Periodic Table & Periodicity in Properties Notes: Class 11, JEE, NEET & AIIMS

CLASSIFICATION OF ELEMENTS AND PERIODICITY IN PROPERTIES IN PERIODIC TABLE

Periodic Table is the most important concept in chemistry, both in principle and in practice. It is the everyday support for students, it suggests new avenues of research to professionals, and it provides a better organization of the whole of chemistry.

Earlier attempts to classify elements:

Prout’s hypothesis 1815	The atomic weights of all other elements are exact multiples of that of hydrogen and hence hydrogen is the primary substance from which the other elements have been formed. It is also called the unitary method.
Dobereiner’sTriads 1829.	The  atomic mass of the middle element of a Triad is the arithmetic mean of the atomic masses of the other two elements. Ex.. Elements                 Li   Na    K    Mean of atomic mass = 7+39/2 Atomic weight       7     23     39            =23
Newland’s octaves  1864	He called this relationship as the Law of octaves. According to Newlands’ law of octaves when the elements are arranged in order of increasing atomic weights then every eighth element has properties similar to that of the first element. analogy with the intervals of the musical scale. Ex..
LotherMeyer’s 1869	He classifies elements in the form of curve b/w atomic volume and atomic mass. The properties of the elements are the periodic function of their atomic volumes.

Mendeleev’s Periodic Table

• He observed that the properties of elements, both physical and chemical, were periodically related to the atomic mass of the elements.
• The Periodic Law (also referred to as Mendeleev’s Law), states that the chemical properties of elements are a periodic function of their atomic weights

 The advantages of Mendeleev’s Periodic table are:

• The inclusion of these newly discovered elements did not disturb the periodic table. Examples include germanium, gallium, and scandium.
•  It was used to correct the wrong atomic weights in use at that time.
• A variance from atomic weight order was provided by Mendeleev’s table.

The limitations of Mendeleev’s Periodic table are:

1.  Hydrogen position was in the group of alkali metals but hydrogen also exhibited halogen like qualities.
2.  Isotopes were positioned differently since this type of classification of elements was done by considering the atomic weight of the element. Therefore – protium, deuterium, and tritium would occupy varying positions in Mendeleev’s table.
3. An anomalous positioning of a few elements showed that the atomic masses did not increase regularly from one element to the next. the atomic mass of 58.9) before nickel (atomic mass of 58.7).
4. He does not explain the electronic arrangement of elements.

Modern Periodic Table (By Moseley)

Modern periodic law:

•  According to the properties of elements all periodic function of their atomic numbers.
• So, when elements are arranged according to increasing atomic no.’s there is periodicity in the electronic configuration that leads to periodicity in their chemical properties.

Introduction of the modern periodic table: It consists of:

horizontal rows (Periods)
Vertical column (Groups)

There are 7 periods and 18 groups in this long form of the periodic table.
The number of elements in each period:

•  Ist period has 2 elements
•  IInd period has 8 elements
•  IIIrd period has 8 elements
• IVth period has 18 elements
•  Vth period has 18 elements
• VIth period has 32 elements
• VIIth period has rest of elements
•  Number of valence electron in atom of elements decides which elements will be first in period and which will be last.

Families of different elements

• 1 to 2 group and 13 to 17 contain normal or representative elements
• 3 to 12 group –  transition elements.
• 57 to 71   lanthanides
•  89 to 103  Actinides.Left hand side – metals.
  Right hand side – non-metals

Valence electrons of different groups:

• Hydrogen element has been placed at top of Ist group. Electronic configuration of H is similar to alkali metal both have 1 valence electron.

Valence electron of group		Element
1	→	1
2	→	2
13	→	3
14	→	4
15	→	5
16	→	6
17	→	7
18	→	8

IUPAC  Nomenclature of elements with atomic numbers more than 100

Digit	Root	Abbreviation
0	Nil	N
1	Un	U
2	Bi	B
3	Tri	T
4	Quad	Q
5	Pent	P
6	Hex	H
7	Sept	S
8	Oct	O

Recommended and official names of elements with Z more than 100

Z	Recommended	Symbol	IUPAC official  name	IUPAC Symbol
101	Unnilunium	Unu	Mendelevium	Md
102	Unnilbium	Unb	Nobelium	No
103	Unniltium	Unt	Lawrencium	Lr
104	Unnilquadium	Unq	Rutherfordium	Rf
105	Unnilpentium	unp	Dubnium	Db

