Periodic Classification of Elements Class 10
Periodic Classification of Elements Class 10
| With the discovery of a large number of elements, it became difficult to study the elements individually, so the classification of elements was done to make the study easier. |
Earlier Classification:
Dobereiner’s Triads: In the year 1817, Johann Wolfgang Döbereiner, a German chemist, tried to arrange the elements with similar properties into groups.
In triads, the atomic mass of the middle element is approximately the average of the other two elements. This is known as the Law of Triads. This classification was applicable to very few elements and so it was rejected. For example:
New Lands Law of Octaves: The attempts of Döbereiner encouraged other chemists to correlate the properties of elements with their atomic masses. In 1866, John Newlands, an English scientist, arranged the known elements in the order of increasing atomic masses and showed that the properties of every eight elements were similar to those of the first one. The relationship is just like the resemblance of first and eighth musical notes. He named this the Law of Octaves. But his classification was rejected since the law of octaves was applicable to elements up to calcium. The properties of the eighth element become not similar for the rest of the element.
Limitations:
(1) Law of octaves was applicable only upto calcium (only for lighter elements).
(2) Newland adjusted two elements in the same slot (e.g. Co and Ni), having different properties. For example; Co and Ni with Fluorine, Chlorine, Bromine and Iodine.
(3) According to Newland, only 56 elements existed in nature and no more elements would be discovered in future.
Periodic Classification of Elements Class 10
MAKING ORDER OUT OF CHAOS – Mendeleev’s Periodic Law: Dimitri Mendeleev classified the elements in the increasing order of their atomic weights. He founded that the properties of elements repeat after a regular interval. Based on this observation, he proposed a Periodic Law which states that “The properties of elements are the periodic functions of their atomic weights.” That is, when elements are arranged in the increasing order of their atomic weights, their properties repeat after a regular interval.
Mendeleev arranged elements in horizontal rows called Periods and vertical Columns called Groups. When Mendeleev proposed his periodic table, some of the elements were not discovered. He left some vacant places (gaps) for them in the periodic table and predicted some of their properties. For e.g. both Gallium and Germanium were not discovered at that time. He named these elements as Eka-Aluminium and Eka-Silicon respectively and predicted their properties.
Merits of Mendeleev’s periodic table:
(i) It was the first comprehensive classification of elements.
(ii) He corrected the wrong atomic weights of some elements and placed them in the correct position in the periodic table.
(iii) He left vacant places for undiscovered elements and predicted some of their properties.
(iv) Elements with similar properties are placed in the same group.
Drawbacks of Mendeleev’s periodic table:
(i) Elements with dissimilar properties are found in the same group.
(ii) He could not give an exact position for hydrogen.
(iii) He could not give an exact position for Lanthanoids and Actinoids and also for isotopes.
(iv) Mendeleev’s periodic table did not strictly obey the increasing order of atomic weights.
Periodic Classification of Elements Class 10
Modern Periodic Law: Moseley performed experiments and studied the frequencies of the x-rays emitted from the elements. With these experiments, he concluded that the atomic number is a more fundamental property of an element than its atomic mass.
Based on this observation, he modified Mendeleev’s periodic law as “the physical and chemical properties of elements are the periodic functions of their atomic numbers”. This is known as Modern Periodic law.
Long-form of Periodic Table (Bohr’s table):
(i) It is divided into two categories
Vertical columns – Groups
Horizontal rows – Periods
(ii) There are 18 groups
(iii) There are 7 periods
There are 7 periods in the Modern periodic table. The period number corresponds to the highest principal quantum number of the elements.
The first period contains 2 elements (H and He). Here the subshell filled is 1s. This period is called a very short period.
The second period contains 8 elements (Li to Ne). Here the subshells filled are 2s and 2p. The third period also contains 8 elements (Na to Ar). Here the subshells filled are 3s and 3p. These 2 periods are called short periods.
The fourth period contains 18 elements (K to Kr). Here the subshells filled are 4s, 3d, and 4p. The fifth period also contains 18 elements (Rb to Xe). Here the subshells filled are 5s, 4d, and 5p. These 2 periods are called long periods.
The sixth period contains 32 elements (Cs to Rn). Here the subshells filled are 6s, 4f, 5d, and 6p. This period is the longest period in the periodic table and is called the Monster period.
