Periodic Classification of Elements

 

1.1 Historical Background of Periodic Classification

 

Need and Search for Classification

With the discovery of more and more new elements, it was necessary to organize these elements systematically and need arose for a frame work in which these elements could be classified and arranged in in order to facilitate their study and make their study simple and systematic. The classification of elements enabled the chemists to understand and interpret the properties of elements in a better way.

 

There could be many ways of arranging the elements; firstly they could be classified by their states (solids, liquids or gases) at a particular temperature, secondly they could be arranged as metals, non-metals and metalloids and thirdly one might find patterns in their reactions with oxygen or water or other chemicals. Would one consider trying to link these properties to the relative atomic masses of the elements?

 

Previously scientist tried to arrange the elements in a scientific, systematic and an organized manner on the basis of their atomic weight (atomic masses) as it was thought that the properties of elements depended upon their atomic masses (the thought was grounded on Dalton’s atomic theory). But recently, the basis of classification has been changed and elements are arranged on the basis of their atomic numbers instead of their atomic masses.

 

Different attempts of Classification

Following attempts were made to classify the known elements:

1.         Al-Razi Classification

2.         Origin of Classification; Dalton’s Atomic Theory

3.         Dumas Work

4.         Prout’s Attempt

5.         Dobereiner’s Triads

6.         Newland’s Law of Octave

7.         Lother Meyer’s Classification

8.         Mendeleev’s Classification

9.         Modern Periodic table

 

1.         Origin of Classification

The basis of classification of elements was grounded on the Daltons’ atomic theory put forward by an English scientist,  John Dalton  in 1808, according to which:

“Atoms of different elements have different atomic masses.”

 

Thus it was concluded that there is a regular relationship between atomic masses and properties of elements. “This relationship proves to be the corner stone for the future classification of elements”.

 

2.         Dumas Work

Dumas (1800-1884), a French chemist arranged the elements on their combining power with chlorine.

For example, elements that combined with 1 chlorine atom could be arranged in vertical columns in increasing order of their atomic weights and so on.

 

Reasons for Failure

Dumas attempt of classification did not gain success as all elements do not combine with chlorine and few show variable valency.

 

3.         Prout’s Attempt

Prout, an English chemist considered the atomic weight of hydrogen as the basis of his classification. He considered that:

“Atomic weights of all elements are simple multiple of the atomic weight of hydrogen”

 

Reason for Failure

It could not explain the fractional atomic weights of elements.

3.Dobereiner’s Triads

A German chemist, Johann Wolfgang Dobereiner in 1817 noticed an interesting pattern in certain sets of three similar elements and classified the similar elements in the groups of three elements (in the sequence of increasing atomic mass) known as triad. He found that the atomic mass of the middle element lay (fall) roughly half way (midway) between the other two (i.e. the lightest and the heaviest) elements of a triad and the elements of a triad also resemble in properties. He also noticed that the middle elements had properties that were an average of the other two members of a triad when arranged by the atomic weights.

 

e.g.

He found that the density of the middle element in most triad is roughly equal to the average of the densities of the other two elements. The density of strontium (2.6 g/cm³) for example is close to the average of the densities of calcium (1.55 g/cm³) and barium (3.51g/cm³).

 

He put forward Law or rule of Triads, according to which;

 

“Central atom of each set of triad has an atomic mass equal to the arithmetic mean of the atomic masses of the other two elements.”

OR

Each set of triad (group of three elements ordered by increasing atomic weights) has similar properties and atomic weight of the middle element of a triad was approximately equal to arithmetic mean (average) of the atomic weights of other two elements of a triad”.

 

He arranged the elements in triads. The elements of triad resemble in properties.

 

He first found alkaline earth metal triad of Ca, Sr and Ba.

He further noticed the same pattern for the alkali metal triad (Li, Na, K), the halogen triad (Cl, BR, I), Chalcogen (S, Se, Te), metalloid triad (P, As, Sb) and transitional metal triad (Mn, Cr, Fe).

 

Reason for failure

Dobereiner’s law of triad has a very limited application and could not be extended to the classification of all the elements as this rule was valid for only very few elements. It failed as this rule was not applicable for all elements i.e. all elements could not arrange in triads.

 

4.   Newland’s Law of Octave

In 1864, an English (London) industrial chemist John Alexander Newland arranged the 56 (60 or 62) known elements by order of increasing atomic weights into a table along horizontal rows seven element long with seven vertical columns and proposed has law of octave accordingly:

 

“If elements are arranged in the ascending order of their atomic weights, the eighth (8^th) element following any given element in the series has nearly same physical and chemical properties as first one” which means that starting from any element, the properties of every eighth element were similar to those of first

i.e. its properties are a kind of repetition of the first (like the eight notes of an octave of music or by the analogy with the seven intervals of the musical scale).

 

It was compared to octaves (Sa, Re, Ga, Ma, Pa, Da, Ni, Sa) in musical scale and thus the name Newland’s law of octaves (notes of music)

 

Merits

1.  It arranges all 56 elements into tabular form.

2. It arranges all elements with identical properties into same group.

3.    Newland’s classification of elements for the first time showed the existence of periodicity i.e. recurrence of chemical and physical properties of elements at regular intervals.

4.  It also provided a great idea towards the development of modern periodic table.

Objections

1.The Law of Octave holds up well for the first 16 (17) elements, but it failed rather badly beyond calcium  in predicting a consistent trend.

2. The heavier elements could not be accommodated by this arrangement.

3. Moreover hydrogen as not included in his table.

 

 4. Lother Meyer’s Classification

In 1869, a German Physicist Julius Lother Meyer (a contemporary of Mendeleev) classified the known 56 elements on the basis of their increasing atomic weights in graphical form in nine vertical columns or groups from I to IX.  Meyer’s work was based on physical properties of elements like atomic volume. He put forward his periodic law, which states that

 

     ‘‘physical properties of elements are periodic function of their atomic weights’’.

 

The volume occupied by 1 gram atomic weight or 1 gram atom or 1 gram mole (i.e. 6.02 x 10²³ atoms) of any element in solid state is called atomic volume which is a rough measure of the relative sizes of atoms.

Lother Meyer’s Atomic Volume Curve

Meyer arranged the elements by plotting a graph between atomic volumes of elements (on y-axis) against their increasing atomic masses (on x-axis).

 

The plot gave a curve called Atomic Volume Curve, consisted of sharp peak (crests) and broad minima (troughs). The curve exhibits periodicity as similar elements occupy same positions on the curve. For example, the highly reactive alkali metals (Li, Na, K, Rb, Cs) occupy the peak of the curve thereby showing that these elements have largest atomic volumes.

 

According to Meyer, the occupying of similar elements on same positions on the curve was called periodicity. The regular spacing of the highest points and occupying of similar elements on the same positions on the curve confirmed the idea of periodicity, suggested by Newland. [Meyer was the first scientist who considered valency as a period property.]

 

Meyer’s curve showed the following characteristics and periodicity:

 

1.    Chemically similar elements occupy similar position on the curves. For example; Alkali Metals like Li, Na, K etc. occupy the peaks of the curve indicating that they have largest atomic volumes than those of neighbouring elements while ascending portion of the curve just before the peak is occupied by halogens showing their smallest atomic volumes. The crest of each wave is occupied by an alkali metal and trough by an element of small chemical affinity.

 

2.     Alkali metals occupy the peaks or crests of the curves.

 

3.   Weak metals or elements of small chemical affinity or transition metals occupy the troughs or minima of the curve.

 

4.   Electronegative and gaseous volatile elements or acidic oxides forming elements are located on the ascending portions of the curve.

 

5.   Electropositive or transition elements or elements with high melting points are found on the descending portions of the curve.

 

6. Midway of ascending portions of curve is occupied by halogens.

 

7.  Midway of descending portions of curve is occupied by alkaline earth metals.

 

Summary of Meyer’s Atomic Curve

Meyer’s curve shows the following characteristics and periodicity:

Objections

Lother Meyer’s Periodic Classification could not receive proper attention due to following reasons:

1.   Meyer’s Periodic Table was incomplete as he left no blank spaces for undiscovered elements as compared with Mendeleev’s Periodic Table (which was characterized by remarkable predictions of discoveries of certain elements).

 

2.  no logical basis for classification based on various physical properties such as atomic volume.

 

3. Chemical properties of elements were completely ignored.

 

4.   His table was non-reproducible form of periodic table.

 

  

Mendeleev’s Classification

 

Most of the credit of the development of periodic classification of elements must go to a Russian chemist Dmitri Ivanovitch (D.I.) Mendeleev who presented the most useful and most systematic scheme for periodic classification of elements in March 1869. (Mendeleev’s was notorious for cutting his hair only once a year). Up till 1869, only 63 elements were known. Mendeleev arranged the elements in the sequence of their increasing atomic weights. He arranged the elements of similar properties in vertical columns and dissimilar elements in horizontal rows.

 

Basis of Classification

1.Increasing order of atomic mass of elements

2.Similarity in chemical properties of elements

 

Mendeleev’s work was an extension of Newland’s octaves. The basis of his classification was the chemical properties of elements. Mendeleev arranged the known 63 elements in the sequence of their increasing atomic weights, placing the elements with similar chemical properties vertically beneath each other. In his table, similar properties occurred periodically i.e. repeated themselves at intervals as a function of atomic weights.

 

Mendeleev’s Periodic Law

Since similar properties occurred periodically as a function of atomic mass, Mendeleev stated the Periodic Law as;

“The physical and chemical properties of elements are a periodic function of their atomic weights i.e. if the elements are arranged in ascending order of their atomic weights, their properties repeat in a periodic manner.”

 

Features of Mendeleev’s Periodic Table

Following are the main features of table:

 

1.   The elements are arranged in ascending order of their atomic masses.

2.    The Mendeleev’s periodic table consisted of 8 vertical columns called groups (i.e. group I to VIII) containing similar elements and 12 horizontal rows called Series or Periods having dissimilar elements.

3.   The groups are further divided into sub groups A and B. This sub division allowed him to place elements with slightly different properties in same group thereby maintaining periodicity.

4.  The elements in each group have similar chemical properties but their physical properties change gradually down the group.

