Introduction:
Upto the end of seventeenth century, only 31 elements were known. Therefore, it was very easy to study and remember the properties of these elements. However, during the later part of the eighteenth century, the pace of discovery of new elements quickened. Between 1800 to 1869, the number of identified elements had become nearly double to 63. With such a large number of elements, it became difficult to study individually the chemistry of all these elements and their innumerable compounds. At this stage, it was realised that there should be some simple way to study and remember the numerous properties of the elements and their compounds. This gave rise to necessity of classification of the elements into various groups having similar properties. This has been done by arranging the elements in such a way that similar elements are placed together while dissimilar elements are separated from one another. This is known as classification of elements. Such a classification of the elements has resulted in the formulation of the periodic table. Periodic table may be defined as the arrangement of the known elements according to their properties in a tabular form.
At present, about 118 elements are known. Of these, the recently discovered elements are man made. Efforts to synthesise new elements are continuing. The periodic classification of the elements has extremely simplified their study. Not only the periodic classification rationalizes the known chemical facts about elements, but it also helps to predict new ones for undertaking further study
Historical Development of the Periodic Table:
Since the beginning of the nineteenth century, scientists have been trying to find a basis of grouping elements having similar properties. Lavoisier classified the elements simply as metals and nonmetals. However, this classification proved to be inadequate. Some of the earlier important attempts to classify the elements are briefly summed up below:
Dobereiner’s triads:
In 1817, a German scientist, John Dobereiner classified the elements in groups of three elements called triads. The elements in a triad had similar properties and the atomic weight of the middle member of each triad is very close to the arithmetic mean (average) of the other two elements. The common triads of Dobereiner classification were: chlorine, bromine and iodine; calcium, strontium and barium; lithium, sodium and potassium, etc.
Dobereiner’s triads
Li Lithium Ca Calcium Cl Chlorine
Na Sodium Sr Strontium Br Bromine
K Potassium Ba Barium I Iodine
The Dobereiner’s relationship was also referred to as law of triads. However, it seemed to work only for a few elements. It was dismissed as coincidence because all the known elements could not be arranged in triads.
De Chancourtois Classification:
The next reported attempt was made by a French geologist, de Chancourtois in 1862. He arrange the then known elements in order of increasing atomic weights and proposed a cylindrical table of elements display the periodic recurrence of properties. He observed that the elements with similar properties in a vertical line from the centre of the spiral. However, this did not attract much attention.
Newlands law of octaves:
In 1865, an English chemist, John Newlands proposed a new system of grouping elements of similar properties. According to him, when the elements are arranged, in the increasing order of atomic weight, the properties of every eighth element are similar to the first one. Newlands called this relation as the law of octaves due to similarity with the musical scale.
This relationship is just like every eighth note resembles the first in octaves of music. The Newlands law of octaves seemed to be true only for elements up to calcium.
In the above table, sodium (Na), the eighth element from lithium (Li) is similar to Li and similarly next eighth element potassium is similar to Na. The same is true for magnesium resembling beryllium, aluminium resembling boron, etc. In the beginning, the idea of Newlands’s was not widely accepted. However, his work was later recognized by the Royal Society of London and he was awarded Davy Medal in 1887.
The main credit of the periodic law, as we know it today goes to the Russian Chemist Dmitri Mendeleev and the German chemist Lother Meyer who proposed the same law while working independently in 1869.
Mendeleev’s Periodic Law:-
In 1869 Mendeleev made a remarkable contribution to the classification of elements. On the basis of physical and chemical properties of the elements, he gave a law known as the periodic law. The law states that “the physical and chemical properties of elements are periodic function of their atomic mass”.
This means that when the elements are arranged in order of their increasing atomic mass, the elements with similar properties recur at regular intervals. Such orderly recurring properties in a cyclic fashion are said to be occurring periodically. This is responsible for the name periodic law or periodic table.