Element in s,p, d and f block and their electronic configuration in Periodic Table

  S-block Elements:

•  Elements in which the last electron enters the s-orbitals.
• General outer shell electronic configuration of s-block elements – ns^1-2 where n = 2-7

General characteristics of s-block elements in Periodic Table:

They are soft metals with low melting and boiling points,

•  low ionization enthalpies (energies) and are highly electropositive.
•   s -block elements lose the valence (outermost) electrons readily to form +1 (in case of alkali metals) and +2 ions (in case of alkaline earth metals)
•  very reactive metals. The metallic character and the reactivity increase as we move down the group. Because of high reactivity, they are never found pure in nature.
• The compounds of s-block elements with the exception of those of beryllium are predominantly ionic.

    P-Block Elements:

• Elements in which the last electron enters any one of three p-orbitals of their respective outermost shells are called P-Block elements.
•  General outer shell electronic configuration of p-block elements –ns² np^1-6 where n = 2-7
  

General characteristics of p-block elements:

• P-block elements include both metal and non-metals but the number of non-metals is much higher than that of metals. Further, the metallic character increases from top to bottom within a group and non-metallic character increase from left to right along a period in this block.
•  Their ionization enthalpies are relatively higher as compared to those s-block elements.
•  They mostly form covalent compounds.
•  Some of the show more than one (variable) oxidation states in their compounds.
•  Their oxidizing character increase from left to right in a period and reducing character increase from top to bottom in a group.

D-block Elements:

• Elements in which the last electron enters any one of the five d-orbitals of their respective penultimate shells are called d-block elements.
• General outer shell electronic configuration of d-block elements ns^1-2 (n-1) d^1-10
  

General Characteristics of d-block elements :

• They are hard, malleable (i.e. can be converted into sheets) and ductile (i.e. can be drawn into wires) metals with high melting and boiling points,

•  a good conductor of heat and electricity.

• ionization enthalpies are between s- and p- block elements.

•  show variables oxidation states.

• form both ionic and covalent compounds.

F-block elements:

• Elements in which the last electron entre any one of the seven f-orbitals of their respective ante-penultimate shells are called f block elements.
• General outer shells electronic configuration of f-block elements – (n – 2)f^1-14 (n – 1)d^0-2 ns²

General Characteristics of f-block elements:

• They have generally high melting and boiling points,
• show variable oxidation states
•  Their compounds are generally colored,
•  have heavy metals,
•  a high tendency to form complexes.

Atomic Properties in Periodic Table:

Atomic radii:	Covalent radii:	Metallic radii:	Ionic radii:	Vander wall radii:
Distance an atom.from the centre of the nucleus to outermost shell of	The distance between the centre of nuclei of atoms and mean position of shared pair of electrons between the bonded atoms	one half the intern nuclear distances between two neighboring metals in lattice	it is the measure of distance between cation and anion in ionic crystals.	A  half of internuclear distance between the adjacent atoms of substance belonging to two different molecules.
For example: The internuclear distance between the two hydrogen atoms in an H2 molecule is measured to be 74 pm. Therefore, the atomic radius of a hydrogen atom is 742=37 pm 74 2 = 37 pm.	For example : H2 molecule the bond length is 74 pm, therefore, covalent radii is 37 pm	For example: Cu metal, the bond length is 256pm therefore radii are128pm	Radius of Fe2+ is 76 pm, while that of Fe3+ is 65 pm.	Hydrogen= 1.2A Carbon= 1.7 A Nitrogen= 1.55A

TRENDS IN ATOMIC PROPERTIES in Periodic Table:

Atomic radius:	Metallic radius:	IonicRadius:
Atomic size decreases on moving from left to right in a period Due to an increase in effective nuclear charge.Increases on moving from top to bottom in a group. Due to addition of new subshells.	As we move towards right the atomic no. increase so as no. of protons increase, attraction increase. So the order is that it keeps on decreasing Along period the size keeps on decreasing due to an increase in nuclear charge. Nuclear charge is the attraction of positively charged protons towards electron. For example, in 3rdperiod: the order is Li>Be>B>C>N>O>F As we move down in the periodic table metallic radius increases due to the addition of shell along the group. It is because every time a new shell is added so the effect of the addition of shell is more than the pronounced nuclear charge, therefore, the order as we discussed is increasing down the group. For example: In the first group the order is Li<Na<k<Rb<Cs	As you move from left to right across an element period (row) the ionic radius decreases. Even though the size of the atomic nucleus increases with larger atomic numbers moving across a period, the ionic and atomic radius decreases. This is because the effective positive force of the nucleus also increases, As you move from top to bottom down an element group (column) ionic radius increases. This is because a new electron shell is added as you move down the periodic table. This increases the overall size of the atom.

