic repulsions and hence the size decreases from O to F.
The size of the atoms of inert gases are larger than those of the preceding halogens.
Explanation :This is due to the reason that in case of inert gases all the orbitals are completely filled and hence the inter-electronic repulsions are maximum. Moreover in case of inert gases,the atomic size is expressed in terms of vander waal's radius since they do not form covalent bonds while in case of all other elements,the atomic size is expressed in terms of covalent radius.
Variation With In A Group :The atomic radii of elements increase with increase in atomic number as we move from top to bottom in a group.
ExplanationAs we move down a group, new energy shelll is added to each succeeding element and the valence electrons lie farther and the farther away from the nucleus.As a result, the attraction of the nucleus for the electrons decreases and hence the atomic radius increases.
Important Trends
(a) The radius of the catio is always smaller than that of its parent atom.
Explanation: It is due to (i)Decrease in the no. of shells (ii) Increase in the effective nuclear charge resulting in greater force of attraction by the nucleus on the electrons.
(b) The radius of the anion is always larger than that of its parent atom
ExplanationIt is due to decrease in effective nuclear charge i.e. lesser force of attraction by the nucleus on the electrons.
Isoelectronic IonsIons of different elements which have the same no. of electrons but different magnitude of nuclear charge.
Ionic Radii Order Of Isoelectronic Species Al3+ < Mg 2+ < Na+ < F- < O2- < N3-
1. Ionisation Energy : It is the amount of energy required to remove an electron from an isolated gaseous atom of an element to produce cation i.e. M - e- -----> M+
Successive ionisation energy of an atom is greater than previous one.
Explanation :
When an electron has been removed from the neutral gaseous atom, the positively charged ion formed has one electron less than the no. of protons in the nucleus. As a result, the electrostatic attraction between the nucleus and the remaining electrons in the cation increases. Therefore,the energy required to remove another electron from this positively charged ion would be higher than the first and so on.
Factors Affecting Ionisation Energy
(a). Nuclear Charge : Ionisation energy is directly proportional to nuclear charge and it is due to the fact that with increase in nuclear charge ,the electrons of the outer shell are more firmly held by nucleus and thus greater energy is required to pull out that electron.
(b). Atomic Radius : Ionisation energy is inversely proportional to atomic radius/size becouse as we increase the distance of outer electrons from the nucleus ,attractive force on the outer electrons decreases.As a result ,outer electrons are held less firmly and lesser energy is required.
(c). Penetration Effect : Ionisation energy is directly proportional to penetration effect. Penetration effect decreases in the order : s>p>d>f.
(d). Shielding Effect : Ionisation energy is inversely proportional to the shielding effect.The actual charge felt by the valence shell electrons will be decreased due to shielding effect; Zeff = Z - S
So, if (S) will be high,means attraction will be less and ionisation energy will also be less.
(e). Electronic Configuration : Half filled and fully filled orbitals are more stable than other orbitals.
Trends in periodic table
(a). In Period, Ionisation energy will increase with increase in atomic number.
Explanation:As we move across a period,the nuclear charge increases and the atomic radius decreases. As a result of increased nuclear charge and decrease in atomic radius,the valence electrons are more and more tightly held by nucleus.Consequently,more and more energy is needed to remove the electron and hence ionisation enthalpies keep on increasing.
(b). In Group :- Ionisation energy will decrease with in a group.
Explanation:On moving down the group,atomic size increases due to addition of one new shell at each succeeding element.As a result,the force of attractiom by the nucleus decreases and hence ionisation energy should decrease.With the addition of new shells,shielding effect increases.As a result, force of attraction decreases and hence ionisation energy should decrease.Nuclear charge increases with increase in atomic no.As a result,the force of attraction by the nucleus for the valence electrons should increase.The combined effect of the increase in atomic size and screening effect more than the effect of increased nuclear charge hence the ionisation energy gradually decrease.