5.   The group number indicates the highest valency of element that it can attain.

6.  Mendeleev’s table clearly and forcefully proved the concept of periodicity.

7.Mendeleev’s table contained vacant spaces for undiscovered (unknown or missing) elements with predicted atomic masses 44, 68, 72 and 100. He named them eka-boron, eka-aluminium and eka silicon.

 

Advantages (Merits) of Mendeleev’s Table

1. Systematic Study of Elements

2. Special Emphasis on Chemical Similarities

3. Determination & Correction of Doubtful Atomic Weights

4.  Separate Group for A set of 3 Elements

5.  Vacant Spaces for Undiscovered Elements & Prediction of Properties of 3 Unknown Elements

 

1.   Systematic Study of Elements

Mendeleev’s table helped chemist to study the elements more easily and systematically as it had reduced or restricted the study of elements into a study of eight groups only. (The study of chemistry of only one element of any group, is largely enough to predict the properties of the other elements of the same group). For instance, the study of sodium metal helped chemist to a large extent to predict the properties of its other group elements like K, Rb, Cs.

 

2.   Special Emphasis on Chemical Similarities/ Properties take precedence over atomic weights

Mendeleev disregarded atomic masses as the only criteria for assigning places to elements because at that time, the atomic weights of many elements were not accurately known; nor was it certain that all the elements has been discovered. Great emphasis was laid on chemical similarities of elements. Thus if the properties of an element suggested that it was out of place in the sequence of atomic masses, it was placed according to its properties rather than its mass.

He placed 3 misfit pairs of element in his table denying his own periodic law to maintain periodicity which was supposed to be more important than following law. Hence some elements of higher atomic mass were placed before elements of lower atomic mass. e.g.

 

3. Determination & Correction of Doubtful Atomic Weights

Mendeleev’s classification helped in correcting the doubtful atomic weights of a number of elements which has been assigned incorrect values and put them in proper places in the periodic table.

Mendeleev used the formula; Atomic weight  =  Equivalent weight  x  valency for calculating atomic weights of elements. The equivalent weight of elements can be calculated by any of the known methods and the valency can be obtained by consulting the periodic table.

For example:

(i)   atomic weight of Be was correct from 13.5 to 9. With this atomic weight, Be was given a position between  Li–7 and B–11.  The properties of Be justify this position in periodic table.

 

(ii)   Similarly atomic weight of indium was readjusted from 75.80 to 113. (Indium was supposed to have the valency 2 and equivalent weight 38, so its atomic weight would be 2 x 38 = 76 and with this atomic weight it would be placed between Zn-65 and Sr-87. There was no place between Zn and Sr. Mendeleev suggested if indium were taken as trivalent, its atomic weight would be 3 x 38 = 114 and thus would get the place between Cd – 112 and Sn -118 that justified its position).

 

(iii)   Also, atomic weight of Cr that had been an atomic weight of 43 was recalculated and found to be 52 and allocated proper place to it.

 

4. Separate Group for a set of 3 Elements in Group VIII

Mendeleev observed that a set of 3 elements i.e. Fe, Co, Ni, and Ru, Rh, Pd and Os, Ir, Pt had very similar properties and could not be assigned to any particular group. He, therefore, placed these elements at one place in Group VIII.

 

5. Vacant Spaces for Undiscovered Elements & Prediction of Properties of 3 Unknown Elements

In order to maintain families of chemically similar elements, he left blank spaces in his table for undiscovered elements after boron, aluminium and silicon which allowed his theory to be tested. Comparing the properties of their group elements, he successfully predicted the three unknown elements, which he named Eka-Boron, Eka-Aluminium and Eka-Silicon (eka means first i.e. eka-silicon means literally first comes silicon and then comes unknown element). This prediction helped in their discovery. By 1886, chemists had discovered all the three elements and had been named as scandium (Sc), gallium ( ) and germanium ( ) respectively. Mendeleev’s predictions of their properties proved to be remarkably accurate. e.g.

 

Comparative Properties of eka-Aluminium and Gallium

Comparative Properties of eka-Silicon and Germanium

 

 

Defects or Demerits or Limitations of Mendeleev’s Periodic Table

1.   Failure to explain atomic structure

2.   No place for isotopes of elements

3.   Anomalous position of hydrogen

4.   Failure to place rare earth (Lanthanides and Actinides) in the main body of periodic table

5.   Group number does not represent valency

6.   Neglection of variable valency

7.   Unable to give cause of Periodicity

8.   Anomalous or Misfit Pairs of elements

  (i)  Elements Cu, Ag, Au were placed with dissimilar elements Li, Na, K, Rb, Cs.

(ii) Similar elements Cu and Hg were placed separately.

(iii) Elements of higher atomic weight placed earlier than elements of lighter atomic weights.

 

1.   Failure to explain atomic structure

Mendeleev’s periodic table failed to account for atomic structure as it was based on atomic weight and not on atomic number. Also Mendeleev’s table was silent about electronic configuration of elements.

 

2.   No place for isotopes of elements

No separate position has been given to isotopes of an element having different atomic masses although the basis of classification is atomic mass. There was no room for isotopes in Mendeleev’s table as it was not possible to accommodate the large number of isotopes in the periodic table.

 

3.   Anomalous position of hydrogen

He could not assign a correct position to hydrogen in his table. The placement of hydrogen in group I along with alkali metals was a matter of dispute. The position of hydrogen was not justified.  

 

4.   Failure to place rare earth in the main body of periodic table

Lanthanides (elements with atomic numbers 58 to 71) and actinides (elements with atomic numbers 90 to 103) had not been placed in the main body of the periodic table. Rather they had been given a separate position at the bottom of the periodic table.

 

5.   Group number does not represent valency

Group number did not represent the valency of the elements e.g. excepting osmium, elements in group VIII did not show a valency of 8. Also the elements in the middle of the long periods (e.g. Mn, Cr etc.) exhibited variable valency.

 

6.   Neglection of variable valency

Elements with variable valencies were considered to have fixed valency.

 

7.   Unable to give cause of Periodicity

Cause of periodicity was not given by Mendeleev.

 

7.   Anomalos or Misfit Pairs of elements/ Wrong order of Some Elements

(i).        Dissimilar elements placed in the same group

Many elements with dissimilar properties had been placed in the same group e.g. Alkali metals and coinage metals were place in same group in spite of their entirely different properties. Also Mn had been placed with halogens. However division of groups into sub-groups solved the issue later.

(ii).Similar elements placed in different groups

Similar pairs of elements were placed in different groups. For instance Ba and Pb resemble in many properties but they were kept in different groups. Moreover, similar elements Cu and Hg were also placed separately.

(iii). Misfit position of elements of group VIII

Group VIII has 9 elements placed in three available columns. These elements did not fit in the system.

(iv)Position of 4 anomalous pairs of elements  

Increasing order of atomic mass could not be maintained. For placing elements in the proper groups, certain elements of higher atomic masses precede those of lower atomic masses in Mendeleev’s table. This was against Mendeleev’s Periodic Law. These misfit pairs of elements were Ar-K, Co-Ni, Te-I and Th-Pa

⁴⁰Ar – ³⁹K                     ⁶⁰Co–⁵⁹Ni                     ¹²⁷Te–¹²⁶I

 

1.2 Modern Periodic Table or Bohr’s Long Form of Periodic Table

 

Discovery of Atomic Number

Mendeleev’s Periodic Table based on atomic masses left many anomalies in the position of different elements in his table. Moreover the existence of isotopes showed that the atomic mass of an element is not the fundamental property of an element.

 

A British physicist, Henry Moseley in 1914 showed that frequency of X-rays emitted by different metal anodes varies directly with its number of protons (or electrons) or positive charge which is called its atomic number. [He showed by investigation of X-ray spectra of elements that this positive charge was in a definite amount and increased regularly for one element to the next by one unit. Thus if the charge of hydrogen nucleus is +1,then the relative charge on the nucleus of next element helium would be +2 and that on the nucleus of third element lithium would be +3, and so on.]

 

The X-ray spectra of elements showed that the physical and chemical properties of elements depend upon the number of electrons i.e. atomic number and their arrangement in different orbitals of the atom. Moseley pointed out that atomic number of an element is the fundamental property and properties of element are related to its atomic number and not their atomic weights. Thus Moseley predicted that most of the defects of Mendeleev’s table could be removed successfully if elements were arranged according to their atomic numbers rather than atomic weights. Thus Moseley modified the periodic law as:

 

Moseley’s Modern Periodic Law

“The physical and chemical properties of all elements are a periodic function of their atomic numbers

i.e. if the elements are arranged in order of their increasing atomic numbers, the properties of elements or similar elements are repeated after definite regular intervals.”

 

With replacement of basis of classification from atomic weight to atomic number, many inconsistencies and irregularities in the Mendeleev’s table disappeared.

 

Basis of Classification

The method of arranging similar elements in one group and separating them from dissimilar elements placing them in periods or horizontal rows based on periodicity of elements is called Periodic Classification of elements. It is so named as it is grounded on periodic recurrence of physical and chemical properties of elements i.e. periodicity.

 

The periodic classification of elements is based on periodicity, due to which the elements having similar properties are repeated at regular intervals. When elements are arranged in ascending order of their atomic numbers, their properties show a repeating pattern after intervals of 2,8,8,18,18 and 32. This is called periodicity in properties. The repetition or recurrence of similar properties among elements after specific intervals or periodically due to repetition of similar valence shell electronic configuration is called periodicity.

 

Chemical properties of elements depend upon the number of valence electrons, hence the elements with similar valence shell electronic configuration tend to show similar chemical behaviour. As atomic number is related to the number of protons and number of electrons in an atom, so the real basis of periodicity of properties is due to recurrence of identical valence shell electronic configuration of the next element in the same group after regular intervals of 2,8,8,18,18 and 32 in atomic numbers].

                                                                                                                                        

Moseley’s Modern Periodic Law in terms of electronic configuration

Now, Moseley’s Periodic Law may be restated as:

 

“The physical and chemical properties of elements are periodic function of the electronic configuration of their atoms which vary with increasing atomic number in a periodic manner”.

Long Form of Periodic Table or Bohr’s Long Form of Periodic Table

The periodic table is an orderly arrangement of the known chemical elements in a tabular form in which elements are placed in the increasing order of their atomic number or electronic structure (configuration) so that many chemical properties vary regularly across the table.

