Periodic Table – Modern, Mendeleev and Periodic Properties

Chemist

1 year ago

Periodic table: The periodic table is an arrangement of the chemical elements into rows (periods) and columns (groups).

Father of the periodic table = Dmitri Mendeleev (a Russian chemist)

Contents

Mendeleev’s Periodic table
Modern Periodic Table
Classification of elements in periodic table
Nuclear charge, Shielding effect and Effective nuclear charge
Periodic trends and periodicity (Atomic properties)

Mendeleev’s Periodic table

Dmitri Mendeleev arranged the known elements in the increasing order of their atomic weight in the form of the table called Mendeleev’s periodic table.

Mendeleev’s periodic law : Mendeleev’s periodic law states that “the physical and chemical properties of elements are periodic functions of their atomic weights.”

It means that if the elements are arranged in order of increasing atomic weight, the elements with similar properties are repeated after a regular interval.

Features of Mendeleev’s periodic table:

The modified form of Mendeleev’s periodic table consists of:

1. Nine vertical columns- Groups:

There are nine vertical columns numbered as I, II, III, IV, V, VI, VII, VIII and zero (noble gases). Except for groups VIII and zero, each group from I and VII is subdivided into two subgroups A and B.

2. Seven horizontal rows- Periods:

There are seven horizontal rows numbered as 1 to 7. The periods are divided into short periods and long periods. The first three periods are called short periods as they contain fewer elements. Fourth, fifth and sixth periods are long periods.

3. Position of lanthanides and actinides:

The lanthanides (14 elements after Lanthanum) and actinides (14 elements after Actinium) are kept separate in two horizontal rows at the bottom of the periodic table.

Note : Only 63 elements were known when Mendeleev’s periodic table was created.

Limitations of Mendeleev’s periodic table:

Some main limitations in Mendeleev’s periodic table are as follows:

Position of hydrogen: Hydrogen forms both the positive ion like alkali metal (group-IA) and negative ion like halogen (group-VIIA), hence it resembles the elements of group IA and group VIIA. Therefore, the position of hydrogen in the periodic table is controversial.
Position of isotopes: Isotopes of the same element have different atomic weight. According to Mendeleev’s periodic law, isotopes of an element should be placed at different places in the periodic table. For example, hydrogen needs three separate positions for protium, deuterium and tritium with atomic weight 1, 2 and 3 respectively. But isotopes were not given a separate place in the periodic table.
Position of lanthanides and actinides: Lanthanides and actinides group consists of 14 elements each. All these elements have been grouped in a single place in the III^rd group respectively in the periodic table despite their different atomic masses. They have not been given proper position within the main frame of the table but are placed outside in two separate rows at the bottom of the periodic table.
Position of anomalous pairs of elements: Certain elements with higher atomic weight were placed before the elements having lower atomic weight. For example: Ar (at.wt.=39.9) was placed before K (at.wt.=39.1) ; Co (at.wt.=59.9) was placed before Ni (at.wt.=58.6), etc. These pairs of elements do not obey periodic law.
Separation of chemically similar elements and grouping of dissimilar elements: In Mendeleev’s periodic table, chemically similar elements like Cu and Hg, Au and Pt, Ag and Tl, Ba and Pb, etc have been placed in separate groups while other dissimilar elements have been placed in same group. For example, coinage metals like Cu, Ag, and Au are lesser reactive metals are placed together with higher reactive metals like Li, Na, K.
Cause of periodicity: It does not explain the cause of periodicity. Because of this, atomic weight is not a good basis for the classification of elements.

Modern Periodic Table

Modern periodic table is based on atomic number.
Modern periodic law was proposed by Moseley.

Modern periodic law : Modern periodic law states that “the physical and chemical properties of elements are periodic functions of their atomic number.”

It means that if the elements are arranged in order of increasing atomic number, the elements with similar properties are repeated after a regular interval.

Long form/ Extended form/ Present form of periodic table:

It was constructed by Bohr based on atomic number and Bohr’s- Bury electronic configuration concept.