2. Electron Gain Enthalpy : It is the amount of the energy released when an electron is added to an isolated gaseous atom.
M + e- -----> M-
It is the property of Non-metals. It is of two types : +ve and -ve.
(A) The elements like halogens have -ve electron gain enthlpy becouse they have strong tendency to accept electrons to get noble gas configuration.
(B) The elements having stable configuartion have positive value of electron gain enthalpy becouse they don't want to add electron anymore.
Factors Affecting Ionisation Energy
(a). Atomic Size : With increase in atomic size ,it will become less negative becouse force of attraction will be less.
(b). Nuclear Charge : As the nuclear charge increases, attraction will be increaed so, its value become more -ve.
(c). Electronic Configuration : For symmetrical configuration like half filled and fully filled configurations have not any tendency to gain electron so, it will be +ve.
Trends in periodic table :
In Period :- Electron gain enthalpy generally becomes more -ve in period.
Explanation:As we move across a period,the atomic size decreases and the nuclear charge increases. Both these factors tend to increase attraction by the nucleus for the incoming electron and hence electron gain enthalpy becomes more and more negative.
Halogens have the most negative electron gain enthapies.
Explanation:The valence shell electronic configuration for halogens is ns2ns5 and they require one more electron to aquire stable noble gas configuration.As a result,they have a strong tendency to accept an additional electron and hence their electron gain enthalpy are highly negative.
The electron gain enthalpy of noble gases is positive.
Explanation: Atoms of these elements have completely filled subshells. As a result,there is no room in their valence orbitals and the additional electron has to be placed in an orbital of next higher shell. As a result,energy has to be supplied to add an additional electron.
In Group :- Electron gain enthalpy generally becomes less -ve in group.
Explanation:With increase in atomic size,the attraction of nucleus for the incoming electron decreases and hence electron gain enthalpy becomes less negative.
The Electron gain enthalpy of some of elements (O,F) are less negative than S and Cl.
Explanation:Elements of second periods have smallest atomic size in their respective groups. As a result,there are electron-electron repulsions with in the atom itself and hence the additional electron is not accepted with the same ease.
3. Electronegativity : The ability of an atom atom to attract the bonding electron pair towards itself. It depends upon ionisation energy and electron affinity.
x = I.E. + E.A./5.6
Factors Affecting Electronegativity :-
(a). Zeff. : Electronegativity is directly proportional to Zeff. value.
(b). Atomic size : Electronegativity is inversely proportional to atomic size.
(c). Oxidation State : i. As positive oxidation state increases,electronegativity increases.
ii. As negative oxidation state increases,electronegativity decreases.
(d). % S-character : Electronegativity is directly proportional to % S-character.
Trends in periodic table (a). In Period :- In period,Zeff. increases therefore electronegavity increases.
(b). In Group :- In Group, atomic size increases therefore electronegativity decreases.
Applications: 1. In prediction of nature of bond.
2. In calculation of % ionic character.
3. In explaining bond angles.
4. In prediction of acidic/basic nature of compounds.
Dobereiner's Triads When the elements in a triad were arranged in the order of increasing atomic masses , the atomic mass of the middle element was found to be approximately equal to the arithmatic mean of the atomic mass of other two elements e.g. Li (7),Na (23), K(39) .
Limitation This idea of classification could not be applied to all the elements known at that time.
Newland's Law Of Octaves The properties of the elements were repeated at every eighth element like the eighth note of an octave in music.
LimitationIt couldn't work for elements having atomic number more than 20.
Mendeleef's Periodic Law "The Properties of the elements are a periodic function of their atomic masses".
Essential Features Of Mendeleef's Periodic Table
1. There were eight groups and seven periods.
2. The elements in each group resembled each other in many properties.
3. He left some gaps to accomodate yet-to-be discovered elements.
Limitations1. Certain pairs of elements were placed in reverse order of atomic masses e.g. Co was placed before Ni.
