Periodic Classification Of Elements

John Alexander Reina Newland was the first to discover the Periodic Table. He made a table of all known elements and arranged them on the basis of their atomic weight. Lithium is the first element and sodium is the eighth element. Similarly, the eighth element after sodium is potassium. Law of octaves says that the elements Li, Na and K must have similar chemical and physical properties. For example, if we start from Lithium (atomic mass = 6) and arrange elements in increasing order of their atomic masses, the eighth element will be sodium (atomic mass = 23). Both elements show the same physical and chemical properties.
This classification was not applicable to all elements. It was only valid for elements having atomic masses lower than Ca.

Dmitri Ivanovich Mendeleev was a Russian scientist. He arranged the elements in increasing order of their relative atomic masses. This law states that the properties of elements are the periodic function of their relative atomic masses.

In 1869, after Newlands Octave Law was rejected, Mendeleev Periodic table was introduced. In this periodic table, elements were arranged on the basis of their atomic masses. A few gaps were left for the elements to be discovered later. Later, Gallium (Ga) and Germanium (Ge) were found that had same properties as eka–aluminum and eka–silicon, respectively.

Henry Moseley removed defects in earlier periodic tables and stated the modern periodic table. The modern periodic table stated that the properties of elements are periodic functions of their atomic number. The elements are arranged on the basis of their increasing atomic number.

This period contains the elements beryllium, boron, carbon, oxygen and fluorine.

Elements A is fluorine (atomic number 9) and C is chlorine (atomic number 17) belong to the same group. These two elements belong to group 17 of the periodic table.

Group 18 in the modern periodic table has noble gases (inert gases). Helium (He), neon (Ne), argon (Ar), krypton (Kr), xenon (Xe) and radon (Rn) are present in group 18. The members of the group have eight electrons in their outermost orbit (except helium which has two electrons). Thus, they have a stable electronic configuration. These gases are chemically unreactive, that is, they don’t react with other elements to form compounds.

Group 14 is known as the carbon family. It contains elements such as carbon (C), silicon (Si), germanium (Ge), tin (Sn), lead (Pb) and flerovium (Fl). Carbon is one of the most common and abundant elements on earth. It is an essential constituent of all organic compounds. Carbon dioxide (CO₂) and methane (CH₄) are a few common compounds containing carbon.

The periodic table has seven rows called periods. It has 18 columns called groups or families. Elements of the same group have similar properties. For example, group 18 consists of the noble or insert gases. These elements (helium, neon, argon, krypton, xenon, radon) have 8 valence electrons.

Periods show the number of outermost shell of any element. Since period 2 has two shells (2s and 2p orbitals), the outermost shell is L shell. The second period contains the elements such as lithium, beryllium, boron, carbon, nitrogen, oxygen, fluorine, and neon.

Among the given elements, phosphorus (P) has maximum number of valence electrons i.e., 5. Na (group 1) has one, Al (group 13) has three, Si (group 14) has four and P (group 15) has five valence electrons.

Atomic radius of fluorine (F) 42. Atomic radius of oxygen (O) 48. Atomic radius of nitrogen (N) is 56. Generally, atomic radius increases down a Group, from top to bottom. Thus, atomic radius will decrease from N to F.

Na and Mg belong to same period whereas Na and K are in the same group. Atomic radius increases in a group and decreases in a period. Potassium (K) has the largest atomic radii (243 pm). Sodium (Na) has 190 pm. Magnesium (Mg) has 145 pm. Calcium (Ca) has 194 pm.

Potassium (K) is more electropositive and hence ionizes easily than sodium (Na) due to bigger size and less electronic attraction. As we move down a group in periodic table, the size of the atom increases and the ionization potential decreases. So, potassium releases electrons easily than Na, Mg and Ca.

Flourine has 7 electrons in the outermost orbit. Its atom size is smallest too. Since it is the most electronegative element, thus does not lose an electron easily.
Sodium has one valence electron, Magnesium has 2 valence electrons, and Aluminum has 3 valence electrons. These are all metals that lose electrons easily.