Mendeleev’s Periodic Table:
On the basis of his periodic law, Mendeleev arranged all the known elements in the form of a table known as periodic table. It was observed that the elements with similar properties recur at regular intervals or periodically. As a result of this, the elements fall in certain groups or families. The elements in each group were similar to each other in many properties. The properties repeated periodically. The horizontal rows in the periodic table are called periods and the vertical columns are called groups. The original periodic table by Mendeleev had six periods and eight groups. In his table, each group is further subdivided into two sub groups marked A and B groups.
Mendeleev’s system of classifying elements was more elaborate. He fully recognized the significance of periodicity and used broader range of physical and chemical properties to classify the elements. Mendeleev in particular mainly relied on the similarities in the empirical formulas and properties of the compounds formed by the elements. He went even further than that. He realized that some of the elements did not fit in very well with his scheme of classification if the order of atomic mass was strictly followed. He showed courage to ignore the order of atomic mass, thinking that the atomic mass measurements might be incorrect. He placed the elements with similar properties together. For example, iodine has lower atomic mass than that of tellurium (of Group VI) but he placed iodine in Group VII along with fluorine, chlorine and bromine because of similarities in their properties. At the same time, keeping his primary aim of arranging elements of similar properties in the same group he noticed that he had to skip several places.
The Mendeleev’s classification gave him so much confidence that he boldly left certain spaces or gaps for undiscovered elements. By considering the properties of the adjacent elements in his table, he predicted the properties of the undiscovered elements. Later on, when these elements were discovered, their properties were found to be exactly similar to those predicted by Mendeleev. For example, gallium and germanium were not discovered at that time, when Mendeleev formulated his periodic table and therefore, he left gaps for these elements. He not only predicted the existence of the elements but he estimated their properties. He tentatively named these elements as eka-aluminium and eka-silicon (word eka meaning ‘next’) because he believed that these would be similar to aluminium and silicon respectively. When chemists discovered these elements, Mendeleev’s prediction of their properties proved to be remarkably correct.
Important Contributions of Mendeleev’s Periodic Table:
Mendeleev’s periodic table was one of the greatest achievements in the development of chemistry. Some of the important contributions of his periodic table are:
- Systematic study of the elements. The Mendeleev’s periodic table simplified the study of chemistry of elements. Knowing the properties of one element in a group, the properties of other elements in the group can be easily guessed. Thus, it became very useful in studying and remembering the properties of a large number of elements.
- Correction of atomic masses. The Mendeleev’s periodic table helped in correcting the atomic masses of some elements based on their positions in the table. For example, atomic mass of beryllium was corrected from 13.5 to 9. Similarly, with the help of this table, atomic masses of indium, gold, platinum etc. were corrected.
- Prediction of new elements. At the time of Mendeleev, only 56 elements were known. While arranging these elements, he left some gaps. These gaps represented the undiscovered elements. Mendeleev predicted the properties of these undiscovered elements on the basis of their positions. For example, he predicted the properties of scandium, gallium and germanium which were discovered later. The observed properties of these elements were found to be similar to those predicted by Mendeleev (Table 3.3).
Defects of Mendeleev’s Periodic Table:
In spite of many advantages, the Mendeleev’s periodic table had certain defects. Some of these are—
- Position of hydrogen– Hydrogen is placed in group I. However, it resembles the elements of group I (alkali metals) as well as the elements of group VIIA (halogens). Therefore, the position of hydrogen in the periodic table is not correctly defined.
- Anomalous pairs– In certain pairs of elements, the increasing order of atomic masses was not obeyed. In these cases, Mendeleev placed elements according to similarities in their properties and not in increasing order of their atomic masses. For example, argon (Ar, atomic mass 39.9) is placed before potassium (K, atomic mass 39.1). Similarly, cobalt (Co, atomic mass 58.9) is placed before nickel (Ni, atomic mass 58.6) and tellurium (Te, atomic mass 127.6) is placed before iodine (I, atomic mass 126.9). Their positions were not justified.
