The elements which were undiscovered and for whom Mendeleev had left vacant places were discovered later. Some of these are scandium (Sc), gallium (Ga) and germanium (Ge). These elements were accommodated in their proper places in the table. The elements helium (He), neon (Ne), argon (Ar), Krypton (Kr), Xenon (Xe) and radon (Rn) became known only towards the end of the nineteenth century. These elements, called noble gases, were placed in the table as a separate group, called 0 group. The periodic table had to be modified then. The modified version of the table is shown below.
|Periods↓||A B||A B||A B||A B||A B||A B||A B|
|Fe Co Ni||Kr|
|Y In||Zr Sn||Nb Sb||Mo Te||Tc I||
Ru Rh Pd
|6||Cs Au||Ba Hg||La* Tl||Hf Pb||Ta Bi||W Po||Re At||Os Ir Pt||Rn|
(along with langhanum)
(along with actinium)
Features of the modified version of Mendeleev’s periodic table :
- Groups into subgroups : Each group of this periodic table is further divided into two subgroups A and B. The properties of elements within a subgroup resemble more markedly but they differ from those of the elements of the other subgroups. For example., lithium (Li), sodium (Na), potassium (K), etc., of subgroups IA have close resemblance of properties but they have hardly any resemblance to the coinage metals (Cu, Ag and Au) of subgroup IB. Mendeleev allowed the subgroups to be represented within the same group.
- Prediction of errors : This periodic table could predict errors in the atomic masses of some elements on the basis of their position in the periodic table. For example, when the periodic table was published, the experimental value of the atomic mass of beryllium (Be) we was supposed to be 13.65 and its valency, 3. So, the position of Be should have been somewhere else, but Mendeleev placed it at its appropriate position on the basis of its properties. He further suggested that the atomic mass of Be needed correction. Mendeleev predicted its atomic mass to be 9.1 and valency, 2. Latter investigations proved him right.
Similarly, the atomic mass of uranium was corrected from 120 to 240. Corrections were also made in the atomic masses of gold, platinum, etc.
- Predictions of properties of higher to undiscovered element : We know that Mendeleev classified the elements in order of their increasing atomic masses. However, this order had to be ignored at some places to make sure that the elements with similar properties fell in the same group. In doing so, he left some vacant places in the table. These vacant places were kept reserved for elements not discovered till then. Mendeleev was confident that these elements would be discovered later and they would occupy these vacant places. Not only this, he also predicted the properties of these undiscovered elements on the basis of this study of his the properties of the neighboring elements. Amazingly, when the missing elements of Mendeleev’s periodic table were discovered subsequently, their properties were found to be very similar to those predicted by Mendeleev.
The elements scandium, gallium and germanium were not known in 1871 but their existence was predicted by mendelev. He named these elements as eka-boron, eka-Aluminium and eka silicon when these elements were discovered, they were named scandium, gallium and germanium respectively and their properties were found to be in good agreement with those predicted by Mendeleev. Properties of ka-aluminium (predicted by Mendeleev) and those of the gallium (discovered later) are given in table.
|Formula of oxide||E2O3||Ga2O3|
|Formula of chloride||ECl3||GaCl3|
Considering its atomic mass, titanium (Ti) should have been placed below aluminium in the periodic table, but Mendeleev placed is below silicon (Si) because the properties of titanium were similar to those of silicon. Thus, a gap was left below aluminium in the periodic table. This gap was filled up by gallium which was discovered later. The properties of gallium (Ga) were found to be similar to those of boron and aluminium.
- Basic features intact : All the basic features of Mendeleev’s periodic table are intact even today. Even when a new class of elements, i.e., noble gases, were discovered, they found place in a separate group called the zero group. The existing order of the periodic table was not at all disturbed.
² Discrepancies in Mendeleev’s periodic table :
Mendeleev’s periodic table has the following defects.
- Position of hydrogen : The position of hydrogen in the periodic table is anomalous. Hydrogen resembles alkali metals (Li, Na, K, etc). So it may be placed in the group of the halogens (VII A).
- Position of lanthanides and actinides : The elements from atomic number 57 to 71 are collectively known as lanthanides. They do not have a proper place in the periodic table. They all have been placed at the same position in group III and period 6. Similarly, the actinides (atomic numbers 89-103) also have no proper place in the periodic table. These elements have also been placed in the same position, in group III and period 7.