 

The modern periodic table is the result of discovery of atomic number by Moseley in 1914. The modern periodic table based on Mosley’s Modern Periodic Law showing periodicity grounded on Bohr’s scheme of classification of elements into 4 types depending on the number of incomplete shells of electrons in the atom, was proposed by Rang (1893) then modified by Werner (1905) and extended by Bury (1921) and adopted by IUPAC in 1984 is called Bohr’s form or Long Form or Extended form of Periodic Table consisting of groups and periods (because it contains 16 groups or 18 vertical columns rather than 8 and 7 periods instead of 12) .

 

With replacement of basis of classification from atomic weight to atomic number, many inconsistencies and irregularities in Mendeleev’s table disappeared.

 

Difference between Mendeleev’s periodic table and Modern periodic table

1.         Modern periodic table is based on the most fundamental property, atomic number of elements, while Mendeleev’s periodic table is based upon the atomic masses of elements.

 

2.  Modern periodic table explains clearly why elements in a group display similar properties and elements of a group differ in properties from elements of other groups. Mendeleev’s periodic table failed to do so.

 

3.   In Mendeleev’s periodic table, there are several anomalies e.g. the position of isotopes, wrong order of atomic masses of some elements etc. In the modern periodic table, these anomalies have been removed.

 

4.   In the long form of the peridoc table, elemetns have been cleraly separated as representative elements, transtion elements and noble gases. Metals and nonmetals are aslo separated. But in Mendeleev’s periodic table there is no such separation of different types of elements.

 

5.   In the modern periodic table the subgroups A and B are clearly separate because the elements belonging to subgroup A differ in properties from those of elements belonging to subgroup B. In Mendeleev’s periodic table, the two subgroups are kept together.

 

Applications of Modern Periodic Table

1. Prediction of Properties of element in a group.

2.Prediction of Molecular formula of compounds (between elements of different groups).

3.Prediction of new or unknown elements has been possible.

4.  Visualization of Reactivities of elements.

5.Suggestions for further research become available.

 

Merits or Advantages of Long Form of Periodic Table

1. Placement of elements according to fundamental property, atomic number

2. Relation of Properties of Element with Electronic Configuration

3. Defects of Mendeleev’s Table Removed

4. Position of Isotopes Solved

5. Exhibition of Distinct Periodicity and Cause of periodicity

6. Demonstration of various types of elements

7. Division of Elements into Four Blocks

8. Separate Position of A and B Groups

9. Showing Trend in Chemical Properties

10. Simple and Easy Study of Elements

11. Memorable and Reproducible Form of Periodic Table

12. Prediction of Properties of element in a group

Demerits of Modern Periodic Table

1.         Controversial Position of Hydrogen

2.         Disputed Position of Helium

3.         Controversial Position of Rare Earths

4.         Three Columns in group VIIIB

5.         Gaps in the Periodic Table

6.         Some Properties Neglected

 

Merits of Long Form of Periodic Table

 

1.   Placement of elements according to fundamental property, atomic number

The modern periodic table is based on more fundamental property, atomic number. It relates to fundamental property i.e. atomic number.

 

2.   Relation of Properties of Element with Electronic Configuration

      It relates the position of an element to its electronic configuration.

It explains why all the elements in a group have identical chemical properties while the elements in a period have different chemical properties.

 

All the elements in a group have similar properties because they have similar valence shell electronic configuration. On the other hand, all the elements in a period have different properties because thyme have different valance shell electronic configuration due to progressive addition of electrons to the valence shell on moving across a period.

 

3.   Defects of Mendeleev’s Table Removed

      The anomalous pair of elements i.e. Ar-K, Co-Ni, Te-I and Th-Pa are found arranged rightly in the table when they are placed in the order of increasing atomic numbers.

 

4.   Position of Isotopes Solved

Different isotopes of an element have been occupied one and the same place in the periodic table as they have same atomic numbers.

 

5. Exhibition of Distinct Periodicity and Cause of periodicity

It clearly exhibits periodicity in properties of elements i.e. recurrence of elements with similar properties. According to modern periodic table cause of periodicity is recurrence of elements with similar outer shell configuration.

 

6.   Demonstration of various types of elements

It clearly illustrates active metals, non-active metals, transition metals, metalloids, non-metals & noble gases.

 

7.   Division of Elements into Four Blocks

It divides the elements into 4 blocks i.e. s, p, d and f-blocks.

 

8.   Separate Position of A and B Groups

      The elements of the two sub-groups have been placed separately and thus dissimilar elements do not fall together.

 

9.   Showing Trend in Chemical Properties

It clearly brings out the trend in chemical properties in a period and group.

 

10. Simple, systematic and Easy Study of Elements

It makes the study of the properties of elements (and their compounds) simple and easy. It systematizes the study of elements. In the periodic table, elements with similar properties have been placed in the same vertical columns or groups. If we know the properties of one element of the group, the properties of other elements in the same group can be predicted. Thus, there is no need of studying the properties of all the elements.

 

11. Memorable and Reproducible Form of Periodic Table

      It is easy to remember, understand and reproduce.

12. Prediction of Properties of element in a group

It is possible to predict the properties of an element from the location of the element in the periodic table. For example, if the element belongs to the group IA, it is likely to be a reactive metal. If the element is the last element of the period, it would be a gas which is almost inert.

Demerits of Modern Periodic Table

1.   Controversial Position of Hydrogen

2.   Disputed Position of Helium

3.   Controversial Position of Rare Earths

4.   Three Columns in group VIIIB

5.   Gaps in the Periodic Table

6.   Some Properties Neglected

 

1.   Controversial Position of Hydrogen

Position of hydrogen in Group IA is disputed and thus its exact position is yet not decided and still remains unsolved.

 

2.   Disputed Position of Helium

 Position of helium in VIIIA group is controversial (configuration of He is 1s² whereas the configuration of other noble gases is ns²np⁶).

 

3.   Controversial Position of Rare Earths

      Lanthanides and actinides have still not been adjusted in the main body of the periodic table.

 

4.   Three Columns in group VIIIB

      Group VIIIB consists of three columns.

 

5.   Gaps in the Periodic Table

      There are large gaps in the periodic table e.g. group IIA is widely separated from group IIIA.

 

6.   Some Properties Neglected

   Some properties of elements such as specific heats have no relationships with the periodic classification.

 

 

1.3 Groups, their Sub-Division and their General Features

Definition and Sub-Division

The vertical columns of elements in the periodic table are called Groups. A group consists of a set or series of elements having identical valance shell configuration. Periodicity of properties of elements gives rise to groups of the periodic table.

 

All the elements belonging to the same group constitute a family. All the elements belonging to a particular group have same number of a valence electrons and hence exhibit similar properties.

 

Total number of groups

There are 18 vertical columns in the modern periodic table, so there are 18 groups in the modern periodic table which are numbered from 1 to 18 according to recommendations of IUPAC.

 

Earlier, the designation of these groups was the same as in the Mendeleev’s periodic table. Thus formerly, there were eight groups (I to VIII) but each group is further sub-divided into A and B sub-groups. But the total groups including A and B sub-groups are 16 as group VIIIB consists of three columns.

 

The relationship between the two ways of numbering the group is given below:

 

Groups are numbered by Roman numericals as IA, IIA, IIIB to VIIIB (comprising of three columns), IB, IIB, IIIA to VIIIA or zero group (or by simple numericals as 1, 2, 3, ……….. 16, 17, 18).

 

Types of Groups

Groups are divided into A-family and B-family. The elements of A-family are chemically different from the elements of B-family.

 

The elements of sub-group A or the elements of groups 1, 2, 13, 14, 15, 16, 17, and 18 are called Main group or Normal or Major or Representative Elements as the properties of these elements are represented by valence electrons. These elements have all their inner shells complete. Their only the outermost shell is incomplete. In these elements the outermost shell gets progressively filled from group 1 to group 18 as we move from left to right in a period. The elements of group IA and IIA are called s-block elements (which include active metals) while the elements of group IIIA to VIIIA are called p-Block elements (which include all non-metals, metalloids and weak metals). The group VIIIA is also called zero group which contains noble gases.

 

The elements of sub-group B or the elements of groups 3 t 12  are called Transition (or outer transition) elements because the properties of these elements show a gradual change or transition between the two sets of representative (s and p-block) elements, on either side of them. In these elements the outermost and the penultimate (next to outermost) shells are incomplete. All transition elements are metals.

 

Lanthanides and actinides are collectively known as inner transition elements. In these elements, the outermost three shells are incomplete.

 

General Characteristics of Groups

1. Representation of Total Valence Electrons

2. Representation of Maximum Valency and the highest Oxidation State

3. Exhibition of Identical Valence Shell Electronic Configuration

4. Exhibition of Same Chemical Properties

5. Exhibition of Regular Gradation in Physical Properties

6. Different Behaviour of First Congeners of Each Group

7. Increasing Electropositivity and Decreasing Electronegativity

8. Identical number of Valence electrons and valence orbital on descending a group

9. Difference in Properties of Sub-groups A and B

10. Representative and Transition elements

1. Representation of Total Valence Electrons

Group number of an element represents the total number of valence electrons in its valence shell e.g. oxygen belongs to VIA group as it has six valence electrons.

 

2.Representation of Maximum Valency and the highest Oxidation State

Group number of an element represents its maximum valency and the highest oxidation states. (Group number of an element is equal to its valency with respect to oxygen).

 

3. Exhibition of Identical Valence Shell Electronic Configuration

Elements of a same group have identical valency shell electronic configuration.

4. Exhibition of Same Chemical Properties

Elements in same group show same chemical properties due to same valency shell configuration.

 

5. Exhibition of Regular Gradation in Physical Properties

Elements of a group show regular gradation (change) in its physical properties on descending a group due to gradual change in their atomic sizes and electronegativities.

 

6.  Different Behaviour of First Congeners of Each Group

The first member of each group shows slightly different behaviours from other members of that group due to its small atomic size.

 

7.   Increasing Electropositivity and Decreasing Electronegativity

Electropositivity (metallic character) increases while electronegativity decreases down each group with increasing atomic numbers due to increasing atomic size.