2. Isotopes have similar chemical properties and different atomic mass but, isotopes were not given seperate places.
3. Chemically dissimilar elements have been grouped together e.g. Cu and Ag have been grouped together in group I.
4. Position of hydrogen was not made clear.
Modern Periodic Table "The properties of elements are periofic function of their atomic number".
Features Of Modern Periodic Table
It has 7 horizontal rows called Periods and 18 vertical columns called Group.
* Group 1 elements are called Alkali metals.
* Group 2 elements are called Alkaline earth metals.
* Group 15 elements are called Pnicogens.
* Group 16 elements are called Chalcogens.
* Group 17 elements are called Halogens.
* Group 2 elements are called Noble gases.
3.* 1st period - 2 elements
* 2nd and 3rd period - 8 elements
* 4th and 5th period - 18 elements
* 6th period - 32 elements
* 7th period - incomplete (32 elements)
4.* s block elements - groups 1 and 2.
* p bock elements - groups 13 to 18 .
* d block elements - groups 3 to 12 .
* Two f-block series lanthanoids and actinoids are placed in the bottom of periodic table.
Advantages Of Modern Periodic Table
1. It made the study of chemistry systematic and easy.
2. It is based upon atomic number which is more fundamental property.
3. It removes all the drawbacks of Mendeleef's Table.
IUPAC Nomenclature Of Elements With Atomic No. > 100 The abbreviations are used for digits are
0 - nil
1 - un
2 - bi
3 - tri
4 - quad
5 - pent
6 - hex
7 - sept
8 - oct
9 - enn
s-block elementsElements in which the last electron enters the s-orbital of outermost shell.
General Electroic Configuration : ns1-2
1. The elements in which last electron enters in s-orbital ,are known as s-block elements.
2. Their general electronic configuration is ns1-2.
3. They are highly electropositive in nature.
4. They impart characteristic colour to the flame.
5. They behave as good reducing agent.
6. They are soft metals.
7. They have low melting and boiling points.
8. They form ionic compounds.
p-block elementsElements in which the last electron enters p-orbitals of outermost shell.
General Electroic Configuration : ns2np1-6
1. The elements in which last electron enters in p-orbital,are known as p-block elements.
2. Their general electronic configuration is ns1-2np1-6.
3. They include both metals and non-metals but the no. of non-metals is much higher than metals.
4. They have ionisation enthapy than s-block elements.
5. They mostly form covalent compounds.
6. They can show variable oxidation states.
d-block elements Elements in which the last electron enters d-orbitals of penultimate shell.
General Electroic Configuration :(n-1)d1-10ns0-2
1. The elements in which last electron enters in d-orbital, are known as d-block elements.
2. Their general electronic configuration is ns0-2(n-1)d1-10.
3. Ther are hard and high density metals.
4. They show variable oxidation states.
5. They form coloured complexes.
6. They are metallic in nature.
7. Ther have high melting and boiling points.
8. They are used as catalysts.
f-block elementsElements in which the last electron enters f-orbitals of ante-penutimate shell.
General Electroic Configuration :(n-2)f0-14(n-1)d0-2ns2
1. The elements in which last electron enters in f-orbital,are known as f-block elements.
2. Their general electronic configuration is (n-2)f1-14(n-1)d0-1ns2.
3. No. of radioactive elements is more than other blocks.
4. They show variable oxidation states.
5. They are paramagnetic in nature.
6. They show shielding effect.
7. They all are metals.
General Properties And Their Trends:
Atomic Radius The distance from the centre of the nucleus to the outermost shell containing electrons.
(a) Atomic Radius : It is one-half the distance between the nuclie of two covalently bonded atoms of the same element in a molecule.
(b)Vander Waals Radius : It is one- half the distance between the nuclie of two identical non-bonded isolated atom.
(c)Metallic Radius : It is one-half the intermolecular distance between the two adjacent metal ions in the metallic lattice.