Main features of modern periodic table:

Long form or extended form of periodic table (modern periodic table) was constructed by Bohr and is also called Bohr’s periodic table. This is an improved and extended form of Mendeleev’s periodic table. Main features of the modern periodic table are as follows:

1) There are 18 vertical columns called groups.

The elements of group 1, 2 and 13 to 17 are called typical or representative elements.
The elements of group 3 to 12 are known as transition elements.
Elements of group 18 are called noble gases

2) There are 7 horizontal rows called periods. They are denoted by 1, 2, 3, 4, 5, 6, and 7.

3) 14 elements after Lanthanum(La) i.e. elements with atomic number 58 to 71 (Ce to Lu) are called lanthanides whereas, the 14 elements after Actinium(Ac) i.e. elements with atomic number 90 to 103 (Th to Lr) are called actinides. These elements are collectively called f-block elements or inner transition elements. These elements are given two separate rows below the main periodic table.

4) Elements are classified into s-block, p-block, d-block and f-block.

5) Metals and non-metals are separated from each other.

Merits (Advantages) of modern periodic table:

The modern periodic table has overcome the drawbacks of Mendeleev’s periodic table by choosing atomic number as the basis of classification. The main advantages of modern periodic table are as follows:

Atomic number basis: Elements are arranged on the basis of atomic number (number of protons), a fundamental property, rather than atomic mass which can be less consistent for isotopes of elements.
Position of isotopes: Atomic number of isotopes is same, so different isotopes can be placed at same place in periodic table. Thus position of isotopes is completely justified.
Proper solution of Mendeleev’s misfit points: The position of Ar, Co and Te before K, Ni and I respectively in Mendeleev’s table was not according to his periodic law. Long form of periodic table justified this anomaly by choosing atomic number as the basis of classification.
Separate position for subgroups: Separate positions for subgroups A and B is provided in modern periodic table. Separation of subgroups removed the anomalies found in Mendeleev’s periodic table such as grouping of chemically dissimilar elements and separation of chemically similar elements.
Separation of metals and non-metals: Metals are placed on the left and non-metals are placed on the right side of the periodic table.
Division of elements into four blocks: Division of elements into s, p, d, and f-blocks based on their electronic configuration has made their study easier.
Successful to explain periodicity: Modern periodic table has been successful to explain the periodicity in certain atomic properties like atomic radius, ionization potential, etc.

Demerits (Defects) of long form of periodic table:

Position of hydrogen is still controversial.
Position of helium along with the p-block elements is not completely justified as its electronic configuration is 1s².
Lanthanides and actinides are still not placed in main body.
Isotopes have different physical properties but have same place in periodic table.

Classification of elements in periodic table

Bohr’s Classification: Depending on the number of incomplete shell in an atom, the elements in the modern periodic table can be classified into four types.

Inert gas elements:

These elements have completely filled ultimate(valence) shell .
General electronic configuration is ns²np⁶
Because of most stable configuration, they are very less reactive. Hence, known as noble gas or inert gas elements.
These elements are present in ‘0’ group or 18^th group.

2. Representative/main group/normal elements:

These elements have incomplete valence (ultimate) shell.
These elements lie in group IA to VIIA
Elements of group IA and IIA are known as alkali metals and alkaline earth metals respectively.
Elements of group VA, VIA and VIIA are known as pnicogen, chalcogen and halogen family.

3. Transition elements:

These elements have incompletely filled ultimate (n) and penultimate (n-1) shell.
These elements lie in group IB to VIIB and group VIII.
The name transition is due to their properties lies between highly reactive metals on the left side and non-metals on the right side.
General outer electronic configuration is : (n-1)d^1-10ns^1-2

4. Inner transition elements:

These elements have incompletely filled ultimate (n), penultimate (n-1) and antepenultimate (n-2) shell.
These elements lie in group IIB and period 6^th and 7^th.
These are 28 in number.
General outer electronic configuration is : (n-2)f^1-14(n-1)d^0-1ns²
Electronic configuration of Lanthanides is 4f^1-14,5d^0-1,6s²and Actinides is 5f^1-14,6d^0-1,7s²

Classification of elements into blocks:

On the basis of sub-shell (orbital) in which last (differentiating) electron enters, the elements are classified into four blocks: s, p, d and f-blocks.