 

8.  Identical number of Valence electrons and valence orbital on descending a group

On moving down a group, the number of shells increase but the total number of valence electrons and valence orbital remain same.

 

9.  Difference in Properties of Sub-groups A and B

The elements of A-family are chemically different from the elements of B-family.

 

10. Representative and Transition elements

The elements of family A (IA to VIIA) are called normal or representative elements having ns¹ to ns² np⁵ valence shell electronic configuration. Elements of group VIIIA are called inert or noble gases with valence shell configuration of ns² np⁶.

 

The elements of family B (IB t VIIIB) are called outer transition elements or d-block elements as outer electrons fall in (n-1)d-orbital.

 

Groups Learning Key/ Mnemonic for learning groups

 

 

 

 

Group IA or Lithium Family (Alkali Metals)

Members

This group includes lithium (₃Li), sodium/natrium (₁₁Na), potassium/kalium (₁₉K), rubidium (₃₇Rb), cesium (₅₅Cs) and francium (₈₇Fr). Francium is radioactive. Sodium and potassium are the 6^th and 7^th most abundant elements in the earth crust with % abundance of 2.6% and 2.4% respectively.

Groups Learning Key/ Mnemonic for group IA

            

                    

General Characteristics

1.   Electropositive Nature, High Reactivity, Solid State and Low Volatility

2.   valence shell electronic configuration of ns¹ , monovalent Nature, Fixed Oxidation state of +1

3.   Formation of Monovalent Cation

4.   Reducing Behaviour

5.   Forming only Ionic bonds

6.   Basic Nature of Oxides

7.   Trend of Physical Properties

 

1.   Electropositive Nature, High Reactivity, Solid State and Low Volatility

They are highly reactive and strongly electropositive metallic elements relatively soft solids having low melting and boiling point.

 

2.   valence shell electronic configuration of ns¹ , monovalent Nature, Fixed Oxidation state of +1

They are associated with valence shell electronic configuration of ns¹ (where n is the number of orbits ranges 2-7) showing that they contain only 1 valence electron, so they are monovalent, exhibiting a fixed oxidation state of +1.

 

3.   Formation of Monovalent Cation

They have tendency to lose their single valence electron on reaction to get respective inert gas like configuration of the previous period to form monovalent positive ion (M⁺) showing their electropositive character.

4.   Reducing Behaviour and Low I.P

They are powerful reducing agent due to their low ionization energies.

 

5.   Forming only Ionic bonds

They can form only ionic bonds.

 

6.   Basic Nature of Oxides

They themselves, their oxides, hydroxides, hydrides, peroxides are basic in nature and when dissolve in water forming alkalis. That is why they are known as alkali metals.

7.   Trend of Physical Properties

They have largest atomic size in their respective period. Their atomic radii, ionic radii, atomic volumes increase down the group from Li to Cs due to the addition of extra shell to each element and due to same reason melting and boiling points decrease while electropositivity increases downward.

  

        Group IIA or Beryllium Family (Alkaline Earth Metals)

Members

This group comprises of Beryllium (₄Be), magnesium (₁₂Mg), calcium (₂₀Ca), strontium (₃₈Sr), barium (₅₆Ba) and radium (₈₈Ra). Radium is radioactive. Calcium is the 5^th most abundant element with % abundance of 3% while magnesium is the 8^th most abundant element with % abundance of 2% in the earth crust.

Groups Learning Key/ Mnemonic for group IIA

 

General Characters

1.   Electropositive Nature, High Reactivity, Solid State and Low Volatility (but less than alkali metals) 

2.   valence shell electronic configuration of ns², divalent Nature, Fixed Oxidation state of +2

3.   Formation of Monovalent Cation

4.   Reducing Behaviour

5.   Forming only Ionic bonds

6.   Basic Nature of Oxides

7.   Trend of Physical Properties

 

1.   Electropositive Nature, High Reactivity, Solid State and Low Volatility (but less than alkali metals)  

They are reactive and electropositive metallic elements but less reactive and less electropositive than alkali metals, a bit harder having relatively high melting and boiling points than the alkali metals.

 

2.   valence shell electronic configuration of ns², divalent Nature, Fixed Oxidation state of +2

They have valency shell electronic configuration of ns² showing that they contain two valence electrons, so they are divalent exhibiting a fixed oxidation state of +2.     

 

3.   Formation of divalent Cation

They have tendency to lose their both valence electron on reaction to get respective inert gas like configuration of the previous period to form divalent positive ion (M²⁺) showing their electropositive character.

 

4.   Reducing Behaviour and high hydration energy

They are powerful and more stronger reducing agents than alkali metals because of high hydration energy of M²⁺ ions.

 

5.   Forming only Ionic bonds except Be

They form ionic bond except Be and Mg (however Mg can form some ionic compounds like MgO, MgSO₄).

 

6.   Basic Nature of Oxides

Their oxides are basic giving weak alkaline solution on dissolution in water. That is why they are called alkaline earth metals as these metals exist as their oxides (lime; CaO, strontia; SrO; Baryta; BaO) in the earth’s crust and are alkaline in nature.

7.   Trend of Physical Properties

They have relatively smaller atomic radii, ionic radii and atomic volumes due to their greater nuclear charge. However down the group they do not show a systematic and regular trend in melting points, boiling points and densities.

 

Group IIIA or Group 13 (Boron Family/Triels)

Members

This group includes boron (₅B), aluminium (₁₂Al), gallium (₃₁Ga), indium (₄₉In) and thallium (₈₁Tl). Boron is a metalloid showing dual characteristics of both metals and non-metals while rest of them are weak metals. Aluminium is the 3^rd most abundant element in the earth’s crust with % of abundance of 7% (7.6%).

Groups Learning Key/ Mnemonic for group IIIA

General Characteristics

 

1.   Electropositive Nature, High Reactivity, Solid State and Low Volatility (but less than alkali metals) 

2.   valence shell electronic configuration of ns² np¹, trivalent Nature, Common Oxidation state of +3

3.   Formation of Monovalent Cation

4.   Reducing Behaviour

5.   Forming only Ionic bonds

6.   Acidic Nature of Oxides

7.   Trend of Physical Properties

 

1.   Electropositive Nature, Low Reactivity, Solid State and Low Volatility

Except boron, they are highly electropositive elements showing metallic character which increases down the group due to increase in atomic size (or atomic volume) relatively hard having high melting and boiling points (except Ga with m.p= 29.9^○C).

 

2.   valence shell electronic configuration of ns² np¹, trivalent Nature, Common Oxidation state of +3

They are associated with ns² np¹ valency shell configuration i.e. they contain 3 valence electrons, so they are trivalent showing a valency of 3 exhibiting a most common oxidation state of +3. Later members also show 1 valency and +1 oxidation state due to inert pair effect.

 

3.   Formation of Trivalent Cation

Besides boron, they have tendency to lose three valence electrons acquiring noble gas configuration to form trivalent positive (M³⁺) ions showing their metallic behaviour.          

           M_(g) ¾¾® M³⁺(g) + 3e^–     (except boron)

4.   Reducing Behaviour

They are reducing agent especially aluminium powder.

 

5.   Forming only Covalent bonds

They preferably form covalent bond. However some ionic compounds of aluminium are known like Al₂O₃, Al₂(SO₄)₃.

 

6.   Basic Nature of Oxides

They mostly form acidic oxides except Al which forms amphoteric oxide. 

 

Group IVA or Group 14 (Carbon Family/Tetrels)

Member

This group includes carbon (₆C), silicon (₁₄Si), germanium (₃₂Ge), tin or stannum (₅₀Sn) and lead or plumbum (₈₂Pb). Of these elements carbon is a typical non-metal, silicon and germanium are metalloids and tin and lead are metals.

Carbon is the 16^th (14^th or 17^th in some books) most abundant element in the earth crust (0.18%) and in human body carbon is the 2^nd most abundant element (18%). Silicon is the 2^nd most abundant element (26%) in the earth crust.

 

In this group, there is smooth transition from non-metal to metal through metalloid. This group occupies the middle part of the periodic table and forms a link between more electropositive and more electronegative elements.

Mnemonic For Group IVA

General Characteristics

1.   Nature, Low Reactivity, Solid State and Low Volatility

2.   valence shell electronic configuration of ns² np²

3.   Formation of Cation

4.   Reducing Behaviour

5.   Forming only Ionic bonds

6.   Acidic Nature of Oxides

7.   Trend of Physical Properties

 

1.   Nature, Low Reactivity, Solid State and Low Volatility

They all are monoatomic solids elements having high melting and boiling points.

 

2.   valence shell electronic configuration of ns² np²

They are associated with valency shell configuration of ns² np² showing that they contain 4 valence electrons and so they are mostly tetravalent showing a valency of 4. Sn and Pb exhibit a variable valency of 2 and 4 due to inert pair effect. Carbon exhibit a variety of oxidation states in organic compounds like –4, –2, –1, +2, +4, 0 etc.

 

3.   Electronegative and electropositive character and Formation of Cation

Only carbon can form anions like carbide ion (C^4–) or dicarbide ion (C^2– or [Cº]^2–). Ge, Sn and Pb form divalent and tetravalent cations like Sn²⁺, Sn⁴⁺, Pb²⁺, Pb⁴⁺ due to inert pair effect.

 

4.   Forming both covalent and Ionic bonds

First three elements C, Si and Ge form covalent compounds while Sn and Pb preferably form ionic compounds. The nature of the compounds M²⁺ and M⁴⁺ cations can be predicted by Fajan’s rule which states that smaller the cations, the greater would be the covalent character. In general, compounds M⁴⁺ are covalent while that of M²⁺ are ionic in nature.

 

5.   Exhibition of allotropy

Except lead, all elements exhibit the property of allotropy e.g. carbon exists in a variety of allotropic forms like crystalline forms such as diamond, graphite and amorphous forms such as coal, etc. Silicon exists in crystalline and amorphous forms. Tin is found as grey tin (diamond type structure), white tin (tetragonal crystals) and brittle tin (rhombic crystals).

 

6.   Acidic Nature of Oxides

Oxides of C and Si (CO₂ and SiO₂) are acidic.