Order : Vander Walls Radius > Metallic Radius > Covalent Radius
Variation Along A Period : The Atomic radius decreases with increase in atomic number.
Explanation : As we move from left to right in a period, nuclear charge increases by one unit in each succeeding element while the no. of shells remains the same.So, the electrons of all the shells are pulled closer to the nucleus. This results in a decrease of atomic radius.
Important Trends :
(a) N > O > F
Explanation :When we move from N to O , the nuclear charge increases by one. But at the same time one of p-orbital has now two electrons which repel each other. However, in case of O inter - electronic repulsions outweigh the effect of increased nuclear charge and hence the atomic size increases from N to O.
On further moving from O to F, the nuclear charge further increases by one and at the same time two of the p-orbitals now have two electrons each which repel each other. However, in case of F , the enhanced nuclear charge outweighs the effect of inter-electronic repulsions and hence the size decreases from O to F.
The size of the atoms of inert gases are larger than those of the preceding halogens.
Explanation :This is due to the reason that in case of inert gases all the orbitals are completely filled and hence the inter-electronic repulsions are maximum. Moreover in case of inert gases,the atomic size is expressed in terms of vander waal's radius since they do not form covalent bonds while in case of all other elements,the atomic size is expressed in terms of covalent radius.
Variation With In A Group :The atomic radii of elements increase with increase in atomic number as we move from top to bottom in a group.
ExplanationAs we move down a group, new energy shelll is added to each succeeding element and the valence electrons lie farther and the farther away from the nucleus.As a result, the attraction of the nucleus for the electrons decreases and hence the atomic radius increases.
Important Trends
(a) The radius of the catio is always smaller than that of its parent atom.
Explanation: It is due to (i)Decrease in the no. of shells (ii) Increase in the effective nuclear charge resulting in greater force of attraction by the nucleus on the electrons.
(b) The radius of the anion is always larger than that of its parent atom
ExplanationIt is due to decrease in effective nuclear charge i.e. lesser force of attraction by the nucleus on the electrons.
Isoelectronic IonsIons of different elements which have the same no. of electrons but different magnitude of nuclear charge.
Ionic Radii Order Of Isoelectronic Species Al3+ < Mg 2+ < Na+ < F- < O2- < N3-
1. Ionisation Energy : It is the amount of energy required to remove an electron from an isolated gaseous atom of an element to produce cation i.e. M - e- -----> M+
Successive ionisation energy of an atom is greater than previous one.
Explanation :
When an electron has been removed from the neutral gaseous atom, the positively charged ion formed has one electron less than the no. of protons in the nucleus. As a result, the electrostatic attraction between the nucleus and the remaining electrons in the cation increases. Therefore,the energy required to remove another electron from this positively charged ion would be higher than the first and so on.
Factors Affecting Ionisation Energy
(a). Nuclear Charge : Ionisation energy is directly proportional to nuclear charge and it is due to the fact that with increase in nuclear charge ,the electrons of the outer shell are more firmly held by nucleus and thus greater energy is required to pull out that electron.
(b). Atomic Radius : Ionisation energy is inversely proportional to atomic radius/size becouse as we increase the distance of outer electrons from the nucleus ,attractive force on the outer electrons decreases.As a result ,outer electrons are held less firmly and lesser energy is required.
(c). Penetration Effect : Ionisation energy is directly proportional to penetration effect. Penetration effect decreases in the order : s>p>d>f.
(d). Shielding Effect : Ionisation energy is inversely proportional to the shielding effect.The actual charge felt by the valence shell electrons will be decreased due to shielding effect; Zeff = Z - S
So, if (S) will be high,means attraction will be less and ionisation energy will also be less.
(e). Electronic Configuration : Half filled and fully filled orbitals are more stable than other orbitals.
Trends in periodic table
(a). In Period, Ionisation energy will increase with increase in atomic number.