- Some similar elements are separated, while some dissimilar elements have been placed in the group : Some similar elements are separated in the periodic table. For example, copper (Cu) and mercury (Hg), silver (Ag) and thallium (Tl), and barium (Ba) and lead (Pb). On the other hand, some dissimilar elements have been placed together in the same group. For example, copper (Cu), silver (Ag) and gold (Au) have been placed in group I along with the alkali metals. Similarly, manganese (Mn) is placed in the group of the halogens.
- Presence of some anomalous pairs of elements : In Mendeleev’s periodic table, elements are arranged in order of increasing atomic mass. In some places, this order has been ignored.
(a) The atomic mass of argon is 40 and that of potassium is 39. But argon is placed before potassium in the periodic table.
(b) The positions of cobalt and nickel are not in proper order. Cobalt (at. mass = 58.9) is placed before nickel (at. mass = 58.6).
(c) Tellurium (at. mass = 127.6) is placed before iodine (at. mass = 126.9).
(d) Thorium (at. mass = 232.12) is placed before protactinium (at. mass = 231)
- Position of isotopes : The isotopes of an element have no place in the periodic table.
The failure of Mendeleev’s periodic law to explain the wrong order of the atomic masses of some elements and the position of isotopes led scientists working in this field to conclude that atomic mass cannot be the basis for the classification of elements. There must be a more fundamental property of elements which can be the basis of classification.
Anomalous pairs of elements
|Group||0||IA||VIII||VIII||VI B||VII B||III B||III B|
² Modern Periodic Table :
Henry Moseley, an English physicist found that the atomic number (Z) was the fundamental property of an elements and not the atomic mass for classification of elements.
² Modern Periodic Law :
‘‘Properties of elements are periodic functions of their atomic numbers, i.e., the number of protons or electrons present in the neutral atom of an element.’’
² Long form of Periodic Table :
Arranged in increasing order of their atomic numbers.
The prediction of properties elements and their compounds can be made with precision. All drawbacks of Mendeleev’s Periodic Table vanish when the elements are arranged on the basis of increasing atomic numbers.
² Elements in a Group :
(1) They show similar chemical properties due to similar outer electronic configuration, i.e., same number of valence electrons.
(2) They have gradation in properties due to gradually varying attraction of the nucleus and the outer valence electrons as we go down the group.
² Main Features of the Long Form of the Periodic Table :
(1) It shows arrangement of elements based on modern periodic law.
(2) There are 18 vertical columns known as groups.
(3) There are 7 horizontal rows known as periods.
(4) Elements having similar outer electronic configurations, i.e., having same valence electrons have been placed in same groups, e.g.,
- In periods, elements in which the number of electrons in the outermost shell increases gradually in step one are placed, e.g.,
(K, L, shells)
(6) Each group in the table signifies identical outer shell electronic configuration i.e., same valence electrons, e.g., group 1 has 1 valence electron, group 2 has 2 valence electrons, group 13 has 3, group 14 has 4 valence electrons.
(7) Each period starts with filling of new shell, e.g.,
1st Period – K shell (1st shell) starts filling with Hydrogen and ends at Helium.
2nd Period – L shell (2nd shell) starts filling from Li (3) upto Ne (10)
3rd Period – M shell (3rd shell) start filling from Na (11) upto Ar (18)
4th Period – N shell (4th shell) starts filling from K (19) upto Kr (36) and so on.
(8) The periodic table is divided in four blocks :
(a) s-block elements : Group 1 and 2 elements are called s-block elements.
(b) p-block elements : Group 13 to 18 elements are called p-block elements
(c) d-block elements : Group 3 to group 12 are called d-block elements or transition elements (in between s- block and p-block elements)
(d) f-block elements : The elements placed at the bottom of the periodic table are known as f-block elements. The fourteen elements after La(57) (Lanthanum) are called Lanthanoides and 14 elements after Actinium Ac (89) are called Actinoides.
² Naturally occurring Elements :
Elements upto atomic number 92 occur in nature except Technetium, Tc (Z = 43) and Promethium, Pm (Z = 61) which are formed from radioactive elements where ‘Z’ represents atomic number.