Isoelectronic species:

• They are those that have the same number of electron
• The species with a higher positive charge is smaller and the species with a lesser positive charge is bigger similarly the species with a higher negative charge is bigger than the species with a lesser negative charge.
• For example: O^2-, F^–, Na⁺etc. are isoelectronic

Atomic radii:

Metallic radii:

Ionic radii:

• Cation is always smaller than parent atom as it is formed by removal of the electron due to which the same nuclear charge acts on a lesser number of electrons, therefore, the nuclear charge increases and the size decreases
• Anion is always bigger than parent atom as it is formed by gaining electrons due to which the same nuclear charge acts on more number of electrons and moreover the magnitude of repulsions increase, therefore the nuclear charge decrease and size increase

Covalent radii:

Ionization energy in Periodic Table:

•  It is the amount of energy required to remove electron from valence shell of isolated gaseous atom.
• The word required is used because it means ionization energy is positive that is it means it is always given from outside to remove electron.

Successive ionization energies:

It is the amount of energy required to remove the second electron from the ion

For example:
If we compare ionization energy 1, 2 and 3 we found that:
M (g) +∆_iH 1àM+(g) + e- (g)

M¹(g) +∆_iH² àM²⁺ (g) + e^– (g)

M²⁺(g) +∆_iH³ aM³⁺(g) + e^– (g)

There are certain factors on which ionization energy depends:

• Size of atom:
  If the size is small, nuclear charge will be high so it will attract electron more effectively, therefore, ionization energy will be high.
  This means ionization energy is inversely proportional to the size of an atom.
• Charge on nucleus:

If nuclear charge is more than attraction for electron will be more, therefore, ionization energy will be high or vice versa.

Screening effect:
Due to filling of inner orbital’s the nuclear charge is somehow reduced for outermost electron. As a result, the outermost electron will be loosely bounded and therefore ionization energy decreases.

•  Penetration effect:
  The orbital’s which are closer to the nucleus will experience high nuclear charge so the penetrating effect will be more. The order it follows is:

S>P>d>f

•  Electronic configuration:
  Half-filled and fully filled orbitals are more stable therefore they have high ionization energies.

  Variation along period:

• Along period it increases It is because along period the size decreases, nuclear charge increase, therefore, ionization energy increases but the anomalous behavior is seen let us see that:
  Li<Be>B<C<N<O<F<Ne

In this Be has high ionization energy than B because of fully filled shell N has high ionization energy than O because N is half-filled, more stable therefore ionization energy is high.

Variation along group:

• Ionization energy decreases along the group as the size increases, nuclear charge decreases, as a result, less ionization energy is required to remove the electron.
• For example in 1st group the order is:
• Li>Na>K>Rb>Cs out of these cesium has the largest size and least ionization  Energy.

Electron gain enthalpy in Periodic Table:

•  It is the amount of energy released when the electron is added to the outermost shell of an atom.
• This energy release depends upon the extent of attraction for an incoming electron.
• In the case of inert gases, they have almost no attraction Because of stable configuration.

Therefore it is negative for inert gases this means energy should be supplied in order to make them accept an electron.
This energy can be mathematically shown as:
M + e^–à M¹ (negatively charged ion)
Successive electron gain energies:
Factors affecting electron gain enthalpies:

Atomic radius :

More is the size, less is the nuclear charge, therefore, less attraction for incoming electron therefore electron gain enthalpy is less negative for those atoms or vice versa

 Nuclear charge :

less is the nuclear charge, therefore, less attraction for incoming electron, therefore, electron gain enthalpy is less negative for those atoms or vice versa

Electronic configuration:

atoms with stable electronic configurations have less negative electron gain enthalpies.

