12th Class Chemistry Chapter 1 Periodic Classification of Elements And Periodicity Short Question Answer

On the basis of similarity in chemical properties.Mendeleev arranged certain elements in certain groups.He placed Be above Mg Ca Sr and Ba.Its atomic mass was though to be 13.5 a.m.u.which was wrong.He gave correct position to it.Its atomic mass was proposed to be 9.00 a.m.u.

According to Moseley’s law the physical and chemical properties of elements are the periodic function of their atomic number.So moseley said that Ar should be placed before K, Co before Ni and Te before I.Moreover he adjusted the positions of the rares earth noble gases and coinage metals.

i)Position of hydrogen was not decided by him. ii)His table did not give the idea of structure of atom. iii)Positions of Lanthanides and actinides could not be decided by Mendeleev.

They are placed in separate group i.e.VIII-A and are the end elements of d-block series.Their outer most orbitals are d.

d- and f-block elements are in between s and p-block elements.Their properties are in between them and they vary from left to the right.It means that their properties show a transition in a systematic way.That is why they are called transition elements.

The elements of group I-A are alkali and alkaline earth metals.Their outermost sub-shells are s and have one and two electrons respectively in s-orbitals.So they are called s-block elements.

The elements of groups III-A to VII-A have atomic number that their outermost sub-shells are p which are in the process of completion.So they are all called p-block elements.

The first d-block series is 3d in period 4 and ranges from Sc₂₁– Zn₃₀.The second last element Cu₂₉ is a coinage metal in this series.

The elements of group I-A and II-A have low ionization energies as compared to the rest of the periodic table.Their outermost electrons are loosely held and they get free in the crystal lattice.This property makes them good metals.

Period six contains 32 elements as first two as s-block next 14 as d – block and last six as p-block elements.

Hydrogen has one electron is s-orbital like alkali metals.Hydrogen and alkali metals show positive oxidation states of+1 give positive ions in solution state and are deposited at thhe cathode during the electrolysis of their halides.Hydrogen has half filled outermost orbital but I-A group donot.H2 is a gas but I-A group are solids.

Hydrogen has outermost s-orbital unlike halogens.Hydrogen is a gas but Bromine is a liquid while iodine and astatine are solids.H- is unstable but X- are stable.Halogens can show variable valency but hydrogen cannot.

The outermost orbitals of elements of group IV-A are half –filled and make the covalent bond with other elements.They resemble in thermodynamic properties and also agree with each other in reduction properties.

The increase of nuclear charge and no change of shielding effect decrease the atomic radii from left to right.The increasing number of shells and increasing shielding effect increase the atomic radii from top to the bottom.

Shielding effect is due to the electrons in between the nucleus and the outermost electron.Greater the number of electrons in these shells lesser the forces of attraction between the nucleus and the outermost electrons and greater the sizes of atoms.

Actually the shielding effects increase down the groups.Moreover due to increasing number of shells,the forces of attraction between the nucleus and the outermost electrons decrease.So the atomic radii increase inspite of increasing nuclear charges.

When the electron is removed from the outermost orbital of the atom then the remaining electrons are pulled with a greater force so the size decreases.Sometimes the removed electron makes the out most orbit or shell vacant.

When the electron is given to the outermost shell of the atom it becomes negatively charged.Due to greater number of electrons to be accommodated in the outermost shell it has to expand so the size of the negative ion is greater.

Greater the amout of the positive charge on the ion greater the difference of number of protons in the nuclei as compared to the total electrons.So there is greater force of attraction for the electrons by the nucleus and smaller the size.

When the atom is in the isolated state then its electronic cloud is not being disturbed by the neighboring atoms or molecules.In this way we can measure the force of attraction between the nucleus and the outermost electron in the form of ionization energy.

The increasing size increasing number of shells and increasing shielding effects are the important factors which cause the decreasing ionization energy values down the group.

The increasing nuclear chars deceasing sizes and no change in the shielding effects are the important factors which control the increasing ionization energies from left to the right.

The outermost electron to be removed may be in s,p,d,or f orbitals is most penetrating and f-orbital is least.So the electron removal from s-orbital is most difficult and easiest from f-orbital.

When we shift from groups II-A to III-A the nature of orbital changes from s to p orbital.p orbital is less penetrating so electron removal from p –orbital of group III-A is easier.

When we move from V-A to VI-A then we move from half filled p-sub shell of V-A to unstable p-sub shell of VI-A Half filled sub-shell is extra stable.So the electron removal from VI-A is easier.

It increases from left to right in a period and decreases from top to bottom in a group.Along the period sizes decrease and along this group sizes of atoms increases.

In most of the elements the outermost orbital is in the process of completion.In this way they have demand of electrons and feel happier in having the electron.So they accept the electron and get stability.

