Chemical elements organised on the basis of electron configurations, atomic numbers, and recurring chemical properties, forms the periodic table.
Elements are arranged in the order of their increasing atomic number. The table rows are called periods and columns are called groups.
The history of the periodic table dates back over a century of growth in the understanding of chemical properties. In the year 1869, Dmitri Mendeleev published the table. He built his work based on earlier discoveries by scientists such as Antoine-Laurent de Lavoisier and John Newlands.
Dobereiner’s triads
Certain elements could be placed in groups of three such that the atomic mass of the middle element was approximately equal to the average of the atomic masses of the other two. Such groups were called "Dobereiner's triads."
For Example
- Strontium (atomic mass-88) has an atomic mass equal to the average of the atomic masses of barium and calcium (137+40/2= 88.5)
- Bromine (atomic mass 80) has an atomic mass equal to the average of the atomic masses of chlorine and iodine (35.5 + 127/2= 81.25)
Drawbacks of Dobereiner’s triads:
- All elements could not be grouped into triads; the rule did not hold for a few combinations even within the same family.
Newland's Law of Octaves
Every eighth element (starting from a given one) has properties similar to the first element, like the pattern seen in musical notes This law gave the first idea about periodicity in elements.
For Example
- Lithium and sodium (placed exactly eight places from it) were found to have similar properties.
- Beryllium and magnesium (just eight places behind) were also found to have identical traits.
Drawbacks of Law of Octaves:
- This law was applicable only till calcium.
- The discovery of several new elements made it difficult to justify the positioning of the elements based on this rule.
- Dissimilar elements were near each other while similar ones quite far away.
Mendeleev's Periodic Law
The properties of elements are a periodic function of their atomic masses. Mendeléev’s Periodic Table contains vertical columns called ‘groups’ and horizontal rows called ‘periods’.
Merits of Mendeleev's classification:
- A systematic study of elements: He arranged known elements in order of their increasing atomic masses considering the fact that elements with similar properties should fall in the same vertical column.
- Correction of atomic masses: It could predict discrepancies in the atomic masses of certain elements and make corrections. For example, the atomic mass of beryllium was changed from 13.5 (incorrect) to 9 (correct).
- It predicted the existence of elements not yet discovered.
- The position of noble gases: One of the greatest merits of Mendeleev’s classification was that when noble gases were discovered, they could be placed in a new group without disturbing the existing order.
Drawbacks of Mendeleev’s periodic table:
- Could not assign an appropriate position for hydrogen.
- Isotopes were discovered long after Mendeléev proposed his classification and posed quite a challenge to Mendeleev’s Periodic Law.
- The atomic masses did not increase in a regular manner in going from one element to the next, so it wasn't possible to predict how many elements could be discovered between two existing ones (especially when we consider the heavier elements).
The modern periodic table is designed depending on the atomic number of elements. According to Henry Moseley, the scientist who devised modern periodic table, the properties of elements are a periodic function of their atomic numbers. Elements, when arranged in the ascending order of their atomic numbers, the elements with the same number of valence electrons show periodicity.
Modern Periodic law:
- Moseley modified Mendeleev’s Periodic law to form the modern periodic law.
- The physical and chemical properties of elements are the periodic function of their atomic numbers.
- Atomic number is the total number of protons present in the nucleus and is responsible for the chemical properties of the element.
- When elements are arranged according to increasing atomic number, there is periodicity in the number of valence electrons and corresponding periodicity in chemical properties
- Periodicity refers to trends or recurring variations in element properties with increasing atomic number. Periodicity is caused by regular and predictable variations in element atomic structure.
- The cause of periodicity is the recurrence of similar electronic configuration.
- Elements have electrons arranged in specific energy levels around the nucleus:
- First shell- K
- Second shell- L
- Third shell- M…
- The number of electrons in each shell is calculated by the formula 2n2 where n is the number of the shell:
- K shell can have 2 x 1 2 = 2 electrons
- L shell can have 2 x 2 2 = 8 electrons
- M shell can have 2 x 3 2 =18 electrons
Periodic trends are specific patterns that are present in the periodic table that provides different aspects of a certain element, such as size, electronic property and so on. Periodic trends which are derived from the periodic table, are very helpful to chemists to estimate an element’s properties. The periodic nature of elements, in turn, leads to periodic trends.
Periodic Trends and Properties of elements
Important terms:
- Atomic size: This is the distance from the nucleus to the valence shell where the valence electrons are located.
- Ionization energy: The amount of energy required to remove the valence electron of an isolated gaseous atom to form a cation. It is measured in kJ/mole.
- Electron affinity: The electron affinity of an atom or molecule is defined as the amount of energy released when an electron is added to a neutral atom or molecule in the gaseous state to form a negative ion. It is measured in kJ/mole.
- Metallic character: The metallic character of an element can be defined as how readily an atom can lose an electron.
- Non-metallic character: The non-metallic character of an element can be defined as how readily an atom can gain an electron to attain stability.
- Electronegativity: Tendency of an atom in a molecule to draw the shared pair of electrons towards itself is called electronegativity. it does not have units since it is only a tendency.
Trends in Periodic Properties:
- Atomic size: Increases down the group (new shells are added).
- Atomic radius: Decreases across period (increase in nuclear charge).
- Valency increases from 1-4 and then decreases from 4-0 across the period.
- Valency remains the same down the group.
- Metallic nature and basic nature of oxides decrease across the period.
- Metallic nature decreases across the period (electropositivity decreases) and increases down the group (electropositivity increases).
- Non-metallic nature increases across period (electronegativity increases) and decreases down the group (electronegativity decreases).
- Ionization energy decreases down the group.
- Electron affinity increases down the group.