 

7.   Trend of Physical Properties

Down the group, atomic radii and atomic volumes increase due to addition of a new shell and from the same reason metallic character increases down the group. Thus Sn and Pb are typical metals.

Group VA or Group 15 or Nitrogen Family (Pnictogens/Pnicogen)

Member

This group contains nitrogen (₇N), phosphorus (₁₅P), arsenic (₃₃As), antimony or stibium (₅₁Sb) and bismuth (₈₃Bi). Among these elements, N and P are non-metals, As and Sb are metalloids and Bi is a metal.

 

The group VA elements are also called pnictogens (pnico=suffocation ; gens = producing) due to specific smell of nitrogen and other. The term pnictogen (or pnicogen) is derived from the Ancient Greek word pnigein meaning "to choke or stiffle", referring to the choking or stifling property of nitrogen gas in absence of oxygen. It can also be used as a mnemonic for the two most common members, P and N. The term "pnictogen" was suggested by the Dutch chemist Anton Eduard van Arkel in the early 1950s. It is also spelled "pnicogen".

 

Nitrogen is the 10^th most abundant element (0.6%) in the earth crust constituting fourth-fifths (4/5) of the air (78%).

Phosphorus is the 12^th most abundant element in the earth crust (0.2%).

 

Both N and P are essential to living organisms.

 

This group is often selected for systematic studies because among its members there is essentially a regular change with atomic weight and size from characteristics of a true non-metal (N) to a typical metal (Bi). There is a large variation of properties in going down the group.

Mnemonic For Group VA

 

General Characteristics

 

1.   Nature, State, volatility and atomicity

Nitrogen exists as diatomic molecules (N₂), phosphorus as tetraatomic molecule (P₄) while rest of them exists in monoatomic form. Nitrogen is gas while all other members are solid.

 

2.   valence shell electronic configuration of ns² np³

They have valency shell configuration of ns² np³ i.e. they have total 5 valence electrons exhibiting a variable valency of 3 and 5 so they are mostly trivalent or pentavalent except Bi which exhibit fixed valency of 3. However nitrogen shows a variety of valencies of 1, 2, 3, 4 and 5. Phosphorus also exhibits more than one valency 1, 3, 4, 5.

 

3.   Electronegative and electropositive character and Formation of Cation

Only nitrogen and phosphorus have tendency to gain 3 electrons to form nitride ion (N^3–) and phosphide ion (P^3–) respectively due to their small atomic size and large ionization potential.

 

4.   Forming both covalent and Ionic bonds

They preferably form covalent bond. However metallic nitrides and phosphides are mostly ionic.

 

5.   Acidic Nature of Oxides

They form all the three types of oxides. Nitrogen forms a variety of oxides which are either acidic or neutral e.g. NO, N₂O, NO₂, N₂O₄, N₂O₃ and N₂O₅.

 

6.   Exhibition of allotropy

All of the elements except nitrogen exhibit the property of allotropy e.g. P has two allotropes mainly white and red phosphorus.

Group VIA or Group 16 or Oxygen Family (Chalcogens)

Members

This group consists of oxygen (₈O), sulphur (₁₆S), selenium (₃₄Se), tellirium (₅₂Te) and polonium (₈₄Po). Of these elements, O and S are non-metals, Se and Te are metalloids and polonium is metal.

This group is also known as chalcogens meaning the ore-forming elements because most the ores of metals occur in nature as oxides and sulphides.

oxygen is the most abundant element in the earth crust (50%) and it is also the most abundant element in the human body (65%) while constituting one-fifth (1/5) of the air (21%).

Sulphur is the 16^th most abundant element (0.04%).

Both oxygen and sulphur are essential to living organisms.

There is a large variation of properties in going down the group.

 

Mnemonic For Group VIA

General Characteristics

1.   Nature, state, volatility and atomicity

2.   valence shell electronic configuration of ns² np⁴

3.   Electronegative and electropositive character and Formation of Cation

4.   Forming both covalent and Ionic bonds

5.   Acidic Nature of Oxides

6.   Oxidizing property

7.   Exhibition of allotropy

8.   Trend of Physical Properties

1.         Nature, state, volatility and atomicity

Oxygen is a gas found as diatomic molecule (O₂), other memebrs are solids which exist as octamolecule (S₈, Se₈) or monoatomic form (Te, Po). However liquid sulphur may exist in S₈, S₆, S₄ forms.

 

2.   valence shell electronic configuration of ns² np⁴

They have valency shell configuration of ns² np⁴ showing that they have 6 valence electrons exhibiting a variable valencies of 2, 4 and 6 (except oxygen which cannot exceed its valency to 2). Most common valency is however 2. Their most common oxidation state is -2. But they also show variable oxidation states.

 

3.   Electronegative and electropositive character and Formation of Cation

Only oxygen and sulphur have tendency to gain 2 electrons to form bivalent anions namely oxide (O^2–) and sulphide (S^2–) ions respectively due to their small atomic size and high electron affinity. Oxygen can also form peroxide [O₂^2–] and superoxide ion [O₂^1–]

 

4.   Forming both covalent and Ionic bonds

They preferably form covalent bond. However metallic oxides, peroxides, superoxides and metallic sulphides are mostly ionic.

 

5.   Acidic Nature of Oxides

They form acidic oxides. E.g. SO₂, SO₃, SeO₂ etc.

 

6.   Oxidizing property

They are oxidizing agents.  

 

7.   Exhibition of allotropy

All the elements exhibit the property of allotropy e.g.

(i) oxygen has two allotropic forms namely ordinary molecular oxygen (O₂) and trioxygen or ozone (O₃).

(ii) Similarly sulphur has a number of different allotropes like rhombic, monoclinic and plastic sulphur.

 

8.   Trend of Physical Properties

Metallic character, ionic and basic nature increase regularly down the group.

Group VIIA or Group 17 or Halogens (Fluorine Family)

 

Members

This group comprises of fluorine (₉F), chlorine (₁₇Cl), bromine (₃₅Br), iodine (₅₃I) and astatine (₈₅At). Except astatine which is a metal and radioactive all others are non-metals.

They are called halogens, a term which means salt formers because they form salts with metals called Halides.

Mnemonic for group VIIA      

 

General Characteristics

1.   Electronegative Nature, High Reactivity, any State and high Volatility

2.   valence shell electronic configuration of ns² np⁵

3.   Electronegative character and Formation of Univalent anion 

4.   Oxidizing Behaviour

5.   Forming Ionic bonds and covalent bonds

6.   Acidic Nature of Oxides

7.   Trend of Physical Properties

 

1.   Electronegative Nature, High Reactivity, any State and high Volatility

They are highly reactive and strongly electronegative non-metallic elements (active non-metals).

 

At room temperature fluorine and chlorine are coloured gases, bromine is a volatile liquid and iodine is a dark coloured sublime solid.

 

They exist as diatomic molecules i.e. F₂, Cl₂, Br₂, I₂ except At. They are found as discrete molecule held together by van der Waal’s forces which accounts for their volatile nature.

 

2.   valence shell electronic configuration of ns² np⁵

They are associated with valency shell configuration of ns² np⁵ showing that they have 7 valence electrons exhibiting a most common valency of 1 and are mostly univalent. However except F, all exhibit a variety of valencies of 1, 3, 5, 7. Their most common oxidation state is -1. However they can also show positive oxidation states of +1, +3, +5 and +7 except fluorine.

 

3.   Electronegative character and Formation of Univalent anion 

They have tendency to accept an electron easily attaining the next noble gas configuration to form univalent halide ions (i.e. X^1– e.g. F^–, Cl^–, Br^–, I^–) due to their high ionization energies and large negative electron affinities. 

4.         Oxidizing Behaviour

They are powerful oxidizing agent due to their smallest atomic radii and high electron affinities.

 

5.   Forming Ionic bonds and covalent bonds

They can form ionic as well as covalent bonds.

 

6.   Acidic Nature of Oxides

Their oxides are acidic in nature. The strength of acidic nature of oxides increases with the increase in oxidation state of halogen but decrease on descending a group.

7.         Trend of Physical Properties

They have smallest atomic size in their respective period.

Group VIIIA or Group 18 or Zero Group or Aerogens (Inert or Noble Gases)

Members

This group includes helium (₂He), neon (₁₀Ne), argon (₁₈Ar), krypton (₃₆Kr), xenon (₅₄Xe) and radon (₈₆Rn). Radon is radioactive.

The clue of existence of noble gases was provided by Cavendish in 1785.

The first noble gas discovered by an English scientist Ramsay (and Raleigh) in 1892 (1894) was argon (Greek meaning idle or lazy) from air.

In the same year, Ramsay isolated the lightest of all noble gases helium (meaning the Sun) from uranium ores.

During 1898, Ramsay and Rayleigh and travers isolated three additional noble gases, Neon (new), Kr (hidden) and Xe (stranger) from air.

Noble gases are found in the atmosphere in very small quantities. As these elements were not known at the time of Mendeleev, so no place was kept for them in the periodic table. They had, therefore, been placed in an additional group called zero group which was inserted in between the most electronegative halogens and group VII and the most electropositive alkali metals of group I.

 

% of Noble gases in Air

Mnemonic for group VIIIA     

General Characters

1.   Nature, State, Volatility, Atomicity

2.   valence shell electronic configuration of ns² np⁶ and valency

3.   No electropositive or electronegative Character

4.   Inertness

1.   Nature, State, Volatility, Atomicity

They are monoatomic and low boiling point, diamagnetic, colourless, odourless and tasteless gases. They do not resemble with any other elements in the periodic table either to the left or right. Thus they act as a bridge between electronegative and electropositive elements in the periodic table.

 

2.   valence shell electronic configuration of ns² np⁶ and valency

They are associated ns² np⁶ (where n = 2 – 6) outer shell electronic configuration is i.e. they have 8 valence electrons showing that they have completely filled outer shell (or fully filled s and p sub-shells or orbitals) except helium which has only two valence electrons (i.e. 1s²).

All noble gases are non-valent elements (i.e. possess zero valency indicating that they have no combining tendency with other elements) which is attributed due to completely filled valence shell.

 

3.No electropositive or electronegative Character

They have complete valence shell in the form of either complete octet (ns² np⁶) or complete duplet for He (1s²) of electrons in their valence shell. Due to completely filled valence shell, they are exceptionally stable and are unable to gain or lose electron due to zero electron affinity and very high I.P respectively. No atom has complete outer shell with the exception of He and Ne).