Isotopes are two or more forms of an element in which the atoms have the same number of protons. Also different isotopes of an element have different mass numbers, hence the number of neutrons in the nuclei of isotopes of an element are different. All isotopes have similar chemical properties.

For example, hydrogen (H) has three isotopes having mass numbers 1, 2 and 3, but all having atomic number equal to 1.

Metallic character decreases on going from left to right in a period. This is due to decreased atomic size which increases the tendency of an element to gain electrons. The metallic character increases on going down in a group due to tendency to lose electrons.

Non-metallic character increases on going from left to right in a period. This is due to decreased atomic size which increases the tendency of an element to gain electrons. Non-metallic character decreases on moving down a group.

Aluminum (Al) has the oxidation state: +3. It can form Al³⁺ as it has valency of +3. So it will form Al₂ O₃.

Boron (B), silicon (Si), germanium (Ge) are metalloids. A few others are arsenic (As), antimony (Sb), tellurium (Te), polonium (Po) and astatine (At). These elements are found between metals and non-metals of the periodic table. They have properties of both metals and non-metals.

Non-metals form acidic oxides. The element with atomic number 7 is a non-metal (nitrogen). All others are metals.

Non-metal silicon has atomic number 14. It possess properties of both metals and non-metals. It forms acidic oxides like other non-metals.

Atomic radius is the distance between centre of the nuclei and the outermost shell.

Elements that belong to the same group has same number of valence electrons. Thus the valency of elements remains the same.

As we move from left to right in a period, the atomic number increases but the number of shell remains same. Thus electronic attraction increases thus reducing the size.

As we move down in a group the metallic nature of elements increases. Be Mg Ca are present in the same group (group 2 of modern periodic table). Thus, Ca has the highest metallic character.

The arrangement of these elements is known as Dobereiner triad. This triad is a group of three elements.

In 1817, Dobereiner suggested relationship between properties of elements and their atomic weights. The atomic weight of the middle element is approximately equal to the average weights of other two elements.

One example of such as set of elements is lithium, sodium and potassium.

Here, atomic weight of sodium is average weights of the lithium and potassium.

7 (lithium) + 39 (potassium)/2 = 23 (sodium)

(i) F and Cl have similar properties. (ii) Na and K have similar properties..

(b) The given sequence represents Newland’s law of octaves. Newland’s law of octaves states that every eighth element will show similar properties as that of the first element. Therefore, above two sets will have similar properties.

No. Na, Si, Cl cannot be classified as Dobereiner's triad. Although the atomic mass of silicon (Si) is average of atomic masses of sodium (Na) and chlorine (Cl), all these elements do not have similar properties.

23 (Na) + 35 (Cl)/2 = 29

(b) Yes. Be, Mg, Ca cannot be classified as Dobereiner's triad. They have similar properties and the atomic weights of magnesium (Mg) is roughly the average of the atomic weights of Be and Ca.

9 (Be) + 40 (Ca)/2 = 24.5

Dobereiner’s triad suggested relationship between properties of elements and their atomic weights. The atomic weight of the middle element is approximately equal to the average weights of other two elements.

In Mendeleev’s Periodic Table, elements with similar properties can be grouped together.

Cobalt was placed in the group of rhodium (Rh) and iridium (Ir) because cobalt has properties similar to these elements.

Likewise, nickel was placed with elements such as pallidium (Pd) and platinum (Pt) because nickel has properties similar to these elements.

Hydrogen occupies a unique position in Modern Periodic Table because of the following reasons:

a. The valence shell configuration of hydrogen and alkali metals is same. Both have only one electron. Thus some of the properties of hydrogen and alkali metals are similar and hence they are placed in group 1.

b. Both hydrogen and halogens need just one electron to complete their valence configuration. Thus they shows similar properties and are placed in group 17.

c. Apart from these properties, hydrogen shows some other unique properties such as they form oxides that are neutral in nature. Metals form basic oxides and halogen form acidic oxides.

