Periodic Classification of Elements Notes - Dobereiner Triads, Newlands Octaves, Mendeleev and Modern Periodic Table

• Elements are substances that contain only one type of atom.
• Currently, 118 elements are known, of which 94 are naturally occurring.

MAKING ORDER OUT OF CHAOS – EARLY ATTEMPTS AT THE CLASSIFICATION OF ELEMENTS

• Elements were initially grouped as metals and non-metals.

Döbereiner’s Triads

• In 1817, German chemist Johann Wolfgang Döbereiner identified groups of three elements, called triads.
• In a triad, when elements are arranged by increasing atomic mass, the atomic mass of the middle element is roughly the average of the other two.
• Examples of triads:

• However, only three triads were identified, making this classification system limited. For example, the elements N (14.0), P (31.0), and As (74.9) do not form a triad, as the average atomic mass of N and As (44.45) does not match P’s atomic mass.

Newlands’ Law of Octaves

• In 1866, English scientist John Newlands arranged elements in order of increasing atomic masses, from hydrogen to thorium (the 56^th element).
• He observed that every 8^th element had properties similar to the first, resembling musical octaves, hence named the Law of Octaves.
• Examples: Sodium (8^th after lithium) resembles lithium; beryllium and magnesium resemble each other.
• A part of Newlands’ Octaves:

• Drawbacks of Newlands’ Octaves:
• It was applicable only up to calcium. After calcium, every 8^th element did not share similar properties with the first.
• Newlands assumed only 56 elements existed. New elements discovered later did not obey the Law of Octaves. To fit elements, he placed two elements in the same slot (e.g., cobalt and nickel) and grouped unlike elements together (e.g., cobalt and nickel with fluorine, chlorine, and bromine).

Iron, which resembles cobalt and nickel, was placed far away.

• The discovery of noble gases made the Law of Octaves irrelevant, as it worked well only with lighter 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 the 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 the 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 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.

Achievements of Mendeleev’s Periodic Table

• Mendeleev left some gaps in the Periodic Table and predicted the existence of some elements.
• He named them by prefixing Eka (one) to the name of the 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.

• 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

• No fixed position for 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.
• Isotopes of an element have similar chemical properties, but different atomic masses. So they cannot be placed in the same slots. E.g., Cl-35 & Cl-37 are isotopes of chlorine.
• 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 the case of heavier 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 the 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 the 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.

Position of Elements in the Modern Periodic Table

• 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 the same number of valence electrons. The number of shells increases as you 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 the same number of valence electrons (i.e., 1).
• Group 17 elements are fluorine (F), chlorine (Cl), etc. Their outermost shells contain 7 electrons.
• There is an anomaly in the case of the position of hydrogen. It can be placed in group 1 or 17 in the first period.
• Like group 1 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 the 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.:

2^nd period elements & their electronic configuration:

3^rd period elements & their electronic configuration:

• Atoms of different elements with the same number of shells are placed in the same period.
• Number of elements in periods is 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, 1^st period has 2 elements. They have only one shell (K).
• L Shell: 2 × (2)² = 8 electrons. Hence, 2^nd period has 8 elements. They have 2 shells (K & L).
• M Shell: 2 × (3)² = 18 electrons. 3^rd period has 3 shells (K, L & M). Last shell can accommodate only up to 8 electrons. Hence, 3^rd period has only 8 elements.
• 4^th, 5^th, 6^th & 7^th 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 the 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.

• In a period, from left to right, valency increases from 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., the distance between the center of the nucleus and the outermost shell.
• E.g., the atomic radius of a hydrogen atom is 37 pm (picometer, 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 pulls the electrons closer to the nucleus, reducing atomic size. E.g.:

• Here, Li has the largest atom, and O has the smallest atom.
• Atomic size increases down the group due to the addition of new shells. This increases the distance between the outermost electrons and the nucleus, so the atomic size increases despite the increase in nuclear charge. E.g.:

• Here, Li has the smallest atom, and Cs has the largest atom.

Metallic & Non-metallic Properties

• In the Periodic Table, metals are found on the left side, and non-metals are found on the right side towards the top. A zig-zag line separates metals from non-metals. E.g.:

• 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, and 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, the tendency to gain electrons increases from left to right up to the 17^th group. It decreases in the 18^th group.
• The tendency to gain electrons decreases down a group.
• These trends help predict the nature of oxides formed by the elements because generally, metal oxides are basic, and non-metal oxides are acidic.