All atoms tend to get inert gas like configuration in order to acquire stability. Thus they are used as standard for comparing electronic configuration of elements.

 

4.   Inertness

They are chemically inert or non-reactive due to completely filled outer s- and p-sub-shells (orbitals). Their inertness is attributed due to their small atomic volumes, high I.P and zero electron affinity.

The word inert is strictly used for He, Ne and Ar but should not be used for Kr, Xe and Rn as they have large atomic volumes and thus form few compounds like KrF₂, XeO₂ etc under drastic conditions. 

 

Periods, their General Features and their types          

Definition of Periods

The horizontal rows of elements in periodic table arranged in the ascending order of their atomic numbers are called Periods which are designated by simple numericals. In each period, the elements have been placed in the increasing order of their atomic numbers.

 

There are 7 periods in the periodic table. Period second and third are called short periods while 4^th and 5^th are called long periods and 6^th periods is called longest periods and 7^th period is called incomplete period. Elements of period 1 and 2 are called typical elements.

 

Recently period 1 is called shortest period, 4^th and 5^th periods are called longs periods, 6^th period is termed as longest period and 7^th period is referred as incomplete period.

 

General Features of Periods

 

1. Constant number of shell and increasing number of valence electrons

2. Period number equals to the total number of shells

3. Each period begins with new energy level

4. Each period starts with an alkali metal and ends up with a noble gas

5. Increasing valency with respect to hydrogen from 1 to 4 and then falling from 3 to 1.

6. Different chemical properties due to different valence shell configuration

7. Variation in properties of elements change from metallic to non-metallic.

8. Placement of metals are at far left side and non-metals at the right side of table.

9. Increasing Metallic character and decreasing non-metallic character increases across each period.

10. Decreasing Atomic volume or atomic radius (i.e. size of atom) across each period.

11. The number of elements in each period is twice the number of atomic orbitals available

1.         In each period, number of shell remains the same but the number of valence electrons increases (from 1 to 8) across each period.

 

2.         Period number of an element represents the total number shells in that element e.g. Iron belongs to fourth period as it has 4 electronic shells in its atom. The period indicates the value of n for the outermost or valence shell.

 

The period of an element can be predicated by observing the electronic configuration of the element. The period of the element is same as the number of the valence shell. For example, if in an element third shell is the valence shell, then it belong the third period (period 3).

 

3.  Each period starts with the filling of electrons in a new energy level (quantum shell) and continues till the p-orbital of the same shell.

 

4.  Each period starts with an alkali metal (except 1^st period which begins with hydrogen) with one valence electron and ends up with a noble gas (except 7^th period) with 8 valence electrons except He which has only 2 electrons.

 

5. In each period (especially short periods), the valency of elements with respect to hydrogen increases from 1 to 4 and then falls from 3 to 1.

 

6.   All the elements in a period have different valence shell configuration and hence have different chemical properties. The elements within a period have dissimilar properties from left to right across any period.

 

7.  The physical and chemical properties of elements change from metallic to non-metallic along each period.

 

8.   Electropositive elements (metals) are at far left side while electronegative elements (non-metals, gases, metalloids) are at the right side of table.

 

9.  Metallic character decreases while non-metallic character increases from left to right across each period. e.g. Na is a metal while Cl is purely a typical non-metal.

 

10. Atomic volume or atomic radius (i.e. size of atom) decreases from left to right across each period.

 

11. The number of elements in each period is twice the number of atomic orbitals available in the energy level that is being filled. There are 2 elements in the 1^st period, 8 in the 2^nd, 8 in the 3^rd, 18 in the 4^th, 18 in the 5^th, 32 in the 6^th and 7^th period is incomplete.

 

Short and Long Periods

Period 2 and 3 are called short periods while period 4, 5 and 6 are called long periods. Elements of period 1 and 2 are called typical elements.

 

First Period (Shortest Period)

It is the shortest period of the periodic table containing only two elements hydrogen and helium (both of them are gaseous non-metals). This period corresponds to filling up of K-shell.

Second and Third Periods (Short Periods)

Period 2 and 3 are called Short Periods each contains 8 elements (2 s-block and 6 p-block elements) from ₃Li to ₁₀Ne and ₁₁Na to ₁₈Ar respectively.

 

Period 2 signifies the filling up of L-shell (containing two energy levels 2s and 2p) and period 3 corresponds to filling up of M-shell up to 8 electrons (containing two energy levels 3s and 3p).

 

Period 2 includes Li, Be, B, C, N, O, F & Ne while period 3 includes Na, Mg, Al, Si, P, S, Cl & Ar.

 

Their general valence shell configuration is ns¹ np⁰ to ns² np⁶.

 

Fourth and Fifth Periods (Long Periods)

Period 4 and 5 are called Long Periods each contains 18 elements (2 s-block, 6 p-block and 10 d-block elements) from ₁₉K to ₃₆Kr and ₃₇Rb to ₅₄Xe respectively.

 

Period 4 corresponds to filling up of N-shell and also M-shell (containing three energy level 4s, 3d, 4p i.e. period 4 starts with filling up of 4s orbital followed by 3d and finally 4p orbital) while period 5 signifies the filling up of O-shell and also N-shell (containing three energy levels 5s, 4d, and 5p i.e. period 5 starts with filling up of 5s orbital followed by 4d and finally 5p orbital).

 

Out of 18 elements, 8 elements (two s-block and eight p-block elements) are representative elements and the remaining ten are transition elements. Fourth period contains 10 transition elements (called 1^st transition series) from Scandium to Zinc i.e. ₂₁Sc to ₃₀Zn. Fifth period also contains 10 transition elements (called 2^nd transition series) from Yttrium to Cadmium i.e. ₃₉Y to ₄₈Cd.

 

Sixth Period (Longest Period)

Period 6 is the longest period and comprising of 32 elements from cesium to radon (i.e. ₅₅Cs to ₈₆Rn). This period corresponds to filling up of P-shell along with N-shell, and O-shell (containing four energy level 6s, 4f, 5d, and 6p i.e. period 6 starts with filling up of 6s orbital and after that electron should enter 4f, but after 6s one electron enters 5d (5d¹) and after that 4f orbital starts filling and completely filled followed by 5d and finally 6p orbital).

 

In 6^th period, there are 8 representative elements (two s-block and six p-block elements), of the remaining 24 elements, 10 are outer transition or d-block elements from lanthanum (₅₇La) and continues to  ₇₂Hf to ₈₀Hg and 14 are f-block or inner transition elements from ₅₈Cerium to ₇₁Lutetium (Lu).

 

For the sake of convenience, 14 f-block elements are placed at the bottom of periodic table. The 6^th period series of 14 inner transition or f-block elements that follows lanthanum (₅₇La) placed at the bottom of the periodic table is called Lanthanide Series or Rare Earth Elements and it is from cerium to lutetium (₅₈Ce to ₇₁Lu). In this series electrons are being added to the 4f sub-levels.

 

Seventh Period (Incomplete Period)

Period 7 is the second longest period and it is incomplete (as to date about 109 elements have been discovered) comprising of 26 elements (but expected to contain 32 elements) starts from francium (₈₉Fn). This period corresponds to filling up of Q-shell along with O-shell and P-shell containing three energy levels 7s, 5f, 6d].

 

It starts with the filling up of 7s orbital. Again after completing 7s, one electron enters 6d orbital (6d¹) then 5f starts filling up (5f¹ to 5f¹⁴) forming Actinide Series of Inner Transition Metals. After 5f¹⁴, electrons again occupying 6d.

 

It contains two representative elements (s-block), 10/8 outer transition or d-block elements and 14 inner transition or f-block elements placed at the bottom of the periodic table called Actinides. The 7^th period series of 14 inner transition or f-block elements that follows actinium (₈₉Ac) in which electrons are being added to the 5f sub-level placed at the bottom of the periodic table is called Actinide Series and it is from thorium to lawrencium (₉₀Th to ₁₀₃Lr). In this series electrons are being added to the 5f sub-levels.

 

 

 

 

The elements following Uranium (₉₂U) in 7^th period with atomic number greater than 92 (Z > 92) are known as trans-uranium elements. They are from Neptunium to Lawrencium (₉₃Np to ₁₀₃Lw). e.g. Am, Cf, Es. All these elements do not occur in nature and are artificially synthesized.

 

The elements following fermium (₁₀₀Fm) in 7^th period with atomic number greater than 100 (Z>100) are known as trans-fermium elements. Fermium is the 100^th element, so trans-fermium means beyond fermium.

 

 

Classification of Elements based on Electronic Configuration

 

The elements in the periodic table has been divided into 4 blocks or groups on the basis of electronic configuration whether the last electron called differentiating electrons enters into s, p, d or f-orbitals. The periodic table has been divided into following 4 blocks on the basis of electronic configuration:

1. Representative Elements (s-block and p-block elements)
2. Noble Gases.
3. Outer transition elements (d-block Elements)
4. Inner transition elements (f-block Elements)

General Valence Shell Electronic Configuration of different Blocks of the Periodic Table

Representative Elements ----  ns^1-2  np^1-5  (ns^1-2  np^1-6)

s-Block Elements -----     ns^1-2 

p-Block Elements ---   ns² np¹⁻⁶

d-Block Elements ----    (n–1)d^1-10,  ns^0-2 

f-Block Elements -------   

(n–2)f^2-14,  (n–1)d^0-1,  ns²  OR (n–2)f^1-14,  (n–1)d^0-1,  ns²  

(n–2)f^0-14,  (n–1)d^0-2,  ns²  OR (n–2)f^1-14,  (n–1)d^0-2,  ns²  

 

 

 

 1. Representative Elements

The elements of family A of the periodic table having incomplete outermost shell are called representative or normal or typical elements. They are so named because the properties of these elements are represented by valence electrons. They form group IA to VIIA and are present at left and right side of the periodic table. Their general valence shell electronic configuration of ns¹, ns² to ns² np⁵ (ns^1-2 to ns² np^1-5). They include s-block elements and p-block elements.