GeCl₄, GaCl₃

‘Eka-element’ term was introduced by D. I. Mendeleev in 1871. The elements were arranged in the periodic table. But all the eka-elements were discovered later. Eka-silicon discovered in 1886 was called germanium (Ge). Eka-aluminum discovered in 1875 was called gallium (Ga). Eka-boron discovered in 1879 was called scandium. The prefix ‘eka’ means ‘beyond’ on his table.

The valency of germanium is 4 so the chemical formula of its chloride is GeCl_4.

The valency of gallium is 3 so the chemical formula of its chloride is GaCl_3.

Element A has 3 valence electrons. Thus it has valency 3. It should be in group 13 that contains elements such as B, Al, Ga, In and Tl.

Element A has 4 valence electrons. Thus it has valency 4. It should be in group 14 that contains elements such as C, Si, Ge, Sn and Pb.

Element A has 2 valence electrons. Thus it has valency 2. It should be in group 2 that contains elements such as Be, Mg, Ca, Sr and Ba.

If an element X is placed in group 14, its chemical formula is XCl₄.

The nature of bonding of its chloride is covalent bonding.

Element X has 4 valence electrons. Thus, it has valency 4. This element must show covalent bonding in order to complete octet configuration. So the chemical formula would be XCl4.

Since this compound is form by sharing of electrons, the bond is covalent in nature.

Radii of Y is less than X.

Here, X has 12 electrons and 12 protons, therefore it is a neutral atom. Y contains 12 protons and 10 electrons so there are 2 more protons that gives Y a charge of +2. The electronic configurations of the two species is as follows:

The atomic size of a positive ion (cation) is always smaller than a neutral atom. This is because cation has less number of electrons and hence more nuclear attractions on electrons in cation. Thus the size of a cation is smaller than that of the neutral atom. Hence, atomic radius of Y is smaller than that of X.

F < N < Be < Li (Li, Be, F and N belong to same period of modern periodic table. As we move from left to right in a group, the atomic radii of elements decreases due to high atomic charge but the number of shell remain constant).

(b) Cl < Br < I < At (Cl, At, Br and I belong to same group of periodic table that is group 17. As we move down in a group, the atomic radii of elements increase due to increase in the number of shell of an element).

Elements with 1 to 3 valence electrons are usually metals. Elements with 4 or more valence electrons are usually non-metals or metalloids.

(a) 2, 8, 2 – Magnesium (12)

(b) 2, 8, 1 – Sodium (11)

(d) 2, 1 – Lithium (3)

Element A is K (Potassium).

The electronic configuration of element A (atomic number 19) would be 2, 8, 8, 1. As it has only one valence electron therefore it must be a metal. Thus it is potassium.

Element B is Cl (Chlorine).

The electronic configuration of element B (atomic number 17) would be 2, 8, 7. As it has 7 valence electrons therefore it must be a non-metal. Thus it is chlorine.

A metal and a non-metal usually combine with an ionic bond. Metals have tendency to lose electrons and form cations whereas non-metals can accept electrons to form anions.

Dot structure:

Potassium and chlorine will combine with an ionic bond to form potassium chloride (KCl).

The electron dot structure of KCl is as given below:

Ge < Ga < Mg < Ca < K

The metallic character decreases from left to right in a period. All the above elements, except Mg, are placed in the same period so the metallic character will decrease from K to Ge and order of metallic character would be Ge < Ga < Ca < K

As we move down in a group, metallic nature of elements increases. So, Ca is more metallic than Mg. Similarly, Mg will be less metallic than K because in period, metallic nature decreases from left to right. Hence the increasing order of metallic character would be

Ge < Ga < Mg < Ca < K

The arrangement of some of these elements in the Modern Periodic table is shown as below.

