PERIODIC CLASSIFICATION OF ELEMENTS
The elements a
At present, 118 elements are known. Of these, only 94 are naturally occurring.
MAKING ORDER OUT OF CHAOS – EARLY ATTEMPTS AT THE CLASSIFICATION OF ELEMENTS
Elements have earlier grouped as metals & non-metals. Other attempts are given below:
DÖBEREINER’S TRIADS
A German chemist, Johann Wolfgang Döbereiner (1817) identified some groups having three elements each. So he called these groups triads.
He showed that when the three elements in a triad are written in the order of increasing atomic masses; the atomic mass of middle element was roughly the average of the atomic masses of the other two elements. E.g.
Triad 1 | Atomic mass |
Li | 6.9 |
Na | 23.0 |
K | 39.0 |
Triad 2 | Atomic mass |
Ca | 40.1 |
Sr | 87.6 |
Ba | 137.3 |
Triad 3 | Atomic mass |
Cl | 35.5 |
Br | 79.9 |
I | 126.9 |
But he could identify only 3 triads. Hence, system of classification into triads was not useful. E.g.
Elements | Atomic mass |
N | 14.0 |
Here, average of atomic mass of N & As is 44.45. So, it is not a Döbereiner triad.
NEWLANDS’ LAW OF OCTAVES
An English scientist, John Newlands (1866) arranged the elements in the order of increasing atomic masses. He started with hydrogen (lowest atomic mass) and ended at thorium which was the 56th element.
He found that every 8th element had properties similar to that of the first. He compared this to the octaves found in music. Therefore, he called it ‘Law of Octaves’.
E.g. Sodium is 8th element after lithium. So they resemble each other. Similarly, Be & Mg resemble each other.
A part of Newlands’ Octaves
Notes of music: | sa (do) | re (re) | ga (mi) | ma (fa) | pa (so) | da (la) | ni (ti) |
| H | Li | Be | B | C | N | O |
F | Na | Mg | Al | Si | P | S | |
Cl | K | Ca | Cr | Ti | Mn | Fe | |
Co & Ni | Cu | Zn | Y | In | As | Se | |
| Br | Rb | Sr | Ce & La | Zr | — | — |
Drawbacks of Newland’s Octaves:
o It is applicable only up to calcium. After this, every 8th element had no properties similar to that of the first.
o Newlands assumed that only 56 elements existed in nature. But new elements were discovered which did not obey Law of Octaves. To fit elements into his Table, he adjusted two elements in the same slot, and put some unlike elements under the same note.
E.g. Cobalt & nickel are in the same slot and are placed in the column of F, Cl & Br. But their properties are very different. Iron resembles Co & Ni. But it was placed far away from them.
o Discovery of noble gases made Law of Octave irrelevant. Thus, this law worked well with lighter elements only.
PERIODIC CLASSIFICATION OF ELEMENTS
MAKING ORDER OUT OF CHAOS – MENDELÉEV’S PERIODIC TABLE
Dmitri Ivanovich Mendeleev (1834-1907, a Russian chemist) was the most important contributor to the development of Periodic Table.
Mendeleev’s Periodic Table proved to be the unifying principle in chemistry. It was the motivation for the discovery of new elements.
Mendeleev started his work with 63 known elements. He examined relationship between the atomic masses (fundamental property) of the elements and their physical & chemical properties.
Mendeleev concentrated on the compounds (hydrides & oxides) formed by elements with oxygen & hydrogen. He selected these elements as they are very reactive and form compounds with most elements. The formulae of the hydrides & oxides were treated as a basic property.
He wrote down the properties of each element on 63 cards and sorted out them with similar properties. He found that most elements were arranged in the increasing order of atomic masses.
There was a periodic recurrence of elements with similar physical & chemical properties. Thus, he formulated a Periodic Law. It states that ‘the properties of elements are the periodic function of their atomic masses’.
In a few cases, Mendeleev had placed an element with slightly greater atomic mass before an element with slightly lower atomic mass. The sequence was inverted to group the elements with similar properties together. E.g. Cobalt (atomic mass 58.9) appeared before nickel (atomic mass 58.7).
MENDELEEV’S PERIODIC TABLE
Mendeleev’s Periodic Table contains groups (vertical columns) and periods (horizontal rows).
R= elements.
Mendeleev left some gaps in Periodic Table and predicted the existence of some elements.
He named them by prefixing Eka (one) to the name of preceding element in the same group.
E.g. scandium, gallium & germanium, discovered later, have properties similar to Eka–boron, Eka–aluminium and Eka–silicon, respectively.