1. s-block elements:

Elements in which the last (differentiating) electron enters into the s-orbital of the outermost shell are called s-block elements.
S-block consists of elements of group IA (alkali metals) and IIA (alkaline earth metals).
These elements are very reactive metals.
They are very soft and malleable metals.
They have low ionization energy and are highly electropositive.
They have low melting and boiling point.
General outer electronic configuration is ns^1-2
Alkali metals have largest atomic size in corresponding periods.
Their hydroxides are strong bases.
Most of the s-block elements (except Be and Mg) impart characteristics colour to the flame (i.e. flame test).

2. p-block elements:

Elements in which the last (differentiating) electron enters into the p-orbital of the outermost shell are called p-block elements.
General outer electronic configuration is ns²np^1-6
p-block consists of elements of group IIIA (13), IVA(14), VA(15), VIA(16), VIIA(17) and zero(18).

IIIA (13) = Boron family

IVA (14) = Carbon family

VA (15) = Nitrogen family (pnicogen family)

VIA (16) = Oxygen family ( Chalcogen family >> ore forming family)

VIIA (17) = Halogen family (salt forming family)

Zero (18) = inert gases/ noble gases/ rare gases/ aerogens

3. d-block elements:

Elements in which the last (differentiating) electron enters into the d-orbital of the penultimate shell are called d-block elements.
d-block consists of the elements of groups IIIB(3), IVB(4), VB(5), VIB(6), VIIB(7), VIII(8-10), IB(11) and IIB(12).
These are hard metals with high melting and boiling points.
They are called transition elements as they exhibit transition behavior intermediate between the properties of s-and p-block elements.
General outer electronic configuration is : (n-1)d^1-10ns^1-2
They show variable oxidation state.
Most of them form coloured salts and their ions or compounds are paramagnetic in nature. These properties are due to presence of unpaired electrons. For example, CuSO₄.5H₂O is blue in colour.
Most transition elements and their compounds possess catalytic properties. Eg. Fe, Ni, Mo, Pt, etc.

4. f-block elements: Inner transition elements

Elements in which the last (differentiating) electron enters into the f-orbital of the ante-penultimate shell are called f-block elements.
These elements lie in group IIB and period 6^th and 7^th.
These are 28 in number.
General outer electronic configuration is : (n-2)f^1-14(n-1)d^0-1ns²
Electronic configuration of Lanthanides is 4f^1-14,5d^0-1,6s²and Actinides is 5f^1-14,6d^0-1,7s²
They have high melting and boiling point.
They are heavy metals.
They show variable oxidation state, commonly +3 state.
They form coloured salts.
Lanthanides (4f-series) are called rare earth elements since they occur rarely in the earth crust.
Actinides (5f-series) are radioactive elements.

Nuclear charge, Shielding effect and Effective nuclear charge

Nuclear charge:

The total positive charge present in the nucleus is called the nuclear charge.
Its value is always positive and depends on the number of protons present in the nucleus.
It is denoted by the letter ‘Z’.
For example, the value of Z for oxygen is +8.

Shielding effect and Effective nuclear charge:

In multi-electron atoms, the electrons in the valence shell are pulled by the nucleus and repelled by the electrons of inner shells. Thus, outermost electrons experience less attraction from the nucleus under the combined effect of attractive and repulsive force acting on the valence electrons. This effect is called shielding effect or screening effect.
Thus larger the number of electrons in inner shell, the larger will be the screening effect.
The actual charge felt by the valence electron as a result of shielding effect is called an effective nuclear charge (Z_eff).

