CLASS 10 Science
CHAPTER 5
Periodic Classification Of Elements (NCERT Notes)
CHAPTER 5
Periodic Classification Of Elements (NCERT Notes)
The earliest attempt to classify the elements resulted in grouping the known elements as metals and non-metals. Later further classifications were tried out as our knowledge of elements and their properties increased.
Döbereiner’s Triads
In 1817 Johann Wolfgang Döbereiner, a German chemist, tried to arrange the elements with similar properties into groups. He identified some groups having three elements each. So he called these groups ‘triads’. Döbereiner showed that when the three elements in a triad were written in the order of increasing atomic masses; the atomic mass of the middle element was roughly the average of the atomic masses of the other two elements. For example triad consisting of lithium (Li), sodium (Na)
and potassium (K) with the respective atomic masses 6.9, 23.0 and 39.0.
What is the average of the atomic masses of Li and K is 26.4 which is nearly equal to that of Na.
Limitations of Döbereiner
Döbereiner could identify only three triads from the elements known at that time. Hence, this system of classification into triads was not found to be useful.
Newlands’ Law of Octaves
In 1866 John Newlands, an English scientist, arranged the known elements in the order of increasing atomic masses. He started with the element having the lowest atomic mass (hydrogen) and ended at thorium which was the 56th element. He found that every eighth element had properties similar to that of the first. He compared this to the octaves found in music. Therefore, he called it the ‘Law of Octaves’. It is known as ‘Newlands’ Law of Octaves’.
In Newlands’ Octaves, the properties of lithium and sodium were found to be the same. Sodium is the eighth element after lithium. Similarly, beryllium and magnesium resemble each other.
Limitations of Newlands’ Law of Octaves:
It was found that the Law of Octaves was applicable only upto calcium, as after calcium every eighth element did not possess properties similar to that of the first.
It was assumed by Newlands that only 56 elements existed in nature and no more elements would be discovered in the future. But, later on, several new elements were discovered, whose properties did not
fit into the Law of Octaves.
In order to fit elements into his Table, Newlands adjusted two elements in the same slot, but also put some unlike elements under the same note. Note that cobalt and nickel are in the same slot and these are placed in the same column as fluorine, chlorine and bromine which have very different properties than these elements. Iron, which resembles cobalt and nickel in properties, has been placed far away from these elements.
MAKING ORDER OUT OF CHAOS – MENDELÉEV’S PERIODIC TABLE
The main credit for classifying elements goes to Dmitri Ivanovich Mendeléev, a Russian chemist. He was the most important contributor to the early development of a Periodic Table of elements wherein the elements were arranged on the basis of their fundamental property, the atomic mass, and also on the similarity of chemical properties. When Mendeléev started his work, 63 elements were known. He
examined the relationship between the atomic masses of the elements and their physical and chemical properties. Among chemical properties, Mendeléev concentrated on the compounds formed by elements with oxygen and hydrogen. He selected hydrogen and oxygen as they are very reactive and formed compounds with most elements. The formulae of the hydrides and oxides formed by an element were treated as one of the basic properties of an element for its classification.
He then took 63cards and on each card he wrote down the properties of one element. He sorted out the elements with similar properties and pinned the cards together on a wall. He observed that most of the elements got a place in a Periodic Table and were arranged in the order of their increasing atomic
masses. It was also observed that there occurs a periodic recurrence of elements with similar physical and chemical properties.
On this basis, Mendeléev formulated a Periodic Law, which states that ‘the properties of elements are the periodic function of their atomic masses’.
Mendeléev’s Periodic Table contains vertical columns called ‘groups’ and horizontal rows called ‘periods’.
Mendeléev’s Periodic Table was published in a German journal in 1872. In the formula for oxides and hydrides at the top of the columns, the letter ‘R’ is used to represent any of the elements in the group. Note the way formulae are written. For example, the hydride of carbon, CH4, is written as RH4 and the oxide CO2, as RO2.
Achievements of Mendeléev’s Periodic Table
Mendeléev sequence an element with a slightly greater atomic mass before an element with a slightly lower atomic mass. The sequence was inverted so that elements with similar properties could be grouped together. For example, cobalt (atomic mass 58.9) appeared before nickel (atomic mass 58.7).
Mendeléev left some gaps in his Periodic Table. Instead of looking upon these gaps as defects, Mendeléev boldly predicted the existence of some elements that had not been discovered at that time.
Mendeléev named them by prefixing a Sanskrit numeral, Eka (one) to the name of preceding element in the same group. For instance, scandium, gallium and germanium, discovered later, have properties similar to Eka–boron, Eka–aluminium and Eka–silicon, respectively. The properties of Eka–Aluminium predicted by Mendeléev and those of the element, gallium which was discovered later and replaced Eka- aluminium.
This provided convincing evidence for both the correctness and usefulness of Mendeléev’s Periodic Table.