Properties of eka–aluminium and gallium
Property | Eka-aluminium | Gallium |
Atomic Mass | 68 | 69.7 |
It proves the correctness and usefulness of Mendeleev’s Periodic Table.
Noble gases like helium (He), neon (Ne) & argon (Ar) were discovered very late because they are very inert and rarely present in the atmosphere. They could be placed in a new group without disturbing the existing order.
Limitations of Mendeleev’s Classification
1. No fixed position to hydrogen because:
· Its electronic configuration resembles that of alkali metals. Like alkali metals, hydrogen combines with halogens, oxygen & sulphur to form compounds having similar formulae. E.g.
Compounds of H | HCl | H2O | H2S |
Compounds of Na | NaCl | Na2O | Na2S |
· Like halogens, hydrogen exists as diatomic molecules and combines with metals and non-metals to form covalent compounds.
2. Isotopes of an element have similar chemical properties, but different atomic masses. So they cannot be placed in same slots. E.g. Cl-35 & Cl-37 are isotopes of chlorine.
3. Atomic masses do not increase in a regular manner from one element to the next. So it was not possible to predict how many elements could be discovered between two elements especially in case of heavier elements.
PERIODIC CLASSIFICATION OF ELEMENTS
MAKING ORDER OUT OF CHAOS – THE MODERN PERIODIC TABLE
Henry Moseley (1913) showed that the atomic number (Z) is a more fundamental property than atomic mass.
Atomic number= Number of protons in an atom’s nucleus.
Thus, Mendeleev’s Periodic Law was modified as follows:
‘Properties of elements are a periodic function of their atomic number.’
Arrangement of elements based on increasing atomic number led to Modern Periodic Table. In this, more precise prediction of properties of elements is possible.
Modern Periodic Table rectified three limitations of Mendeleev’s Periodic Table.
- Positions of Co & Ni resolved based on atomic number.
- Isotopes have same atomic number so they are placed in the same group.
- Atomic number is a whole number. So, there is no confusion about the presence of an element between two elements. E.g. There is no element with atomic number 1.5 placed between hydrogen and helium.
The Modern Periodic Table has 18 vertical columns (groups) and 7 horizontal rows (periods).
Groups signify an identical outer shell electronic configuration. All elements in a group contain same number of valence electrons. The number of shells increases as go down the group. E.g.
· Group 1 elements are H, Li, Na, K, Rb, Cs & Fr.
Electronic configuration of H = 1.
Electronic configuration of Li = 2, 1.
Electronic configuration of Na = 2, 8, 1.
Here, all elements have same number of valence electron (i.e., 1).
· Group 17 elements are fluorine (F), chlorine (Cl) etc. Their outermost shells contain 7 electrons.
There is an anomaly in case of the position of hydrogen. It can be placed in group 1 or 17 in the first period.
- Like group I elements (alkali metals), hydrogen has only one valence electron. Thus, it can lose an electron to achieve a stable configuration like alkali metals. Hence it can be placed in group 1.
- Like group 17 elements, it needs only one electron to complete its valence shell. Thus, it can gain an electron to achieve a noble gas configuration.
Elements in a period do not have the same number of valence electrons, but contain same number of shells. Also, the number of valence shell electrons increases by one unit, as the atomic number increases by one unit on moving from left to right. E.g.
2nd period elements & their electronic configuration:
Li | Be | B | C | N | O | F | Ne |
2,1 | 2,2 | 2,3 | 2,4 | 2,5 | 2,6 | 2,7 | 2,8 |
3rd period elements & their electronic configuration:
Na | Mg | Al | Si | P | S | Cl | Ar |
2,8,1 | 2,8,2 | 2,8,3 | 2,8,4 | 2,8,5 | 2,8,6 | 2,8,7 | 2,8,8 |
Atoms of different elements with the same number of shells are placed in the same period.
Number of elements in periods are based on how electrons are filled into various shells.
Maximum number of electrons that can be accommodated in a shell depends on the formula 2n2 (n= number of the shell). E.g.
- K Shell: 2 × (1)2 = 2 electrons. Hence 1st period has 2 elements. They have only one shell (K).
- L Shell: 2 × (2)2 = 8 electrons. Hence 2nd period has 8 elements. They have 2 shells (K & L).
- M shell: 2 × (3)2 = 18 electrons. 3rd period has 3 shells (K, L & M). Last shell can accommodate only up to 8 electrons. Hence 3rd period has only 8 elements.
- 4th, 5th, 6th & 7th periods have 18, 18, 32 & 32 elements respectively.
Mendeleev used formulae of compounds as a basic property to decide the position of an element. This was a good choice because elements are arranged in groups based on the number of valence electrons and valency. Since valency in a group is same, they will form similar formulae with hydrogen, oxygen etc. Thus they show similar chemical properties.
