At presently, 118 elements are known to us. Out of these 118, only 94 are naturally occurring. All these have different properties.
Classification of Elements
Elements are classified on the basis of similarities in their properties.
Dobereiner’s Triads
In the year 1817, Johann Wolfgang Dobereiner, tried to arrange the elements with similar properties into groups of three elements each. So, he called these groups ‘triads’.
Dobereiner's Triad:
Li | Ca | Cl |
Na | Sr | Br |
K | Ba | I |
He showed that when the three elements in a triad 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, take triad consisting of Lithium (Li), Sodium (Na) and Potassium (K) with the respective atomic masses 6.9, 22.9 (23.0) and 39.0.
Elements | Atomic Masses | Average |
Lithium (Li) Sodium (Na) Potassium (K) | 6.9 23.0 39.0 | 6.9 + 39.0 / 2 = 22.95 ~23.0 |
Calcium (Ca) Strontium (Sr) Barium (Ba) | 40.1 87.6 137.3 | 40.1 + 137.3 / 2 = 88.65 |
Chlorine (Cl) Bromine (Br) Iodine (I) | 35.5 79.9 126.9 | 35.5 + 126.9 / 2 = 81.2 |
Groups | Traids | Atomic masses |
Group A | N P As | 14.0 31.0 74.9 |
Group B | Ca Sr Ba | 40.1 87.6 137.3 |
Group C | Cl Br I | 35.5 79.9 126.9 |
From above table it was concluded that the only group B and C followed the Dobereiner’s Triads and group A did not follow the Triad rule.
Limitations:
Dobereiner could identify only three triads from the elements known at that time. Hence this system of classification into triads was not found much useful.
Newland’s Law of Octaves
In 1866, John Newland, arranged the then 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 first. He compared this to the octaves found in the music. Therefore, he called it the ‘Law of Octaves’. It is also known as ‘Newland’s Law of Octaves’
For example, the properties of lithium and sodium were found to be same. Sodium is the eighth element after lithium. Similarly, beryllium and magnesium resemble each other.
Limitations:
The Law of Octaves was applicable only upto calcium (i.e., upto atomic mass 40).
Newland assumed that only 56 elements existed in nature. So, properties of new elements did not fit into the Law of Octaves.
In order to fit elements in the table, Newland adjusted two elements in the same slot, but also put some unlike elements under the same note.
Observe 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.
With discovery of noble gases, the Laws of Octaves became irrelevant.
Thus, Newland’s Law worked well with lighter elements only.
Mendeleev’s Periodic Table
Mendeleev tried to arrange the elements on the basis of their fundamental properties, the atomic mass and also on the similarity of physical and chemical properties.
Among chemical properties, Mendeleev concentrated on the compounds formed by elements with oxygen and hydrogen. 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.
Mendeleev formulated a Periodic Law, which states that –
‘The properties of elements are the periodic function of their Atomic Masses’.
Mendeleev’s periodic table contains vertical columns called ‘Groups’ and horizontal rows called ‘Periods’.
Achievements of Mendeleev’s periodic table:
A systematic study of elements: Elements with similar properties grouped together.
Placement of Noble gases: They were easily place in the new group called Zero group in the table.
Prediction of yet to be discovered elements: Mendeleev named certain elements, Eka (one) to the name of preceding element in the same group. For instance, Scandium, Gallium and Germanium have properties similar to Eka-boron, Eka-aluminium and Eka-silicon, respectively.
Property | Eka-aluminium | Gallium |
Atomic mass | 68 | 69.7 |
Formula of Oxide | E2O3 | Ga2O3 |
Formula of Chloride | ECl3 | GaCl3 |
Limitation:
Position of Hydrogen: Hydrogen resembles both alkali and halogen in properties.
Position of Isotopes: Atomic weight differ, but they were not placed in different position.
Anomalous pair of elements: The atomic masses do not increase in a regular manner in going from one element to the next. Cobalt (Co) have higher atomic mass but was placed before Nickel (Ni).
Placement of like elements in different groups: Platinum (Pt) and Gold (Au) both have similar properties but were placed in different groups.
Compunds of H | Compounds of Na |
HCl | NaCl |
H2O | Na2O |
H2S | Na2S |
The Modern Periodic Table
In 1913, Henry Moseley showed that the atomic number (symbolized as Z) of an element is a more fundamental property than its atomic mass. Hence, he modified the Mendeleev’s Periodic table and the law was stated as follows:
‘Properties of elements are a periodic function of their Atomic Number’.
Position of Elements in the Modern Periodic Table:
According to their increasing atomic numbers, elements are organised.
The contemporary periodic table divides the elements into 7 periods and 18 groups.
Vertical columns are known as Groups, and horizontal rows are known as Periods.
Depending on how many atomic shells each element has, it is classified into periods.
The first period, which only has two elements—hydrogen and helium—is the shortest.
The sixth period in the periodic table is regarded as the longest period. It contains substances ranging from Radon to Cesium.
Actinides and Lanthanides are included at the bottom of the periodic table.
Metals like Na, Mg, etc are placed towards left-hand side while the non-metals like S, Cl, etc are placed towards right-hand side in the table.
A zig-zag line in the table separate’s metals from non-metals and have elements like boron, silicon, germanium, arsenic, antimony, tellurium and polonium. These are semi-metals or metalloids.
The elements present in any one group have the same number of valence electrons. For example, elements Fluorine and Chlorine, belong to group 17 and hence each of them contains 7 valence electrons in their outermost orbit.
Hence, groups in the periodic table signify an identical outer-shell electronic configuration. The number of shells increases as we move down the group.
Trends in the Modern Periodic Table
Elements thus arranged show periodicity of properties including atomic size, valency or combining capacity, metallic and non-metallic character, etc.
Trends in Properties | In the Group (from top to bottom) | In the Periods (from left to right) |
Valency (the number of valence electrons present in the outermost shell of its atom or the combining capacity of an element is known as its valency) | Remains same (because the number of valence electrons remains the same.) | Firstly, increases and then decreases  |
Atomic number | Increase | Increase |
Atomic size (refers to the radius of an atom or the distance between the centre of the nucleus and the outermost shell of an isolated atom) | Increases (because new shells are being added. This increases the distance between the outermost electrons and the nucleus) Example; Atomic size of first group element: Li < Na < K < Rb < Cs < Fr | Decreases Example: Size of second period elements: Li > Be > B > C > N > O > F |
Atomic radius | Increases | Decreases (due to an increase in nuclear charge which tends to pull the electrons closer to the nucleus and reduces the size of the atom) |
Metallic character | Increases (effective nuclear charge acting on valence electrons decreases and so the tendency to lose electrons increases) | Decreases (effective nuclear charge acting on valence electrons increases and so the tendency to lose electrons decreases) |
Non-metallic character | Decreases | Increases |
Electronegativity (tendency of an element to attract the shared pair of electrons towards it) | Decreases Example; Li > Na > K > Rb > Cs | Increases Example; Li < Be < B < C < N < O < F. |
Electron Affinity | Decreases | Increases |
Ionization energy (tendency to form ions) | Decreases | Increases |
Metals are electropositive, on the other hand non-metals are electronegative.
The oxides of metals are basic in nature while that of non-metals is acidic.