Synthetic Elements : Elements beyond atomic number 92 are man-made elements. They are also called synthetic elements
(1) Elements in a group have same number of valence electrons.
(2) The chemical properties of valence electrons, e.g., all the group 1 elements have 1 valence electron. They form positively charged ions by losing one electron, when required amount of energy is supplied to them i.e., Li+, Na+, K+.
Group 1 elements are called alkali metals. Group 2 elements are called alkaline earth metals.
Group 2 elements when 2 valence electrons in the outermost shells. They can lose both the valence electrons to form dipositive cations, i.e., Be2+, Mg2+, Ca2+, etc. Positively charged ions are called cations.
Group 13 elements belong to boron family, 14 to carbon family, 15 to Nitrogen family, 16 to Oxygen family.
Group 16 elements contain 6 valence electrons in their outermost shells, i.e., two electrons less than the maximum number of electrons that can be present in the outermost shell. They can gain 2 electrons more easily rather than lose 6 electrons. They change into dinegative ions such as O2–, S2–. Group 17 elements called Halogens contain 7 valence electrons. They can gain one electron to acquire stable electronic configuration, i.e., 8 electrons in the outermost shell and form uni-negative (single negative) ions such as F–, Cl–, Br–, I–.
Negatively charged ions are called anions.
Group 18 elements called noble gases, have their outermost, shell completely filled. The elements of this group have no tendency to lose or gain electrons. Thus, the elements of this group have zero valency and are almost unreactive. Hence they are called Noble gases. However, nowadays, compounds of Kr, Xe and Rn have been prepared.
In any particular group, the number of shells increase but the number of valence electrons remains the same.
² Periods :
(1) The horizontal rows in the periodic table are called periods.
(2) There are 7 periods in the long form of periodic table
(3) The first period contains 2 elements, Hydrogen and Helium. They have only one shell.
(4) The second period contains 8 elements : Lithium Li(3), Beryllium Be(4), Boron B(5), Carbon C(6), Nitrogen N(7), Oxygen O (8), Fluorine F(9) and Neon Ne (10). The second period has 2 shells (K and L) and L shell is progressive filled.
(v) The elements of 3rd period are :
|(K, L, M shells)||2, 8, 1||2, 3, 2||2, 8, 3||2, 8, 4||2, 8, 5||2, 8, 6||2, 8, 7||2, 8, 8|
In 3rd period, 3rd shell (M-shell) is being progressively filled and there are three shells.
4th period has 18 elements
5th period has 18 elements
6th period has 32 elements
7th period has 32 elements
In periods, the number of valence electrons increases from left to right in s and p-blocks
² Periodicity in Properties :
The properties of elements depends upon the electronic configuration which changes along a period and down a group in periodic table.
There is periodicity in properties, i.e., repetition of properties after a regular interval due to similarity in electronic configuration.
² Atomic Size (Atomic radii) :
Atomic size means radius of an atom. It is defined as distance between centre of nucleus and outermost shell (valence shell) of an isolated atoms.
² Covalent Radii :
It is defined as half of the distance between the centres of nuclei of two atoms (bond length) bonded by a single covalent bond, e.g., Bond length in case of
H—H (Hydrogen molecule) is 74 pm.
Covalent radius = 1/2 × 74 pm = 37 pm (picometre)
[1 pm = 10–12 m]
It can be measured in case of diatomic molecules of non-metals.
² Metallic Radii : If is defined as half of the internuclear distance between the two metal ions in a metallic crystal. It is measured in case of metals.
² Variation of Atomic size in a Group :
Size generally increases from top to bottom in a group.
Reason : It is due to addition of a new shell, i.e., number of shells go one increasing, e.g., pm stands for picometre, i.e., 10–12 m.