Sodium (Na) or Potassium (K) [Alkali metals such as sodium (Na) and potassium (K) are soft and reactive.] (b) Calcium (Ca) [The chemical formula of limestone is CaCO₃. Hence calcium (Ca) would be an important constituent of limestone.] (c) Mercury (Hg) [Mercury (Hg) is the only metal that exists in liquid state at room temperature.]

Increasing order of their reactivity: Since reactivity of metals decreases from left to right in a period therefore increasing order of reactivity would be Hg < Ca < Na/K.

(a) Sodium (Na) – Group 1 and Period 3 or Potassium (K) – Group 1 and Period 4. Sodium and potassium are alkali metals. Such metals have high reactivity, hence they are usually kept under kerosene. Alkali metals are placed in group 1 and starts from period 2 to 7.

(b) Phosphorus (P) Group 15 and Period 3. It has a variable valency of 3 and 5. It reacts with air and not water. Thus, it is stored in water. It is placed in group 15 and period 3 of periodic table.

(c) Carbon (C) Group 14 and Period 2. Carbon forms the basis of all organic compounds. It is tetravalent because it has 4 valence electrons. It is placed in group 14 and period 2 of Periodic table.

(d) Helium (He) – Group 18 and Period 1. Helium is the lightest inert gas. It has atomic number 2. It is placed in group 18 and period 1 of periodic table.

(e) Aluminium (Al) – Group 13 and Period 3. Aluminium is used to make other elements corrosion resistant by the process of “anodising”. The oxide layer formed is that of Al₂O₃. . The atomic number of aluminum is 13 and it is placed in group 13 and period 3 of periodic table.

(a) The element placed in 2nd Group and 3rd Period of the Periodic Table is Magnesium (Mg).

(b) Electronic configuration: 1s², 2s² 2p⁶ 3s²

K, L, M

2, 8, 2

(c) When the magnesium metal burns in oxygen, it produces magnesium oxide.

2Mg(s) + O₂(g) → 2MgO(s)

When magnesium oxide is heated with water, magnesium hydroxide is formed.

(d) MgO(s) + H₂O(l) → Mg(OH)₂(aq)

Electron dot structure:

Magnesium loses two electrons and oxygen gains two electrons to complete the octet structure.

(a) Element X belongs to Group 17 and 3rd period. X is Cl (chloride).

Element Y belongs to Group 2 and 4^th period. Y is Ca (calcium).

(b) X is a non-metal and Y is a metal.

(c) Metals always combine with oxygen to form basic oxides. Basic oxides react with acids to form salt and water.

In a reaction, meals lose electrons and non-metals gain electrons. Thus positive and negative ions are formed. This creates a strong electrostatic force between these ions. These strong electrostatic forces form the ionic bonds.

(d) Electron dot structure

Electron dot structure for calcium chloride

The elements are neon (Ne), calcium (Ca), nitrogen (N) and silicon (Si)

(b) Group no.: neon (Ne) - 18, calcium (Ca) - 2, nitrogen (N) -15 and silicon (Si) - 14

(c) Period: neon (Ne) – 2nd period, calcium (Ca) – 4th period, nitrogen (N) - 2nd period and silicon (Si) – 3rd period

(d) Electron configuration: neon (Ne) - (2, 8), calcium (Ca) - (2, 8, 8, 2), nitrogen (N) - (2, 5) and silicon (Si) - (2, 8, 4)
(e) Valency: neon (Ne) - 0, calcium (Ca) - 2, nitrogen (N) - 3 and silicon (Si) - 4

(a) H, He, Li, Be, B, C, N, O, F, Ne, Na, Mg, Al, Si, P, S, Cl, Ar, K, Ca

(b) Group 1 — H, Li, Na, K

Group 2 — Be, Mg, Ca

Group 13 — B, Al

Group 14 — C, Si

Group 15 — N, P

Group 16 — O, S

Group 17 — F, Cl

Group 18 — He, Ne, Ar

(a) The two elements that have taken the place of Eka-silicon and Eka-aluminium are Germanium (Ge) and Gallium (Ga), respectively.