- Position of isotopes– Isotopes are the atoms of the same element having different atomic masses but same atomic number. Therefore, according to Mendeleev’s classification, these should be placed at different places depending upon their atomic masses. For example, isotopes of hydrogen with atomic masses 1, 2 and 3 should be placed at three places. However, isotopes have not been given separate places in the periodic table.
- Some similar elements are separated and dissimilar elements are grouped together– In the Mendeleev’s periodic table, some similar elements were placed in different groups while some dissimilar elements had been grouped together. For example, copper and mercury resembled in their properties but they had been placed in different groups. At the same time, elements of group I A such as Li, Na and K were grouped with copper (Cu), silver (Ag) and gold (Au), though their properties are quite different.
- Cause of periodicity– Mendeleev did not explain the cause of periodicity among the elements.
- Position of lanthanides and actinides– The fourteen elements following lanthanum (known as lanthanides, from atomic number 58—71) and the fourteen elements following actinium (known as actinides, from atomic number 90—103) have not been given separate places in Mendeleev’s table.
Modern Periodic Law:-
Moseley to conclude that atomic number is the fundamental property of the atoms than atomic weight. He, therefore, suggested that atomic number instead of atomic weight should be the basis of classification of elements.
The acceptance of atomic number, as the important characteristic of an atom, led to the modern periodic law. Modern periodic law may be stated as—
“The physical and chemical properties of the elements are periodic functions of their atomic numbers”.
Consequently, when the elements are arranged in the order of their increasing atomic numbers, it is observed that the elements of similar properties recur at regular intervals. As a result of this, the elements fall in certain groups and lead to an arrangement called the periodic table.
General Feature of Long Form (Modern) Periodic Table:
The long form periodic table is based on the atomic number i.e. all the elements have arranged in the order of increasing atomic number, the arrangement consists of the following features:
1) There are seven horizontal rows, called periods.
2) There are eighteen vertical columns, called groups.
3) The elements as arranged in the long form periodic table (modern) can be divided into four different blocks known as s, p, d and f blocks respectively. This division of elements is based upon the electronic configurations of the atoms. In this division the element which involve the filling of a particular orbital (i.e., s, p, d or f) are grouped together.
S-block elements:
The elements in which the last electron enters the -orbital of their outermost energy level are called s-block elements. It consists of elements of groups 1 and 2 having the ground state electronic configurations of outermost shell as ns1 and ns2 respectively. The elements corresponding to ns1 configuration are called alkali metals while those corresponding to ns2 configuration are called alkaline earth metals. Thus, the general electronic configuration of s-block elements may be expressed as: ns1-2
P-block elements:
The elements in which the last electron enters the p-orbital of their outermost energy level are called p-block elements. The elements of groups 13 to 18 involving addition of one (ns2 np1), two (ns2 np2), three (ns2 np3), four (ns2 np4), five (ns2 np5 and six (ns2 np6) electrons respectively in p-orbitals and s-orbitals are already filled in their atoms constitute p-block. The general electronic configuration of p-block may be written as: ns1-2 np1-6
d-block elements:
The elements in which the last electron enters the d-orbital of their outermost energy level are called d-block elements. These elements are also called transition elements. The general electronic configuration of d-block may be written as: (n-1) d1 – 10 ns1 or 2
f-block elements:
The elements in which the last electron enters the f-orbital of their outermost energy level are called f-block elements. These elements are also called inner-transition (lanthanides and actinides) elements. The general electronic configuration of f-block may be written as:
(n-2) f 1-14 (n-1) d 0 – 1 ns 2
Advantages of the Long Form of the Periodic Table:
The important advantages of the long form of the periodic table are given below:
i) This classification is based on the atomic number which is a more fundamental property of the elements.
ii) Since this classification is based on the atomic number and not on the atomic mass, the position of placing isotopes at one place is fully justified.
iii) The position of elements in the periodic table is governed by the electronic configurations, which determine their properties.
iv) It is easy to remember and reproduce.
v) The systematic grouping of elements into four blocks s, p, d and f has made the study of the elements simpler.
vi) The position of some elements which were misfit on the basis of atomic mass is now justified on the basis of atomic number. For example, argon proceeds potassium because argon has atomic number 18 and potassium has 19.
vii) The lanthanides and actinides which have properties different from other groups are placed separately at the bottom of the periodic table.