 

(a) s-block Elements

The elements in which outer electrons enter into s-orbital having ns^1-2 (i.e. ns¹ to ns²) valence shell configuration (where “n” denotes the number of outermost shell or the number of period ranges 2-7) are called s-block elements. Since last electron lies in ns orbital, they are referred to as s-block elements.  They form group IA and IIA of periodic table found at its far left side. There are total 13 s-block elements including hydrogen. They are highly reactive electropositive elements with low ionization potential showing fixed oxidation state of +1 and +2.  Except Li, Be, all form ionic compound. Their oxides are basic in nature.

 

(b) p-block Elements

The elements in which outer electron enters into p-orbitals having ns² np^1-6 (i.e. ns² np¹ to ns² np⁶ denotes the number of outermost shell or the number of period ranges 2-6) valence shell configuration are called p-block elements. Since in these elements outer electron enters into p-orbitals (which are being progressively filled), they are referred to as p-block elements. They from Group IIIA to Group VIIIA of the periodic table and are located at the extreme right of the periodic table. There are total 30 elements in six sub-groups of p-block including noble gases except helium. Mostly they are highly electronegative elements (non-metals). Their oxides are neutral or acidic. Mostly they form covalent compounds.

 

2. d-Block Elements or Outer Transition Elements

The elements having partially filled d-orbitals in their atoms or ions in which last electron enters into (n – 1)d-orbitals in their atomic state or ionized state (i.e. in their common oxidation states) are called d-block elements. Since last electron is in the process of occupying d-orbitals, they are known as d-block elements. In these elements, two outermost shells are incomplete i.e. in these elements, besides the outermost valence shell (s-subshell) penultimate shell (d-subshell) is also incomplete. These elements are called transition elements because they show transitional (intermediate) behaviour between the two sets of representative elements on either side of them i.e. s-block and p-block elements. They form B-family (group IB, IIB, IIIB to VIIIB) of the periodic table and hence are also called Group B Elements and are located in the middle of the periodic table between s-block and p-block elements. They are characterized by ns², (n – 1)d¹ to ns², (n – 1)d¹⁰  OR ns², (n – 1)d^1-10 valence shell electronic configuration (where n= 4-7). They are called outer transition elements as these elements are placed in the upper middle (outer) portion of the periodic table. They all are metals characterized by their variable valencies, forming coloured compounds and their ability to form complex ions by co-ordination through co-ordinate covalent bonds.

 

3.f-Block Elements/Inner Transition Elements                                                

The elements having partially filled f-orbitals (except ₇₁Lu, ₉₀Th, ₁₀₃Lr) in their atoms or ions in which last electron enters into (n – 2)f-orbitals in their atomic state or ionized state are called f-block elements. Since last electron is in the process of occupying f-orbitals, they are known as f-block elements. In these elements, three outermost shells are incomplete i.e. in these elements, besides the outermost valence shell (i.e. ns-subshell) and (n – 1)d-subshell, penultimate f-subshell or (n – 2)f-orbitals are also incomplete. Properly they should be placed after IIIB but these elements are found in a separate position at the bottom of the periodic table. These elements are also called the inner transition elements because the filling of electrons takes place in the inner (n – 2)f-orbitals (4f or 5f sub-shell) i.e. two levels below the outer (n – 1)p-orbitals (5p or 6p) and ns-orbitals (6s or 7s) orbitals which are already filled in these elements. They are characterized by (n – 2)f^0-14, (n – 1)d^0-2 , ns² OR (n – 2)f^2-14/0-14, (n – 1)d^0-1/0-2 , ns² valence shell electronic configuration (where n = 6-7).

 

1. Representative Elements

Definition

The elements of sub-group A or the elements of groups 1, 2, 13, 14,15,16,17, and 18 are called Main group or Normal or Major or Representative Elements as the properties of these elements are represented by valence electrons.

OR

The elements in which all their inner shells are complete but outermost shell is incomplete having less than 8 valence electrons in their outermost shell are known as representative elements i.e. s-block and p-block elements except inert gases are known as representative elements. In these elements the outermost shell gets progressively filled from group IA to group VIIIA as we move from left to right in a period.

 

Reason for Name

They are so named because the properties of these elements are represented by valence electrons.

Group and Position in table

They form group IA to VIIA and are present at left and right side of the periodic table.

General valence shell electronic configuration

Their general valence shell electronic configuration of ns¹, ns² to ns² np⁵ (ns^1-2 to ns² np^1-5).

Blocks Included

They include s-block elements and p-block elements.

General Characters

1.   They consist of some metals (IA and IIA Groups, some elements of IIIA group like Al, Ga, In, Tl and miscellaneous metals like Sn, Pb, Bi, Po), all non-metals and metalloids (B, Si, Ge, As, Sb, Se, Te).

 

2.   There are 42 representative elements including noble gases, of these 20 are metals (12 s-block metals and 8 p-block metals), 14 are non-metals (N₂, O₂, F_2, Cl₂, Br₂, I₂, C, P, S, Ne, Ar, Kr, Xe, Rn) and 8 are metalloids (B, Si, Ge, As, Sb, Se, Te, At). Out of 42 normal elements, 31 elements are solids (19 metals, 4 non-metals and 8 metalloids), 2 elements are liquids (Ga and Br₂) and 9 elements are gases.

 

3. Some elements are diamagnetic and some are paramagnetic. However, their compounds are generally diamagnetic and colourless.

Sub-Division of Representative Elements

Representative elements are of two types namely  s-block elements and p-block elements:

(a) s-block Elements

(b) p-block Elements

(a) s-block Elements

The electropositive elements of group IA and IIA of the periodic table in which outer electrons enter into s-orbital having ns^1-2 (i.e. ns¹ to ns²) valence shell configuration (where n = number of outermost shell or the number of period ranges 2-7) are referred as s-block elements. They form group IA and IIA of periodic table found at its far left side.

Since last electron lies in ns orbital (which is being progressively filled), they are referred to as s-block elements.

 

There are total 13 s-block elements including hydrogen (there should be 14 s-block elements as He is also associated with 1s² electronic configuration).

 

They are highly reactive electropositive elements with low ionization potential showing fixed oxidation state of +1 and +2 forming ionic compound except Li, and Be. Their oxides are basic in nature.

 

Groups Included

They include group IA and IIA. Group IA includes 6 elements namely Li, Na, K, Rb, Cs, Fr. Francium is radioactive. The IA group is also called Lithium Family being Lithium is the first member. Group IA elements are called Alkali Metals because they yield strong alkalis when their oxides, hydroxides or hydrides are dissolved in water which are completely soluble in water.

Group IIA includes 6 elements namely Be, Mg, Ca, Sr, Ba and Ra. Radium is radioactive. The IIA group is also called Beryllium Family being Be is the first member. Group IIA elements are called Alkaline Earth Metals because they and their compounds (CaCO₃) are found abundance in earth crust and yield weak bases when their oxides, hydroxides and hydrides are dissolved in water which are sparingly soluble in water.

Electronic Configuration of s-Block Element

Alkali Metals

Since, their valence shell configuration is ns¹, so their principal oxidation state is +1 and they form monovalent cation.

Alkaline Earth Metals

Since, their valence shell configuration is ns², so their principal oxidation state is +2 and they form divalent cation.

general characters

They have following general characters:

1. Highly electropositive nature

2. Monovalent and fixed oxidation state of +1 and +2

3. Forming Ionic Bond

4. Basic nature of oxides, hydroxides and hydrides

5. Magnetic behaviour

6. Forming Colourless compounds

7. Reducing behaviour

 

1.    They are highly reactive electropositive elements with low ionization potentials.

 

2.     Group IA elements are monovalent exhibiting fixed oxidation states of +1 while group IIA elements are divalent showing fixed oxidation state of +2.

 

4.    Except Li, Be, all form ionic bonds (compound).

 

3.   Their oxides and hydroxides are basic in nature. Group IA elements are called alkali metals as their oxides yield strong alkalis on dissolution. Group IIA elements are called alkaline earth metals as these metals exist as their oxides (magnesia; MaO, lime; CaO, strontia; SrO, baryta; BaO) in the earth which give weak alkaline solution (except BaO) on dissolution.

5.   Group IA elements are paramagnetic and group IIA elements are diamagnetic. However their compounds are mostly diamagnetic.

6.  They form colourless compounds except their manganates, permanganates, chromates, dichromates etc. e.g. KMnO₄, NaMnO₄, Na₂CrO₄, K₂CrO₄, K₂MnO₄ etc.

 

7.    They are powerful reducing agent due to their largest assize, lowest I.P and highest negative reduction potential.

 

8. All s-block elements (except Be and Mg) impart characteristic colours to the non-luminous Bunsen flame when ignited due to their low ionization enthalpies. The formation of coloured flame is explained by the fact that these metals or their salts on heating in a flame, the valence electrons get excited and jumps to higher energy level, when these excited electrons return back to the original ground state, the absorbed energy is liberated as different frequency of visible light in the visible region of the electromagnetic spectrum and hence the flame appears coloured. The different colours arise due to different energies required for electronic excitation and de-excitation. The colour imparted by an element depends upon the I.P. Higher the I.P, higher will be the frequency of radiation absorbed and consequently lower will be frequency of radiation emitted i.e. the colour imparted. Since Li has higher I.P, it emits radiation of lower frequency i.e. red colour.

Be and Mg atoms due to their small size and high I.P, their electrons more strongly bound (because of higher effective nuclear charge) and are not excited by the energy of the flame to higher energy states. Hence, they require high excitation energy and therefore, these elements do not give any colour in Bunsen flame.

(b) p-block Elements

The elements in which outer electron enters into p-orbitals having ns² np^1-6 (i.e. ns² np¹ to ns² np⁶ where n denotes the number of outermost shell or the number of period ranges 2-6) valence shell configuration are called p-block elements. Since in these elements outer electron enters into p-orbitals (which are being progressively filled), they are referred to as p-block elements.