Effective nuclear charge (Z_eff) = Total nuclear charge (Z) – Screening constant (s)
Value of effective nuclear charge is always less than that of the nuclear charge.
For hydrogen, the effective nuclear charge is equal to the nuclear charge as there is no screening effect.

Periodic trends and periodicity (Atomic properties)

Periodicity of elements:

When the elements are arranged in the modern periodic table in order of increasing atomic number, the occurrence of similar properties of elements after a definite interval is termed as periodicity of an element.
These properties include atomic radius, ionization potential, electron affinity, electronegativity, etc.

Causes of periodicity:

The cause of periodicity in properties is due to the same outermost shell electronic configuration coming at regular intervals.
In the periodic table, elements with similar properties occur at intervals of 2, 8, 8, 18, 18 and 32. These numbers are called magic numbers.

The variation of some properties along with group and period are described below:

Atomic radii (size)

Atomic radius is the distance from centre of the nucleus to the outermost shell of the electrons.
Atomic radius cannot be measured directly because the atom cannot be isolated to determine its radius.

It can be measured indirectly from bond length measurement.

The atomic radii are expressed in terms of covalent radii, metallic radii, Vander Waal’s radii and ionic radii.
For most of the elements, atomic radii is measured in terms of covalent radii whereas, Vander Waal’s radii is measured for noble gases.

1. Covalent radii: The half of the distance between two nuclei in a homonuclear diatomic molecule attached to single covalent bond is called covalent radii.

For example, the bond length of H2 molecules is 0.74 Å. According to definition, covalent radius is 0.74/2 = 0.37 Å.

2. Metallic radii: The half of the distance between two nuclei of atoms attached by metallic bond in metals is called metallic radii.

3. Vander Waal’s radii: The half of the distance between two non-bonded nuclei of atoms attached by Vander Waal’s force of attraction is called Vander Waal’s radii.

The strength of various bonds is:

Covalent bond > Metallic bond >> Vander Waal’s force of attraction.

Therefore, bond length increases in the order:

Covalent radii < Metallic radii << Vander Waal’s radii

Variation of atomic radii in the periodic table:

In a group:

In the group, from top to bottom the nuclear charge increases as well as there is increase in principle quantum number or number of shell (orbit).
Effect of adding the new shell is larger than that of increase in nuclear charge.
The effective nuclear charge per electron decreases.
Hence, atomic radius (size) increases.
Example : Atomic radii of alkali metals is: Li < Na < K < Rb < Cs

In period:

In the period, from left to right the nuclear charge increases and electrons are added to the same shell.
The effective nuclear charge per electron increases.
Hence, atomic radius (size) decreases.
Same period inert gas has highest atomic radii due to presence of Vander Waal’s radius.

Ionic radii: The ionic radius is the radii of the ions in crystal.

1. Cationic radius:

The cation is formed by removal of one or more electron from an atom.
The effective nuclear charge per electron in cation is more than the parent atom.
Hence, the size of cation is smaller than that of parent atom.

Example: Fe > Fe⁺² > Fe⁺³

2. Anionic radius:

The anion is formed by addition of one or more electron to an atom.
The effective nuclear charge per electron in anion is less than the parent atom.
Hence, the size of anion is larger than that of parent atom.

Example: O^-2 > O^– > O

Size of isoelectronic species:

Those species having same number of electrons but different nuclear charge are called isoelectronic species.
In isoelectronic species, size (radii) decreases with the increase in nuclear charge.

Example: N^-3 > O^-2 > F^– > Ne > Na⁺ > Mg⁺² > Al⁺³

Ionization energy (I.E.) / Ionization potential

The minimum amount of energy required to remove the most loosely bound electron from an isolated gaseous atom in its ground state to produce a cation is called ionization energy.

M (g) + I.E. → M⁺ (g) + e^–

Note: It is an endothermic process (∆H = +ve) and measured in electron volt (eV) or KJ/mole

Successive ionization energies:

The term ionization enegy (I.E.) is in place of first ionization energy. The energy required to remove second, third, and fourth electrons are called second, third and fourth ionization energies respectively and are denoted by IE₂, IE₃ and IE₄.