|Group I||Electronic Configuration||No. of Shells||Atomic radius (pm)|
|Li(3)||2, 1||(2 shells)||133|
|Na(11)||2, 8, 1||(3 shells)||154|
|K(19)||2, 8, 8, 1||(4 shells)||201|
|Rb(37)||2, 8, 18, 8, 1||(5 shells)||216|
|Cs(55)||2, 8, 18, 18, 8, 1||(6 shells)||235|
² Variation of Atomic size along a Period :
Atomic size goes on decreasing along a period from left to right
Reason : It is due to increase in nuclear charge (number of protons in nucleus) which pulls the electrons towards it, i.e., force of attraction between nucleus and valence electrons increase, therefore atomic size decreases, e.g.,
|Elements of 2nd Period Atomic radius (pm)||Elements of 3rd Period Atomic radius (pm)|
² Ionisation Energy and Electron Affinity :
Chemical nature and reactivity of an element depend upon the ability of its atoms to donate or accept electrons. This can be measured quantitatively with the help of ionisation energy and electron affinity of its atoms :
Ionisation energy : It is defined as the energy required to remove an electron completely from an isolated gaseous atom of an element. The energy required to remove the 1st electron is called first ionisation energy.
A(g) + I.E1 ¾® A+(g) + e–
² Second Ionisation Energy : he energy required to remove an electron from a unipositive ion is called the second ionisation energy.
A+(g) + I.E2 ¾® A2+ (g) + e–
Te second ionisation energy is greater than the first ionisation energy due to increase in positive charge, hence increase in force of attraction between the nucleus and the valence electron.
Ist I.E. < 2nd I.E. < 3rd I.E.
² Variation of Ionisation energy down a Group :
Ionisation energy goes on decreasing down a group.
Reason : It is due to the increase in the distance between the valence electrons and the nucleus as the atomic size increase down a group, the force of attraction between the nucleus and the valence electrons decrease, therefore, the energy required to remove the electrons, i.e., the ioisation energy goes on decreasing
|Group I||Ionisation Energy
(in kJ mol–1)
|Group 2||Second Ionisation Energy
(in kJ mol–1)
|First Ionisation Energy (in kJ mol–1)|
² Variation of Ionisation energy along a Period :
It goes on increasing generally along a period from left to right with decrease in atomic size.
Reason : Due to decrease in atomic size, the force of attraction between the valence electrons and the nucleus increase and, therefore, more energy is required to remove electron.
|Elements of 2nd Period||I.E. in kJ mol–1|
There is a decrease in ionisation energy from Be to B and from N to O, the reason of which you will study in higher classes.
Group 18 elements (noble gases) have the highest ionisation energy in respective periods due to stable electronic configuration, i.e., 8 electrons in their valence shells except He which has 2 electrons.
² Electron Affinity :
It is the energy change when an electron is gained by a neutral gaseous atom converting it into a negatively charged ion.
It is a measure of attraction or affinity of the gaseous atom for an extra electron to be added.
Cl(g) + e– ¾® Cl–(g) + E.A.
² Factors :
It depends upon atomic size as well as electronic configuration.
² Variation down the Group :
Electron affinity goes on decreasing down the group in general.
Reason : Due to the increase in atomic size, the force of attraction between the nucleus and the electron to e added becomes less.
² Variation along a Period :
Electron affinity increase from left to right in period.
Reason : It is due to decrease in atomic size which leads to an increase in the force of attraction between the nucleus and the electrons to be added.
|Group 17||E.A. (kJ mol–1)|
However, deviations to this rule are observed in variation of electron affinity.
² Metallic and Non-metallic Character :
Group 1 to 12 are metals. Group 13 to 18 comprise non-metals, metalloids and metals.
Metalloids : Those elements which resemble both metals and non-metals are called metalloids. They are also called semi-metals, e.g., Boron, Silicon, Germinaium, Arsenic, Antimony, Tellurium and Polonium.
Properties of Metals :
(i) The are malleable.
(ii) They are ductile.
(iii) They are good conductors of heat and electricity.
(iv) They have generally 1 to 3 valence electrons.
(v) They have the same or less number of electrons in their outermost shell than the number of shells.
(vi) They are mostly solids.
Properties of Non-metals :
(i) They exist in solid, liquid or gaseous state.
(ii) Non-metals are generally brittle.
(iii) They are non conductors.
- They have 4 to 8 valence electrons.
² Variation of Metallic Character :
Metallic character increases down a group due to decrease in ionisation energy. It decrease along a period due to increase in ionisation energy from left to right
² Variation of Non-metallic Character :
Non-metallic character decreases down a group because of decrease in electron affinity which is due to increase in atomic size.
Along a period, non-metallic character increases from left to right due to increase in electron affinity which is due to decrease in atomic size