Defects of the Long Form of the Periodic Table:
Although the Long Form of the periodic table has helped in systemising the study of chemistry of elements, yet it has certain defects. The main defects of this table are:
i) The position of hydrogen is not settled. It resembles with alkali metals as well as halogens. However, it has been placed with alkali metals.
ii) Lanthanides and actinides have not been accommodated in the main body of the periodic table.
Periodic trends in properties of elements:
There are numerous physical properties of elements such as melting points, boiling points, enthalpy of fusion, enthalpy of vaporization, enthalpy of atomization, density, etc. which show periodic variations. These are indirectly related to electronic configurations of atoms. However, some physical properties such as valency, atomic size, ionization enthalpy, electron gain enthalpy, electronegativity, etc. are directly related to the electronic configuration of atoms. Some of the important properties and their periodic trends are:
1) Atomic Radii:-
The size of atom is very important property because many physical and chemical properties are related to it. If the atom is assumed to be spherical, the atomic size is given by the radius of the sphere and is called atomic radius. Generally, the term atomic radius means “the distance from the centre of the nucleus to the outermost shell of electrons”. However, it is difficult to determine the exact radius of the atom because of the following reasons
i) The size of an atom is (approximately 1.2 A or 1.2 X 10-10 m in radius) very small.
ii) According to the probability picture of electrons, an atom does not have well defined boundary. The probability of finding the electron is never zero even at large distances from the nucleus.
iii) It is not possible to isolate an atom and measure its radius. The probability distribution of an atom is also affected by the presence of other atoms in its neighbourhood. Therefore, the size of atom may change in going from one set of environment of another.
iv) Size of atom also changes from one bonded state to another.
Thus, we can only arbitrarily define atomic radius as the effective size which means the distance of closest approach of one atom to another atom in a given bonding situation. The approximate radii of atoms can be determined by measuring the distance between the centres of two neighbouring atoms (called internuclear distance) in a covalent molecule by X-ray diffraction, electron diffraction or other spectroscopic techniques.
As shown in Fig., this internuclear distance corresponds to twice the radius of an atom and, therefore, half of this distance gives the atomic radius.
This radius is also referred to as covalent radius and may be defined as one half of the distance between the centres of nuclei of two similar atoms bonded by a single covalent bond.
For homonuclear molecule,
r covalent = Internuclear distance between/ two bonded atoms
For example, as shown in Fig., the internuclear distance between two hydrogen atoms in H2 molecule is 74 pm.
Atomic radius of hydrogen = 74/2 = 37 pm.
Similarly, the atomic radii of chlorine and bromine are 99 pm and 114 pm because the internuclear distances in chlorine Cl—Cl and Br—Br are 198 pm and 228 pm respectively.
In the case of molecules containing different atoms (heteronuclear molecules) the covalent radius of an atom may be defined as: “the distance between the centre of nucleus of the atom and the mean position of the shared pair of electrons between the bonded atoms”.
In general, the atomic radii decrease with increase in atomic number (going from left to right) in a period. For example, in the second period, the atomic radii decrease from Li to Ne through Be, C, N, O and F. This may be explained on the basis of increasing nuclear charge along a period. With the increase in atomic number from lithium to fluorine, the magnitude of the nuclear charge increases progressively by one unit while the corresponding addition of electron takes place in the same principal shell (second). Since, the electrons in the same shell do not screen each other from the nucleus, the increase in nuclear charge is not neutralised by the extra valence electron. As a result, electrons are pulled closer to the nucleus by the increased effective nuclear charge and thereby, cause a decrease in the size of the atom in this way, the atomic size goes on decreasing across the period.