Trends in the Modern Periodic Table
Valency
It is the number of electrons that must be lost or gained by an atom to attain a stable configuration.
It is determined by the number of valence electrons present in the outermost shell of its atom.
Valency of a metal = Number of valence electrons.
E.g. Electronic configuration of Mg (Z= 12) is 2, 8, 2.
∴ Valency of Mg = 2.
Valency of a non-metal = 8 – No. of valence electrons.
E.g. Electronic configuration of S (Z= 16) is 2, 8, 6.
∴ Valency of S = 8 – 6 = 2.
Valencies of first 20 elements:
Elements | Atomic No. | E. Config. | Valency |
H | 1 | 1 | 1 |
He | 2 | 2 | 0 |
Li | 3 | 2, 1 | 1 |
Be | 4 | 2, 2 | 2 |
B | 5 | 2, 3 | 3 |
C | 6 | 2, 4 | 8 – 4 = 4 |
N | 7 | 2, 5 | 8 – 5 = 3 |
O | 8 | 2,6 | 8 – 6 = 2 |
F | 9 | 2, 7 | 8 – 7 = 1 |
Ne | 10 | 2, 8 | 8 – 8 = 0 |
Na | 11 | 2, 8, 1 | 1 |
Mg | 12 | 2, 8, 2 | 2 |
Al | 13 | 2, 8, 3 | 3 |
Si | 14 | 2, 8, 4 | 8 – 4= 4 |
P | 15 | 2, 8, 5 | 8 – 5= 3 |
S | 16 | 2, 8, 6 | 8 – 6= 2 |
Cl | 17 | 2, 8, 7 | 8 – 7= 1 |
Ar | 18 | 2, 8, 8 | 8 – 8= 0 |
K | 19 | 2, 8, 8, 1 | 1 |
Ca | 20 | 2, 8, 8, 2 | 2 |
In a period, from left to right, valency increases 1 to 4 then decreases from 4 to 0.
When going down a group, valency remains the same.
Atomic size (Atomic radius)
It refers to the radius of an atom. i.e., distance between the centre of the nucleus and outermost shell.
E.g. atomic radius of hydrogen atom is 37 pm (picometre, 1 pm = 10–12 m).
In a period, atomic radius decreases from left to right. This is due to an increase in nuclear charge which pull the electrons closer to the nucleus reducing atomic size. E.g.
Atomic radii of 2nd period elements in decreasing order:
Period II elements | Li | Be | B | C | N | O |
Atomic radius (pm) | 152 | 111 | 88 | 77 | 74 | 66 |
Here, Li has largest atom and O has smallest atom.
Atomic size increases down the group due to the addition of new shells. This increases distance between outermost electrons and nucleus so that the atomic size increases in spite of the increase in nuclear charge. E.g.
Atomic radii of 1st group elements in an increasing order. Here, Li has smallest atom and Cs has largest atom.
Elements | Atomic radius (pm) |
Li | 152 |
Na | 186 |
K | 231 |
Rb | 244 |
Cs | 262 |
Metallic & Non-metallic Properties
In Periodic Table, metals are found on the left side and the non-metals are found on the right side towards the top. A zig-zag line separates metals from non-metals. E.g.
Elements of third period:
Elements with Atomic No. | Configuration | Metal / Non-metal |
Na (11) | 2, 8, 1 | Metal |
Mg (12) | 2, 8, 2 | Metal |
Al (13) | 2, 8, 3 | Metal |
Si (14) | 2, 8, 4 | Metalloid |
P (15) | 2, 8, 5 | Non-Metal |
S (16) | 2, 8, 6 | Non-Metal |
Cl (17) | 2, 8, 7 | Non-Metal |
Ar (18) | 2, 8, 8 | Non-Metal |
In the middle, semi-metal or metalloid are found. They show intermediate properties of metals and non-metals. These borderline elements include boron, silicon, germanium, arsenic, antimony, tellurium & polonium.
Metals form bonds by losing electrons. So they are electropositive.
Metallic character decreases across a period and increases down a group because
- Across a period, the effective nuclear charge acting on the valence electrons increases. So, the tendency to lose electrons decreases.
- Down a group, the nuclear charge acting on valence electrons decreases as the outermost electrons are farther away from the nucleus. So, the electrons are lost easily.
Non-metals form bonds by gaining electrons. So, they are electronegative.
In a period, tendency of gaining electrons increases from left to right up to 17th group. It decreases in 18th group.
Tendency of gaining electrons decreases down a group.
These trends help to predict the nature of oxides formed by the elements because generally the metal oxides are basic and non-metal oxides are acidic.