The seventh period is an incomplete period. It can also accommodate 32 elements. Here the subshells filled are 7s, 5f, 6d, and 7p.
The 14 elements of each of the sixth and seventh periods are placed in separate rows below the main body of the periodic table. These are together called inner transition elements. The 14 elements of the sixth period (from cerium to lutetium) are called Lanthanides or Lanthanones or Lanthanoids or rare-earth metals. The 14 elements of the seventh period (from thorium to lawrencium) are called Actinides or Actinones or Actinoids.
Periodic Classification of Elements Class 10
Atomic Properties with Pictorial Representation
Atomic Properties: The properties such as atomic radius, ionic radius, ionisation energy, electronegativity, electron affinity, and valence, called atomic properties.
Atomic Radius: It is defined as the distance from the centre of the nucleus to the outermost shell containing electrons. Atomic radius is commonly expressed in picometre (pm) or angstrom (Å).it is measured by the x-ray diffraction method or by spectroscopic methods.
Covalent radius: One-half of the distance between the nuclei of two covalently bonded atoms of the same element in a molecule (used for non-metals).
Van der Waals radius: One half of the distance between nuclei of two identical atoms of separate molecules called van der Waals radii.
Metallic radius: One-half of the internuclear distance between the two adjacent metal ions in the metallic lattice.
Van der waal’s radius > metallic radius > covalent radius
Periodic Classification of Elements Class 10
Variation of Atomic Radii Along the Period and Down the Group: The atomic size decreases from left to right in a period. This is because, in a period, the electrons are added to the same valence shell. Thus, the number of shells remains the same, but the effective nuclear charge increases. So, the atomic radius decreases. In a given period, alkali metals (group 1) have the maximum size and halogens (group 17) have the minimum size.
Example: Size of second period elements: Li > Be > B > C > N > O > F
Down a group, the atomic radius increases from top to bottom. This is because of the increase in no. of shells and the Shielding Effect (in atoms with higher atomic numbers, the inner electrons partially shield the attractive force of the nucleus. So the outer electrons do not experience the full attraction of the nucleus and this is known as the shielding effect or screening effect).
Example ; Atomic size of first group element : Li < Na < K < Rb < Cs < Fr
Atomic radius of noble gases is larger than that of halogens. This is because noble gases are monoatomic. So van der Waal’s radius is used to express the atomic radius which is greater than covalent radius or metallic radius.
Periodic Classification of Elements Class 10
Ionic Radius: Effective distance from the centre of the nucleus of the ion up to which it exerts its, influence on its electronic cloud.
Generally, a cation is smaller than its parent atom (e.g. Na+ is smaller than Na atom). This is because a cation has fewer electrons, but its nuclear charge remains the same as that of the parent atom. An anion is larger than its parent atom (e.g. Cl– is larger than Cl atom). This is because the addition of one or more electrons would result in an increased electronic repulsion and a decrease in effective nuclear charge.
The size of inert gases is larger than halogens because in inert gases all orbitals are completely filled, inter electronic repulsions are maximum.
Electronegativity Ionisation and Electron Gain Enthalpies
Electronegativity Ionisation and Electron Gain Enthalpies:
Ionization Enthalpy (∆iH): Minimum amount of energy required to remove the most loosely bound valence electron from an isolated gaseous atom so as to convert it into gaseous cation.
It may be represented as: X(g) + ∆iH → X+(g) + e–
Its unit is kJ/mol or J/mol.
The energy required to remove the first electron from the outermost shell of a neutral atom is called first ionization enthalpy (∆iH1)
X(g) + ∆iH1 → X+(g) + e–
Second Ionisation enthalpy (∆iH2) is the amount of energy required to remove an electron from a unipositive ion.
X+(g) + ∆iH2 → X2+(g) + e–
Periodic Classification of Elements Class 10
Energy is always required to remove an electron from an atom or ion. So ∆iH is always positive. The second ionization enthalpy is always higher than the first ionization enthalpy. This is because it is more difficult to remove an electron from a positively charged ion than from a neutral atom.
Similarly, the third ionization enthalpy is higher than the second ionization enthalpy, and so on. i.e. ∆iH1 < ∆iH2 < ∆iH3………… As the ease of removal of electrons increases, the ionization enthalpy decreases.