2) Inert Gas Radii:-
The atomic radius abruptly increases in case of noble gas element, Ne. This is because of the reason that the values for other elements are for covalent radii while the value of Ne is not covalent radius because neon can not form neon (Ne2) molecule. In case of noble gas elements, we measure van der Waals radius because these atoms are held together by weak van der Waals forces in solid state. This is also called inert gas radius. Therefore, inert gas radius may also be defined as: “one half of the internuclear distance between two adjacent atoms of the substance in the solid state”.
The van der Waals forces existing between atoms in the solid state are weak and the atoms are held at larger distances. Therefore, the internuclear distances in case of atoms held by van der Waals forces are larger than that between covalently bonded atoms. Consequently, van der Waals radii are always larger than covalent radii. Therefore, inert gas radii of noble gases are large.
The atomic radii of elements increase from top to bottom in a group. In moving down a group, the nuclear charge is increasing with increase in atomic number. Consequently, the distance of the outermost electron from the nucleus gradually increases down a group.
3) Ionic Radii:-
The ionic radii correspond to the radii of ions in ionic crystals. The ions are formed as a result of addition or removal of electrons from the outermost shells of atoms. The ions formed by the loss of electron acquire positive charge and are called cations while the ions formed by gain of electrons, get negative charge and are called anions. Ionic radius may be defined as the effective distance from the nucleus of the ion upto which it has an influence in the ionic bond.
The ionic radii in a particular group increase in moving from top to bottom because the number of shells increases.
Note:-
i) The radius of positive ion (cation) is always smaller than that of the parent atom.
ii) The radius of negative ion (anion) is always larger than that of the parent atom.
4) Ionisation Enthalpy:-
The electrons in an atom are attracted by the positively charged nucleus. In order to remove an electron from an atom, energy has to be supplied. The quantitative measure of the tendency of an atom to lose an electron is given by ionization enthalpy. It is defined as: “the energy required to remove an electron from isolated gaseous atom in its ground state”.
The first ionization enthalpy, IE1 is the energy required to remove the most loosely electron of the neutral atom and the second ionization enthalpy, IE2 the energy required to remove the second electron from the resulting cation and so on. Thus, first ionization enthalpy of an element (M) may be defined as the enthalpy change (∆H) for the reaction represented as:
M (g) ——–> M+ (g) + e– (g)
In other words, first ionization enthalpy is the enthalpy change when most loosely bound electron is removed from an isolated gaseous atom.
The ionization enthalpy is expressed in units of kJ mol-1. Similarly, we can define second ionization enthalpy as the energy required or enthalpy change to remove the second most loosely bound electron. In other words, it is the enthalpy change for the reaction:
M+ (g) ——-> M+2 (g) + e– (g)
Since energy is required to remove electrons from an atom and therefore, ionization enthalpies are always positive. Thus, the ionisation enthalpy gives the ease with which electron can be removed from an atom.
Note: – Evidently, the smaller the value of ionisation enthalpy, the easier it is to remove the electron from the atom.
Ionisation enthalpy depends on the following factors—
i) Size of the atom: – With increase atomic size it becomes easier to remove the electron and therefore, ionisation enthalpy decreases.
ii) Charge on the nucleus: – With the increase in nuclear charge, it becomes more difficult to remove an electron and hence ionisation enthalpy increases.
iii) Screening effect of the inner electrons: – An increase in the number of inner electrons tends to decrease the ionisation enthalpy.
iv) Electronic configuration: – The ionisation enthalpy depends upon the electronic configuration of the atom. Thus, more stable electronic configuration, the greater is the ionisation enthalpy.
Ionisation enthalpy in general, increases with increasing atomic number in a periods and decreases in moving from top to bottom in a group.