(b) Germanium is placed in Group 14 and Period 4 in the Modern Periodic Table.

Gallium is placed in Group 13 and Period 4 in the Modern Periodic Table.

(c) Germanium (Ge) is metalloid. Gallium (Ga) is a metal.

(d) Valence electrons in Germanium (Ge) are 4. Valence electrons in Gallium (Ga) are 3.

(a) The most electropositive element among them is lithium.

(b) The most electronegative element is fluorine.

(c) The element with smallest atomic size is fluorine.

(d) The element which is a metalloid is boron.

(e) The element which shows maximum valency is carbon. Its valency is 4.

(a) Element X is sulphur (atomic no. 16).

(b) Electronic configuration of ‘X’ is

K, L, M

2, 8, 6

(c) 2FeSO₄ (s) Heat → Fe₂O₃(s) + SO₂ (g) + SO₃(g)

(d) SO₂ and SO₃ are acidic oxides.

(e) It belongs to group 16 and 3rd period.

The element X is nitrogen (atomic no. 7) with electronic configuration as 2, 5.

It has 5 valence electrons.

(b) Electron dot structure of diatomic molecule of nitrogen

It has triple covalent bonds.

(c) Electron dot structure of ammonia

It has 3 single covalent bonds.

Noble gases could be placed in Mendeleev's Periodic Table without disturbing the original order.

According to Mendeleev’s classification, the properties of elements are the periodic function of their atomic masses and there is a periodic recurrence of elements with similar physical and chemical properties.

Noble gases were discovered much later after Mendeleev Periodic Table was discovered. Noble gases were placed in a separate group called Zero Group, after the 8th group. The placement of noble gases did not disturb the arrangement of any elements in the Mendeleev’s Periodic Table. Noble gases are present in very low concentration in the atmosphere and are chemically almost unreactive.

In the Mendeleev’s classification of elements, he arranged all 63 elements in the order of their increasing relative atomic masses in the form of a table. It is known as Mendeleev’s Periodic Table. This table was divided into 8 columns and 7 rows. The columns are known as groups and rows are known as periods. Groups from 1 to 7 comprise normal elements and group 8 comprised a few transition elements. Elements with similar properties had been kept in the same group. For example; lithium, potassium, rubidium are kept in 1st group. These elements usually formed compounds such as oxides and hydrides. Such properties formed the basis of classification of these elements. For example, hydrogen, sodium, and potassium belong to the first group. The general formula of oxides for the elements of 1st group is R₂O (H₂O, Na₂O, K₂O).
He arranged all the known elements in increasing order of their atomic masses. Elements with similar properties fall in same group. However, Mendeleev placed many elements in wrong order of their increasing atomic masses. For example, the atomic mass of nickel is less than that of cobalt but still cobalt was placed before nickel.
Mendeleev left some blank spaces intentionally in his periodic table in order to place the elements having similar properties in the same group in future. For example, titanium has been placed in 4th group, leaving a blank space adjacent to it in 3rd group. Similarly, arsenic has been placed in 5th group; leaving two adjacent spaces blank. These spaces have been occupied by scandium, gallium and germanium after their subsequent discoveries.
Mendeleev discovered some elements and named them as eka-boron, eka-aluminium and eka-silicon. Scandium, Gallium and Germanium were discovered later and took the place of eka-boron, eka-aluminium and eka-silicon, respectively in the gap left in the Mendeleev’s Periodic table. Their properties were exactly similar to the corresponding predicted elements.

Thus, Mendeleev’s Periodic Law states that the properties of elements are the periodic function of their relative atomic masses. Noble gases, being inert, could be placed in a separate group without disturbing the original order in the Mendeleev’s Periodic Table.