 

They from Group IIIA to Group VIIIA of the periodic table and are located at the extreme right of the periodic table. There are total 30 elements in six sub-groups of p-block including noble gases except helium. Among these elements, 2 are liquids (Ga and Br), 9 are gases (N₂, O₂, F₂, Cl₂, Ne, Ar, Kr, Xe, Rn) and 19 are solids. Out of 30 elements, 8 are metals (Al, Ga, In, Tl, Sn, Pb, Bi, Po), 8 are metalloids (B, Si, Ge, As, Sb, Se, Te,At) and 14 are non-metals including first member of group IVA (C), first two members of group VA (N, P), first two members of group VIA (O, S), first four elements of group VIIA (F, Cl, Br, I) group.

 

they are mostly highly electronegative elements (non-metals) mostly forming covalent compounds and their oxides are usually either acidic or neutral.

general characters

They have following general characters:

1. Highly electronegative nature

2. Polyvalent nature and variable oxidation state of

3. Forming covalent bond

4. Variable nature of oxides

5. Magnetic behaviour

6. Forming Colourless compounds

7. Oxidizing behaviour

1.  They are mostly highly electronegative elements with high ionization potentials. They also include less active non-metals, gases, metalloids and some weak non-active metals.

2.   They usually exhibit variable (more than one) or variety of oxidation states ranging from – ½  +1 to +7.

3.  Mostly they form covalent compounds.

4. Their oxides are mostly acidic in nature (e.g. B₂O₃, CO₂, SiO₂, NO₂, P₂O₃, SO₂, SO₃, Cl₂O, Cl₂O₇). But some may be neutral (CO, NO, N₂O, H₂O etc) or amphoteric (Al₂O₃, As₂O₃ etc.) or basic             (Bi₂O₃).

5.  They may be paramagnetic or diamagnetic. However their compounds are mostly diamagnetic.

6.  They mostly form colourless compounds except NO₂.

7. They are mostly oxidizing agent like F₂, Cl₂, O₂ etc. except Al, C etc.

2. The Noble Gases/Aerogens

3. d-Block Elements or Outer Transition Elements

Definition

The elements of sub-group B or the elements of groups 3 t 12 having partially filled d-orbitals in their atoms or ions in which last electron enters into (n–1)d-orbitals in their atomic state or ionized state (i.e. in their common oxidation states) are called d-block elements or outer transition elements.

 

In these elements, two outermost shells are incomplete i.e. besides the outermost valence shell (s-subshell) penultimate shell (d-subshell) is also incomplete.(These are the element whose outermost s-sub-shell and penultimate d-sub-shells are being occupied with electrons).

 

They all are metals characterized by their high melting and boiling points, variable valencies, catalytic property, forming coloured paramagnetic compounds and their ability to form complex ions by co-ordination through co-ordinate covalent bonds.

 

Reason for calling d-Block elements

Since last electron is in the process of occupying d-orbitals, they are known as d-block elements.

 

Reason for calling Transition elements

These elements are called transition elements because they show transitional (intermediate) behaviour between the two sets of representative (s and p-block) elements on either side of them.

 

Reason for calling outer Transition elements

They are called outer transition elements as these elements are placed in the upper middle (outer) portion of the periodic table.

 

Group Formed

They form B-family (group IB, IIB, IIIB to VIIIB) of the periodic table and hence are also called Group B Elements and are located in the middle of the periodic table between s-block and p-block elements.

 

General valence shell electronic configuration

They are characterized by (n–1)d¹, ns^0-2 to (n–1)d¹⁰, ns^0-2 OR (n–1)d^1-10, ns^0-2  valence shell electronic configuration (where n= 4-7).

 

Sub-Division

The outer transition (d-block) elements consist of following 4 series of 10 elements each:

Characteristics of Transition Elements

1.         Metallic Character

2.         Variable nature of Bonding

3.         Variable oxidation states

4.         Variable nature of oxides

5.         Coloured Ions and compounds

6.         Magnetic behaviour

7.         Catalytic Properties

8.         High melting and boiling points

9.         High density and hardness

10.       Complex formation

 

1.         They all are metals are mostly less reactive characterized by partially filled d-orbitals in their atoms or ions in ground state or in any oxidation state (except Zn, Cd, Hg).  

 

2.  They form ionic, covalent and co-ordination (complex) compounds as they can form all the three types of bonds i.e. ionic, covalent, and co-ordinate bonds.

 

3.  Excepting Zn, Cd, Ag, Sc, Y, they show variable valency forming more than one type of ions (e.g. Cu¹⁺, Cu²⁺, Fe²⁺, Fe³⁺ etc.). Excepting Zn, Cd, Ag, Sc, Y, they show more than one oxidation states in their compounds. Cu shows +1 and +2.

 

4.   They form acidic (CrO₃, Mn₂O₇ etc), basic () and amphoteric (ZnO, Cr₂O₃) oxides. Acidic nature of oxides increases with increasing oxidation state.

 

5.  They themselves, their ions and their compounds are highly coloured except Zn, Cd, Hg e.g. ferric salts are brown.

 

6. except Zn, Cd, Hg, they all are paramagnetic owing to presence of unpaired electrons.

 

7.  They exhibit catalytic property and behave as catalyst in a variety of reactions.

8.  Their melting points, boiling points and densities are usually very high except  Sc, Y and La.

9.  They form complex ions.

 

Detailed Characteristics of d-Block Elements

1.   Metallic Character

They all are metals with pronounced electropositive character. However, most of them are relatively less reactive metals. They are hard, malleable, ductile, good conductor of electricity and heat and form alloys.

 

2.   Variable oxidation states

With the exceptions of Zn, Cd, Sc, and Ag, they exhibit variety of variable oxidation states from 0 to +8 in their compounds e.g. Cr⁺³ or Cr⁺⁶, Mn⁺², Mn⁺³, Mn⁺⁴ etc.

 

3.   Variable nature of Bonding

They form covalent, ionic as well as co-ordination compounds or complexes. They are capable to form co-ordination compounds or complexes due to the presence of small, highly charged cation with highly effective nuclear charge and availability of vacant d-orbitals (and tendency of central transition metal to acquire effective magic numbers of next inert gas) which can accept lone pair of electrons donated by other groups called ligands.

 

4.   Variable nature of oxides

They form acidic (CrO₃, Mn₂O₇ etc), basic () and amphoteric (ZnO, Cr₂O₃) oxides. Acidic nature of oxides increases with increasing oxidation state.

                                    2CrO_3(s)  +  H₂O  ----→ 2H⁺  +  Cr₂O₇^2-

 

5.   High melting and boiling points

Their melting and boiling points are generally very high except Zn, Cd and Hg due to strong interatomic binding forces (in the form of strong covalent bonding) between their atoms extending throughout the crystals owing to small atomic size and unfilled d-orbitals containing unpaired valence electrons for interaction.

 

6.   High density and hardness

They are hard having high densities due to their small atomic volumes (except Sc and Y).

 

7.   Magnetic behaviour

They all are paramagnetic due to presence of unpaired electrons with the exceptions of Zn, Cd and Hg which are diamagnetic. Fe and Co are strongly paramagnetic and can be magnetized and are called ferromagnetic.

 

8.   Coloured Ions and compounds

They form variety of highly coloured paramagnetic compounds. Colour is associated with incompletely filled electron shells and ability for electronic transitions from one energy level to another. Sc³⁺, Y³⁺, Zn²⁺, Cd²⁺, Hg²⁺ ions are non-transitional since they either do not have 3d-electrons or ten 3d-electrons (i.e. completely filled  (n-1)d-orbitals)

 

9.   Catalytic Properties

A number of transition metals and their compounds act as catalytic agent as they have ability to provide low energy path for the reaction either by forming intermediate compounds or by the change of oxidation states e.g. Cu, Fe, Ni, Pt, Pd, W, V₂O₅, Cr₂O₃, ZnO, FeO etc.

 

10. Low reactivity

They show an increasing tendency to remain unreactive and this tendency is most pronounced in gold and platinum. This is due to their high I.P. They are relatively more reactive in powdered form.

 

 4. Inner Transition Elements                                                                             

Definition

The elements having partially filled f-orbitals (except ₇₁Lu, ₉₀Th, ₁₀₃Lr) in their atoms or ions in which last electron enters into (n–2)f-orbitals in their atomic state or ionized state (i.e. in their common oxidation states) are called f-block elements or inner transition elements.

In these elements, three outermost shells are incomplete i.e. in these elements, besides the outermost valence shell (i.e. ns-subshell) and (n–1)d-subshell, penultimate f-subshell or (n–2)f-orbitals are also incomplete.

 

Reason for calling f-block elements

Since in these elements, last electron is in the process of occupying penultimate f-orbitals, they are known as f-block elements.

 

Group Formed

Properly they should be placed after IIIB but these elements are found in a separate position at the bottom of the periodic table.

 

Reason for calling Inner Transition elements

These elements are also called the inner transition elements because the filling of electrons takes place in the inner (n–2)f-orbitals (4f or 5f sub-shell) i.e. two levels below the outer (n–1)p-orbitals (5p or 6p) and ns-orbitals (6s or 7s) orbitals which are already filled in these elements.

 

General valence shell electronic configuration

They have following valence shell electronic configuration

(n–2)f^2-14, (n–1)d^0-1, ns²  and (n–2)f^0-14, (n–1)d^0-2, ns²

OR

(n–2)f^2-14/0-14, (n–1)d^0-1/0-2, ns²  (where n = 6-7).

Sub-Division

f-block consists of two series of 14 elements each namely Lanthanide Series and Actinide Series. Lanthanides and actinides are collectively known as inner transition elements.

The 6^th period series of 14 inner transition elements or f-block elements that follows Lanthanum (₅₇La) in which electrons are being added to the 4f-orbitals placed at the bottom of the periodic table is called Lanthanide Series or Lanthanides or Rare earth elements and it is from cerium to lutetium (₅₈Ce to ₇₁Lu). 

 

The 7^th period series of 14 inner transition elements or f-block elements that follows Actinum (₈₉Ac) in which electrons are being added to the 5f-orbitals placed at the bottom of the periodic table (just beneath lanthanides series) is called Actinides Series or Actinides and it is from Thorium to Lawrencium (₉₀Th to ₁₀₃Lr).

 

General Characters of f-block elements

1.         They are metals generally paramagnetic and form coloured ions.

2.         Their compounds are also paramagnetic.

3.  The last natural element in this series is Uranium (Z=92). Elements following uranium are called transuranic or transuranium elements which are artificially prepared by nuclear reactions.