[Exceptional case: B>Be & O>N]
5) Electron Gain Enthalpy:
Energy is released when an electron is added to a neutral gaseous atom. This is called electron gain enthalpy or electron affinity. Thus, electron gain enthalpy is the enthalpy change or the amount of energy released when electron is added to an isolated gaseous atom.
X (g) + e– ———- X– (g) + E.A.
Electron gain enthalpy or electron affinity depends upon the following factors:
There are many factors which govern the electron gain enthalpy but the following are some important factors on which it mostly depends
i) Nuclear charge. The electron gain enthalpy become more negative as the nuclear charge increases. This is due to greater attraction for the incoming electron if nuclear charge is high.
ii) Size of the atom. With the increase in size of the atom, the distance between the nucleus and the incoming electron increases and this results in lesser attraction. Consequently, the electron gain enthalpy become less negative with increase in size of the atom of the element.
iii) Electronic configuration. The elements having stable electronic configurations of half filled and completely filled valence subshells show very small tendency to accept additional electron and thus electron gain enthalpies are less negative.
In general, electron gain enthalpy increases from left to right in a period and becomes less as we go from top to bottom in a group.
Note:-
i) Halogen have the highest electron gain enthalpy. This is due to the fact that halogens have general electronic configuration of ns2np5, one electron is less than the stable electronic configuration (ns2np6).
ii) Electron gain enthalpy values of noble gases are positive while those of Be, Mg, N and P are almost zero. The electron gain enthalpy is positive because they have stable electronic configuration and thus they have absolutely no tendency to accept an additional electron. Similarly, the low electron gain enthalpy or electronic affinity values of Be, Mg, N and P can be explained due to the extra stability of completely filled 2s and 3s orbitals in Be (2s2) and Mg (3s2) and half filled 2p and 3p orbitals in N and P respectively. Therefore, the configuration show little tendency to any electron and hence very low electron gain enthalpy or electronic affinity.
iii) Electron gain enthalpy or electronic affinity of fluorine is unexpectedly less than that of chlorine. This is due to the small size of F-atom.
6) Electronegativity:
Electronegativity is defined as: “the tendency of an atom of an element to attract the shared pair of electrons towards itself in a covalent bond”.
Greater the electronegativity of an atom, greater will be its tendency to attract the shared pair of electrons towards itself. Fluorine atom has the greatest power of attracting electrons and is the most electronegative element. It must be remembered that unlike other atomic properties such as ionisation enthalpy, electron gain enthalpy which are related to individual gaseous atoms, the electronegativity is related to atoms in the bonded state.
In general, the electronegativity increases on moving across a period from left to right and decreases on moving down the group.
Electronegativity of an element can be measures by the following methods:
i) Mulliken Scale: – According to this scale, electronegativity of an atom is taken as average of ionisation energy and electron affinity as-
Electronegativity = (IE + EA)/2.
ii) Pauling Scale: – It is based on bond energy data. If we take energy in eV per mole, then the electronegativity of two atoms A and B may be expressed as:
XA – XB = 0.208(∆E) ½.
Where ∆E is the difference between actual bond energy of AB and mean bond energy
Mulliken and Pauling Scale are related as: χ (Pauling) = χ (Mulliken)/2.8.
7) Valence:
The valence is the most characteristic property of the elements. It has been observed that the chemical properties of elements depend upon the number of electrons present in the outermost shell of the atom. The electrons present in the outermost shell are called valence electrons and these electrons determine the valence of the atom.
In case of representative elements, the valence is generally equal to either number of valence electrons or equal to eight minus the number of valence electrons. However, the transition elements, exhibit variable valency.
a) Variation in a period:- The number of valence electrons increases from 1 to 8 on moving across a period, the valency of the elements with respect to hydrogen and chlorine increases from 1 to 4 and then decreases to zero.
b) Variation in a group: – On moving down a group, the number of valence electrons remains same and, therefore, all the elements in a group exhibit same valency. For example, all the elements of group 1 have valency one and those of group 2 have valency two.