M (g) + IE₁ → M⁺ (g) + e^–

M⁺ (g) + IE₂ → M²⁺ (g) + e^–

M²⁺ (g) + IE₃ → M³⁺ (g) + e^–

Factors affecting I.E.

1. Atomic size: Ionization energy decreases with the increase in atomic size.

2. Nuclear charge: Ionization energy increases with the increase in nuclear charge.

3. Shielding or Screening effect: If the shielding or screening effect of the inner electrons increases then ionization energy decreases.

In multi-electron atoms, the electrons in the valence shell are pulled by the nucleus and repelled by the electrons of inner shells. Thus, outermost electrons experience less attraction from the nucleus under the combined effect of attractive and repulsive force acting on the valence electrons. This effect is called shielding effect or screening effect.
The actual charge felt by the valence electron as a result of shielding effect is called an effective nuclear charge (Z_eff).
With the increase in shielding effect, effective nuclear charge (Z_eff) decreases and hence I.E. decreases.

4. Penetration power of sub shell:

More penetrating (i.e. more closer) are the sub-shells of a shell to the nucleus more tightly the electrons are held by the nucleus and more is the I.E.
The penetrating power follows the order : s > p > d > f

5. Electronic configuration:

Half filled and completely filled orbitals are more stable than others and hence more energy is needed to remove an electron from such atoms. Thus, more stable the electronic configuration, the greater will be the I.E.

Inert gases have highest I.E. due to completely filled orbital. ‘He’ has highest I.E. in the periodic table.
Elements like Be (1s²,2s²) and Mg (1s²,2s²,2p⁶,3s²) have slightly higher I.E. due to extra stability of fully filled s-orbitals.
Elements like N (1s², 2s², 2p³) and P (1s²,2s²,2p⁶,3s²,3p³) have higher I.E. due to extra stability of half-filled p-orbitals.

Variation of Ionization energy in the periodic table:

In group: From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, ionization energy gradually decreases.

For example: Ionization energies of alkali metals is : Li > Na > K > Rb > Cs

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in I.E.

However, some elements show irregularities in the general trend. This may due to the extra stability of half-filled and completely filled electronic configurations.

For example: the variation of I.E. among the elements of II^nd period is:

Li < Be > B < C < N > O < F < Ne

Electron Affinity

The amount of energy released when an electron is added to an isolated gaseous atom in its ground state to form a gaseous anion is called electron affinity (EA).

X (g) + e^–→ X^– (g) + energy (EA)

Note: It is measured in electron volt (eV) or KJ/mole.

Factors affecting electron affinity:

Atomic size: Electron affinity decreases with the increase in atomic size.
Nuclear charge: Electron affinity increases with the increase in nuclear charge.
Screening effect: Electron affinity decreases with the increase in screening effect.
Electronic configuration: Elements having stable electronic configuration like half filled and completely filled orbitals have EA either very low or almost zero as they do not accept additional electrons so easily.

EA of inert gases is zero due to completely filled orbitals.
EA of alkaline earth metals is almost zero due to completely filled s-orbital.
EA of N, P is very low due to half filled orbital.

Variation of electron affinity in the periodic table:

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EA.

However, some elements show irregularities in the general trend. This may due to the extra stability of half-filled and completely filled electronic configurations.

For example: the variation of EA among the elements of II^ndand III^rd period is:

In 2^nd period : Ne < Be < N < Li < B < C < O < F

Halogens possess maximum electron affinity due to small size and maximum effective nuclear charge and after gaining one electron, they attain stable inert gas configuration.

In group: From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, EA gradually decreases.

For example: EA of alkali metals is : Li > Na > K > Rb > Cs

However, the electron affinity of fluorine is lower than chlorine. Due to small size of fluorine, the incoming electron feels more repulsion and less attraction. In case of chlorine, incoming electron feels less repulsion and more attraction than in fluorine. Hence, the EA of fluorine is lower than that of chlorine.