Variation along group:

•  Electron gain enthalpy becomes less negative as we move down the group
• As we down the group, the size increases, and nuclear charge decreases due to which the attraction for incoming electron decreases so as electron gain enthalpy decreases.
• But certain anomalies are seen like out of chlorine and fluorine, fluorine has less negative electron enthalpy because of its extremely small size the incoming electron suffers interelectronic repulsions, therefore, electron gain enthalpy is less negative as compared to chlorine

Variation along period

• Along period it increases the reason being the size decreases and nuclear charge increases.
• But still, the exceptional behavior is seen like Electron gain enthalpy of inert gases is positive because of their stable electronic configuration.

Electronegativity in Periodic Table:

• It is the tendency of an atom to attract a shared pair of electrons more towards itself. This property depends upon the size and the nuclear charge.
• If the size is less then more will be the nuclear charge hence more attraction for shared pair of electrons.

In the whole periodic table, fluorine is maximum electronegative due to its smallest size.

Other factors on which electronegativity depends:

• State of hybridization:
  The order is sp>sp²>sp³

Oxidation state of element:

• Higher is the oxidation state more it is electronegative.
  

  Variation along group:

• Along group it decreases as size increases and nuclear charge decreases
  

  Variation along period:

Along period it increases as size decreases and nuclear charge increases, therefore, the attraction for a shared pair of electrons increase

Applications of electronegativity:
It tells us about the metallic and non metallic character of atom more is the electronegativity lesser is the metallic character and more is the non metallic character or vice versa.
Polar or non-polar:
If the difference in the electronegativity between the two bonded atoms is more than the bond is polar that is it has lost but if electronegativity difference is less or zero than the bonded non polar.
Metallic and non-metallic character

Metallic character is a tendency to lose an electron.

Non-metallic character: It is the tendency to gain an electron

Along group

• Metallic character increase and non-metallic character decreases
  Size increases i.e. valence electrons become more away from nucleus due to which nuclear charge decrease and metallic character increases whereas non-metallic character decreases.
  For example
  Li.        Na        K      Rb         Cs          Fr
  (least)                                                   (Most)
  Metallic)                                            (metallic)

   Along periods

Metallic character decreases and non-metallic character increases. As we move the Size decreases along period due to which Nuclear Charge increases, therefore, tendency to gain electron increases and to lose electron decreases.

Na      Mg      Al      Si      P        S          Cl

• Metallic character decreases
• Non-metallic character increases
  

                    Nature of oxides:-

  Along group:

• Basic oxides increase as we come across more metallic character down the group.

Along period:

• Basic nature of oxides decreases and acidic nature of oxides increases because size decrease and tendency to lose electron decreases, therefore, oxides are acidic

Chemical reactivity in Periodic Table:

Along group

• It decreases (in case of non-metals)
• It increases (in case of metals)
• F                                     Li
  Cl      Decrease               Na
  Br                                   K
  I                                      Rb       Increases
  Cs
  Fr
  ALONG PERIOD
  It decreases and then increases.

Diagonal relationship in Periodic Table

• It is the anomalous behavior of second group elements. It is defined as the similarity of some elements of second group with elements of third group present diagonally.”
  Like few diagonal relationships are given:

• Elements of 3rd period called as bridge elements.

Important Questions With Answer of Periodic table:

1 Considering the atomic number and position in the periodic table, arrange the following elements in the increasing order of metallic character :
Si, Be, Mg, Na, P.

ANS:

Metallic character increases down a group and decreases along a period as we move from left to right. Hence the order of
increasing metallic character is       P < Si < Be < Mg < Na.

2 Which of the following species will have the largest and the smallest size?
Mg, Mg²⁺, Al, Al³⁺

ANS:

Atomic radii decrease across a period. Cations are smaller than their parent atoms. Among isoelectronic species, the one with the larger positive nuclear charge will have a smaller radius. Hence the largest species is Mg;  the smallest one is Al³⁺

3 Are the oxidation state and covalency of Al in [AlCl(H₂O₅)²⁺] same?

ANS:

No. The oxidation state of Al is +3 and the covalency is 6.

4 Show by a chemical reaction with water that Na₂O is a basic oxide and Cl₂O₇is an acidic oxide.

ANS:

Na₂O with water forms a strong base whereas Cl₂O₇ forms strong acid.
Na₂O + H₂O → 2NaOH
Cl₂O7 + H₂O → 2HClO₄
Their basic or acidic nature can be qualitatively tested with litmus paper.