Explanation:As we move across a period,the nuclear charge increases and the atomic radius decreases. As a result of increased nuclear charge and decrease in atomic radius,the valence electrons are more and more tightly held by nucleus.Consequently,more and more energy is needed to remove the electron and hence ionisation enthalpies keep on increasing.
(b). In Group :- Ionisation energy will decrease with in a group.
Explanation:On moving down the group,atomic size increases due to addition of one new shell at each succeeding element.As a result,the force of attractiom by the nucleus decreases and hence ionisation energy should decrease.With the addition of new shells,shielding effect increases.As a result, force of attraction decreases and hence ionisation energy should decrease.Nuclear charge increases with increase in atomic no.As a result,the force of attraction by the nucleus for the valence electrons should increase.The combined effect of the increase in atomic size and screening effect more than the effect of increased nuclear charge hence the ionisation energy gradually decrease.
2. Electron Gain Enthalpy : It is the amount of the energy released when an electron is added to an isolated gaseous atom.
M + e- -----> M-
It is the property of Non-metals. It is of two types : +ve and -ve.
(A) The elements like halogens have -ve electron gain enthlpy becouse they have strong tendency to accept electrons to get noble gas configuration.
(B) The elements having stable configuartion have positive value of electron gain enthalpy becouse they don't want to add electron anymore.
Factors Affecting Ionisation Energy
(a). Atomic Size : With increase in atomic size ,it will become less negative becouse force of attraction will be less.
(b). Nuclear Charge : As the nuclear charge increases, attraction will be increaed so, its value become more -ve.
(c). Electronic Configuration : For symmetrical configuration like half filled and fully filled configurations have not any tendency to gain electron so, it will be +ve.
Trends in periodic table :
In Period :- Electron gain enthalpy generally becomes more -ve in period.
Explanation:As we move across a period,the atomic size decreases and the nuclear charge increases. Both these factors tend to increase attraction by the nucleus for the incoming electron and hence electron gain enthalpy becomes more and more negative.
Halogens have the most negative electron gain enthapies.
Explanation:The valence shell electronic configuration for halogens is ns2ns5 and they require one more electron to aquire stable noble gas configuration.As a result,they have a strong tendency to accept an additional electron and hence their electron gain enthalpy are highly negative.
The electron gain enthalpy of noble gases is positive.
Explanation: Atoms of these elements have completely filled subshells. As a result,there is no room in their valence orbitals and the additional electron has to be placed in an orbital of next higher shell. As a result,energy has to be supplied to add an additional electron.
In Group :- Electron gain enthalpy generally becomes less -ve in group.
Explanation:With increase in atomic size,the attraction of nucleus for the incoming electron decreases and hence electron gain enthalpy becomes less negative.
The Electron gain enthalpy of some of elements (O,F) are less negative than S and Cl.
Explanation:Elements of second periods have smallest atomic size in their respective groups. As a result,there are electron-electron repulsions with in the atom itself and hence the additional electron is not accepted with the same ease.
3. Electronegativity : The ability of an atom atom to attract the bonding electron pair towards itself. It depends upon ionisation energy and electron affinity.
x = I.E. + E.A./5.6
Factors Affecting Electronegativity :-
(a). Zeff. : Electronegativity is directly proportional to Zeff. value.
(b). Atomic size : Electronegativity is inversely proportional to atomic size.
(c). Oxidation State : i. As positive oxidation state increases,electronegativity increases.
ii. As negative oxidation state increases,electronegativity decreases.
(d). % S-character : Electronegativity is directly proportional to % S-character.
Trends in periodic table (a). In Period :- In period,Zeff. increases therefore electronegavity increases.
(b). In Group :- In Group, atomic size increases therefore electronegativity decreases.
Applications: 1. In prediction of nature of bond.
2. In calculation of % ionic character.
3. In explaining bond angles.
4. In prediction of acidic/basic nature of compounds.