When the atom is in the isolated state its outermost orbitals are not disturbed by the neighbouring atoms or molecules.So we only measure the force of attraction between the nucleus and the incoming electron.It gives us correct electron affinity value.

When oxygen is given the first electron it releases energy of 141 kj mol^-1 and becomes O^–. When the second electron is given to O^– ,880 kj is to be given.The reason is that the second incoming electron is repelled by O^–.

The second electron is repelled by the negative ion and energy is absorbed in this way.The obsorbed energy is given positive sign.

The elements of group II-A have fulfilled s-orbitals and the incoming electron has to be accommodated in the next higher p-sub shell.So energy has to be provided to given them the electron.For this reason their electron affinities are less than those of group I-A.

For the non-metallic elements there is least tendency to loose the electrons and maximum tendency to have the electrons.These two properties are due to greater nuclear charges small sizes and tightly held electrons.These properties are associated with oxygen sulphur fluorine and chlorine. These elements are highly electronegative.This property increases from left to right and decreases from top to bottom.

The size increase the valence shell becomes away from nuclei and metallic effect becomes more and more dominant.All these things make the electron removal easier.Hence metallic character enhances down the group.Amount of energy evolved when one mole atoms of an element accept electrons is their outermost orbitals in gaseous state.

When we go from left to the right in a period the number of electrons go on increasing in the outermost shell.The tendency to unpair the electrons increase upto group IV-A.In this way the binding forces increase hence melting and boiling points increase upto group IV-A.After that pairing of electrons starts and binging force become less.There are no binding forces in group VIII elements and that is why they are gases at room temperature.

The atomic sizes of the halogens increase with increasing shielding effects.The diatomic molecules of halogens can create inter-molecular forces due to loosely held electrons down the group.This thing increases the melting and boiling points down the group.

The number of electrons in outermost shells go on changing in periods from left to the right so oxidation states go on changing.The number of electrons in the outermost shells remain the same in a group so the oxidation states remain the same.Anyhow the process of unpairing of electrons may happen in a group and oxidation states may change.

d-orbital is not available to oxygen for the promotion of electrons but for sulphur ,Se and Te d-orbitals in the same principle quantum number are available.In this way the greater number of unpaired electrons make the oxidation number to increase.

The outermost orbitals of noble gases are complete.They show least tendency to share or transfer electron during bonding.

Elements of group VIII-A have eight electrons in their outermost principal quantum number except.He which two electrons.It means that their outermost shells are complete and they do not want to react show any oxidation state.

When potential difference is applied to the substances,then free electrons should be available to conduct the electrical current.These free electrons transfer the potential difference from one place to the other.In group I-A the possibility of free electrons increase down the group.

These elements have low ionization values and their outermost electrons are loosely held.For this reason they are good conductors of electricity.

Graphite has hexagonal layered structure.Free electrons are available in between the layers.These electrons are responsible for making graphite as a good conductor.No free electrons are available in diamond.

Smaller the size of the positive or negative ion greater the charge density of the ion.This creates greater force of attraction between the ion and the water molecules.It evolves greater hydration energy.

Hydration energy of an ion depends upon the charge density.Greater the charge density greater the hydration energy. AI³⁺.Mg²⁺.Na⁺¹ have same number of electrons but different number of protons.AI⁺³ has maximum charge  density due to 13 protons 10 electrons,and +3  charge.Na⁺ has least change density because it has 11 protons 10 electrons and +1 charge.

This is due to the decreasing charge densities of halide ions.I^– has maximum size smallest charge density and smallest hydration energy among the halogens.

When elements of group I-A amd II-A combine with halogens  find greater difference of electronegativities.In the way the % age of ionic character is sufficiently higher for such compounds.The ionic character of halides decreases  from NaCI MgCI₂,AICI₃,SiCI₄, to SiCI₂.

Elements on the left side of the periodic Table are electropositive.They have low I.E.value and make ionic bonds  with halogens which have high electron affinities.Element on the right  side  of the Periodic  Table are electronegative.They have least tendencies  to lose electrons and to give to halogens.So they give  covalent halides.NaCI is ionic  to good extant PCI₃,PCI₅,and S₂CI₂ are  covalent halides.They are produced by sharing  of electrons.
CCI₄,SiCI₄,PCI₃,S₂CI₂ are covalent chlorides.

The covalent halides have low melting and boiling points as compared to the ionic halides.The % age of ionic character of halides decrease from left to the right in a period.

Greater the positive oxidation state of elements greater the charge density greater the polarizing power and greater the % age of covalent character in a bond.

Alkali metals are much electropositive than hydrogen.So these metals lose electrons and give to the hydrogen.Transference of electrons takes places and the bond is ionic.Anyhow the % of ionic character goes on increasing down the group for such hydrides.