Noble gases like helium (He), neon (Ne) and argon (Ar) show different property from other element of periodic table. These gases were discovered very late because they are very inert and present in extremely low concentrations in our atmosphere. One of the strengths of Mendeléev’s Periodic Table was that, when these gases were discovered, they could be placed in a new group without disturbing the existing order.
Limitations of Mendeléev’s Classification
Electronic configuration of hydrogen resembles that of alkali metals. Like alkali metals, hydrogen combines with halogens, oxygen and sulphur to form compounds having similar formulae. For example Compound of H as in HCl and H2S , Compound of Na as in NaCl and Na2S.
Just like halogens, hydrogen also exists as diatomic molecules and it combines with metals and non-metals to form covalent compounds.
Certainly, no fixed position can be given to hydrogen in the Periodic Table. This was the first limitation of Mendeléev’s Periodic Table. He could not assign a correct position to hydrogen in his Table.
Isotopes were discovered long after Mendeléev had proposed his periodic classification of elements. So Mendeléev’s could not explain about the Isotopes in periodic table. Isotopes of an element have similar chemical properties, but different atomic masses.
Another problem was that the atomic masses do not increase in a regular manner in going from one element to the next. So it was not possible to predict how many elements could be discovered between
two elements.
THE MODERN PERIODIC TABLE
In 1913, Henry Moseley showed that the atomic number of an element is a more fundamental property than its atomic mass.
The Modern Periodic Law can be stated as follows: ‘Properties of elements are a periodic function of their atomic number.’ The atomic number gives us the number of protons in the nucleus of an atom and this number increases by one in going from one element to the next. Elements, when arranged in order of increasing atomic number Z, lead us to the classification known as the Modern Periodic Table.
Position of Elements in the Modern Periodic Table
The Modern Periodic Table has 18 vertical columns known as ‘groups’ and 7 horizontal rows known as ‘periods’.
Elements present in any one group have the same number of valence electrons. For example, elements fluorine (F) and chlorine (Cl), belong to group 17, have 9 electrons in fluorine and 17 in chlorine have in their outermost shells. And belongs to same groups in the Periodic Table signify an identical outer-
shell electronic configuration. Because they have same no of valence electrons.
The number of shells increases as we go down the group.
Hydrogen can show anomaly behaviour when it comes to the position of because it can be placed either in group 1 or group 17 in the first period.
The number of valence shell electrons increases by one unit, as the atomic number increases by one unit on moving from left to right in a period.
Atoms of different elements with the same number of occupied shells are placed in the same period. Na, Mg, Al, Si, P, S, Cl and Ar belong to the third period of the Modern Periodic Table, since the electrons in the atoms of these elements are filled in K, L and M shells.
The maximum number of electrons that can be accommodated in a shell depends on the formula 2n2 where ‘n’ is the number of the given shell from the nucleus.
For example,
K Shell – 2 × (1)2 = 2, hence the first period has 2 elements.
L Shell – 2 × (2)2 = 8, hence the second period has 8 elements.
M Shell – 2 × (3)2 = 18, but the outermost shell can have only
8 electrons, so the third period also has only 8 elements.
The position of an element in the Periodic Table tells us about its chemical reactivity. The valence electrons determine the kind and number of bonds formed by an element.
Trends in the Modern Periodic Table
Valency : The valency of an element is determined by the number of valence electrons present in the outermost shell of its atom.
Atomic size: The atomic size refers to the radius of an atom. The atomic size may be visualised as the distance between the centre of the nucleus and the outermost shell of an isolated atom. The atomic radius of hydrogen atom is 37 pm (picometre, 1 pm = 10–12m).
The atomic radius decreases in moving from left to right along a period. This is due to an increase in nuclear charge which tends to pull the electrons closer to the nucleus and reduces the size of the atom.
The atomic size increases down the group. This is because new shells are being added as we go down the group. This increases the distance between the outermost electrons and the nucleus so that the atomic size increases in spite of the increase in nuclear charge.
Metallic and Non-metallic Properties
The metals like Na and Mg are towards the left-hand side of the Periodic Table while the non-metals like sulphur and chlorine are found on the right-hand side. In the middle, we have silicon, which
is classified as a semi-metal or metalloid because it exhibits some properties of both metals and non-metals.
In the Modern Periodic Table, a zig-zag line separates metals from non-metals. The borderline elements – boron, silicon, germanium, arsenic, antimony, tellurium and polonium – are intermediate in properties and are called metalloids or semi-metals.
Metals tend to lose electrons while forming bonds, that is, they are electropositive in nature.
As the effective nuclear charge acting on the valence shell electrons increases across a period, the tendency to lose electrons will decrease. Down the group, the effective nuclear charge experienced by valence electrons is decreasing because the outermost electrons are farther away from the nucleus. Therefore, these can be lost easily. Hence metallic character decreases across a period and increases down a group.
Non-metals are electronegative. They tend to form bonds by gaining electrons.