The EA order of halogens is : Cl > F > Br > I

This type of anomaly is also observed in chalcogens (group 16) i.e. S>O>Se>Te

Anomalous Electron affinity:

Electron affinity is …………….
Generally, From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EA. From top to bottom in a group, atomic size increases and shielding effect also increases . Hence, EA gradually decreases.
However, some exceptions in the general trend in EA are found, which is called anomalous EA. Some of the anomalies are:

1. Zero and very low EA : Elements having stable electronic configuration like half filled and completely filled orbitals have EA either very low or almost zero as they do not accept additional electrons so easily.

EA of inert gases is zero due to completely filled orbitals.
EA of alkaline earth metals is almost zero due to completely filled s-orbital.
EA of N, P is very low due to half filled orbital.

2. Halogens have highest EA: Halogens possess maximum electron affinity due to small size and maximum effective nuclear charge and after gaining one electron, they attain stable inert gas configuration.

3. Electron affinity of fluorine is lower than chlorine: Due to small size of fluorine, the incoming electron feels more repulsion and less attraction. In case of chlorine, incoming electron feels less repulsion and more attraction than in fluorine. Hence, the EA of fluorine is lower than that of chlorine.

The EA order of halogens is : Cl > F > Br > I

This type of anomaly is also observed in chalcogens (group 16) i.e. S>O>Se>Te

Eletronegativity (E.N.)

Electronegativity of an element is defined as the relative tendency of an atom in a molecule to attract shared pair of electrons towards itself.
It has no unit.
Higher the difference in electronegativity, more the polarity in the bond.

Factors affecting Electronegativity:

1. Atomic size: Smaller the size of an atom, the greater is its tendency to attract the shared pair of electrons towards itself. Hence,the electronegativity increases with a decrease in size of the atom.

2. Effective nuclear charge (Z_eff): EN increases with the increase in Z_eff.

3. Ionization energy and Electron affinity: Higher the value of I.E. and EA, higher is the EN.

4. Number and nature of atoms bonded to it: EN of an element depends upon the number and nature of the atoms to which it is bonded. For example, the EN of phosphorous in PCl₅ is higher than in PF₅, since fluorine is more electronegative than chlorine.

5. Type of hybridization: The EN increases as the s-character in hybrid orbital increases.

For example: EN of carbon in methane, ethane and ethyne is:

Ethyne > Ethene > Methane

6. Charge on the ion: Cation is smaller in size while anion has larger size as compared to that of parent atom. Hence, EN increases with the increase in +ve charge and decrease in negative charge.

For example:

Fe < Fe⁺² < Fe⁺³

O^-2 < O^– < O

Variation of electronegativity in the periodic table:

In group: From top to bottom in a group, atomic size increases and shielding effect also increases. Hence, EN gradually decreases.

For example: EN of alkali metals is : Li > Na > K > Rb > Cs

EN of halogens is : F > Cl > Br > I >At

In period: From left to right in a period, nuclear charge increases and atomic size decreases, so, there is gradual increase in EN.

For example: EN order of second period : Li < Be < B < C < N < O < F

Metallic character [Electropositive character]

The tendency of an element to lose an electron to form a cation is called electropositive character.
The electropositive character of metal is called a metallic character.
Lesser the value of ionization energy, more will be the metallic character and vice-versa.

Variation of metallic character in the periodic table:

In group: The metallic character of elements increases in going from top to bottom in a group. This is due to increase in size of the atom.

For example: Group 1: Li < Na < K < Rb < Cs

In period: The metallic character of elements decreases in going from left to right in a period. This is due to increase in effective nuclear charge.

For example: Period 3: Na > Mg > Al

Diagonal relationship

Similarities between certain pairs of elements that are diagonally adjacent to each other in the second and third period of the periodic table is called diagonal relationship.

Diagonal relationship is due to,

Almost same electronegativity (main cause)
Almost same atomic size.

Note: Bridge elements– Bridge elements are the elements of the second period of the periodic table. These elements show a relationship with the third-period, which are diagonal.