Factor affecting Ionization enthalpy: The important factors which affect ionization enthalpy are:
a) Atomic size: Greater the atomic size (atomic radius), the smaller will be the ionization enthalpy.
b) Nuclear charge: The value of ionization enthalpy increases with nuclear charge.
Electron Gain Enthalpy (∆egH): It is the enthalpy change when an electron is added to an isolated gaseous atom. X(g) + e– → X– (g). Its unit is kJ/mol.
The greater the amount of energy released, the higher is the electron gain enthalpy of the element.
It may be positive or negative depending on the nature of the element. For most of the elements, energy is released when an electron is added to its atoms. So ∆egH is negative. Noble gases have large positive electron gain enthalpy because of their completely filled (stable) electronic configuration.
Variation along the Period: From left to right across a period, ∆egH becomes more negative. This is because of a decrease in atomic radius and an increase in nuclear charge. So, the ease of addition of electrons increases and hence the ∆egH.
Variation down the Group: As the size increases, the tendency to add the electron decreases hence electron gain enthalpy becomes less – ve.
Electron gain enthalpy of fluorine is less negative than chlorine. This is because, when an electron is added to F, it enters into the smaller 2nd shell. Due to the smaller size, the electron suffers more repulsion from the other electrons. But for Cl, the incoming electron goes to the larger 3rd shell. So, the electronic repulsion is low and hence Cl adds electron more easily than F.
Thus, in the modern periodic table, alkali metals have the least –ve ∆egH, and halogens have the most –ve ∆egH. Among halogens, the negative ∆egH decreases as follows. Cl> F > Br > I. The negative electron gain enthalpy is also called electron affinity.
Noble gases have positive electron gain enthalpy. They have completely filled orbitals. Additional electrons will be placed in the next higher shell. As a result, energy has to be supplied to add an additional electron.
The formation of O is exothermic but O is endothermic.
O (g) + e– → O– (g); ∆egH = – 141 KJ mol
O– (g) + e– → O2- (g); ∆egH = + 780 KJ mol
∆egH After the addition of one electron the atom becomes negatively charged and the second electron is to be added to a negatively charged ion. The addition of a second electron is opposed by electrostatic repulsion, hence energy has to be supplied for the addition of a second electron to overcome the strong electrostatic repulsion between the negatively charged O ion and the second electron is added.
Electronegativity: The tendency of an atom in a molecule to attract the shared pair of electrons towards itself is termed as its electronegativity. It is not a measurable quantity and so it has no unit. There are different scales for measuring the Electronegativity of elements. The most commonly used is the Pauling Electronegativity scale developed by Linus Pauling.
(i) In period- The electro-negativity increases from left to right in a period.
(ii) In group- The electro-negativity decreases from top to bottom in a group.
Difference between Electronegativity and Electron Gain Enthalpy:
Periodic Classification of Elements Class 10
Diagonal Relationship: Some elements of the second-period show similarities with elements of the third period placed diagonally to each other due to the same charge/radius ratio.
Anomalous Behaviour of the first element of every group is due to
i). Small size and high electronegativity: N can form p π – p π multiple bonds whereas P cannot.
ii) High IE: They form only covalent compounds and not ionic compounds.
iii) Absence of vacant d orbitals: N cannot form NCl5 or R3N = O since it cannot expand its covalence beyond 4 whereas P can form PCl5 and R3P = O.
Valence Electrons: The electrons present in the outermost shell are called valence electrons. Because the electrons in the outermost shell determine the valency of an element.
Valency of an Element: According to the electronic concept of valency, “the number of electrons which an atom loses or gains or shares with another atom to attain the noble gas configuration is termed as its valency.”
Periodicity:
(i) In period- The valency first increases then decreases from left to right in a period.
(ii) In group- The valency remains constant from top to bottom in a group.
Electropositivity: It is the tendency of an atom to lose the most loosely bound electron (valence electron). It is directly related to the metallic character of elements. It depends on the atomic size and nuclear charge. As the atomic radius increases, electropositivity increases.
Along a period, electropositivity decreases from left to right. But down a group, it increases. So, francium is the most electropositive element and fluorine is the least electropositive element.
Variation of Various Properties along with the Periods and Groups
Periodic Classification of Elements Class 10
