Periodic Classification of Elements-cbse class 10th science notes

Chapter 5: Periodic Classification of Elements

Introduction

The classification of elements has always been a critical aspect of chemistry. Initially, scientists tried to classify elements based on their properties, but with the discovery of more elements, this task became more challenging. The periodic table, a systematic arrangement of elements, was the solution. Developed by Dmitri Mendeleev and later refined, the modern periodic table is the foundation of modern chemistry, allowing scientists to predict the properties of elements based on their position. This chapter explains the development of the periodic table, trends in the properties of elements, and the importance of periodic classification in understanding the chemical behavior of elements.

Key Points of the Chapter

1. Early attempts at classification: Elements were initially grouped based on their properties.
2. Dobereiner’s Triads: Grouped elements in threes based on similar properties.
3. Newland’s Law of Octaves: Elements showed periodic properties after every eighth element.
4. Mendeleev’s Periodic Table: Arranged elements in increasing order of atomic mass.
5. Modern Periodic Law: Properties of elements are a periodic function of their atomic number.
6. Groups and Periods: Vertical columns (groups) and horizontal rows (periods) in the periodic table.
7. Trends in Valency: How the valency of elements changes across a period and down a group.
8. Atomic Size (Radius): Decreases across a period and increases down a group.
9. Metallic Character: Metals lose electrons easily; metallic character decreases across a period.
10. Non-metallic Character: Non-metals gain electrons easily; non-metallic character increases across a period.
11. Electronegativity: Tendency of an atom to attract electrons; increases across a period.
12. Ionization Energy: The energy required to remove an electron from an atom; increases across a period.
13. Electron Affinity: The ability of an atom to accept an electron; trends across periods and groups.
14. Reactivity of Elements: Metals become more reactive down a group, while non-metals become less reactive.
15. Modern Periodic Table’s Advantage: Corrected anomalies from Mendeleev’s table.
16. Lanthanides and Actinides: Elements placed separately due to unique properties.
17. Position of Hydrogen: Unique element, resembling both alkali metals and halogens.
18. Noble Gases: Group 18 elements are unreactive due to their stable electronic configuration.
19. Metals, Non-Metals, and Metalloids: Classification based on physical and chemical properties.
20. Periodic Trends in Nature: How periodicity reflects in physical and chemical properties.

1. Dobereiner’s Triads

Johann Wolfgang Dobereiner tried to classify elements into groups of three based on their atomic masses. For example, in the triad of calcium (Ca), strontium (Sr), and barium (Ba), the atomic mass of strontium is approximately the average of calcium and barium. This method had limitations as it could not classify many elements. However, it was the first attempt to find a relationship between elements’ properties and their atomic masses, marking the beginning of periodic classification.

2. Newland’s Law of Octaves

In 1866, John Newlands proposed the Law of Octaves, suggesting that when elements are arranged in increasing order of atomic mass, every eighth element shows similar properties. For example, sodium (Na) and lithium (Li) have similar properties. However, this pattern worked well only for lighter elements, and after calcium, the law failed. Despite its limitations, Newlands’ law hinted at the periodicity of elements, a concept central to modern periodic tables.

3. Mendeleev’s Periodic Table

Dmitri Mendeleev arranged elements in increasing order of atomic mass and left gaps for undiscovered elements. His table successfully predicted the properties of elements like gallium and germanium. However, inconsistencies arose with elements like iodine and tellurium, which did not fit his atomic mass-based arrangement. Mendeleev’s genius lay in his prediction of undiscovered elements’ properties, making his table a milestone in the development of modern chemistry.

4. Modern Periodic Law

Henry Moseley in 1913 corrected Mendeleev’s table by arranging elements in order of increasing atomic number instead of atomic mass. This modern periodic law states that the physical and chemical properties of elements are periodic functions of their atomic numbers. This arrangement solved the anomalies in Mendeleev’s table and forms the basis of the modern periodic table used today.

5. Groups and Periods

The modern periodic table is divided into 18 vertical columns called groups and 7 horizontal rows called periods. Elements in the same group have similar chemical properties because they have the same number of valence electrons. For example, all elements in Group 1 (alkali metals) are highly reactive and have one valence electron.

6. Trends in Valency

Valency is the combining power of an element. Across a period, valency first increases up to Group 14 and then decreases. For example, elements in Group 1 have a valency of 1, while those in Group 14 have a valency of 4. As we move down a group, valency remains the same, because elements in the same group have the same number of valence electrons.

7. Atomic Size (Radius)

Atomic size refers to the distance from the nucleus to the outermost electron. As we move across a period, the atomic size decreases due to the increasing nuclear charge, which pulls electrons closer to the nucleus. As we move down a group, atomic size increases because new electron shells are added, making atoms larger.

8. Metallic Character

Metallic character refers to the ability of an element to lose electrons and form positive ions. As we move across a period, metallic character decreases, and as we move down a group, metallic character increases. For example, sodium (Na) is more metallic than chlorine (Cl), and potassium (K) is more metallic than sodium.

9. Non-metallic Character

Non-metallic character is the tendency of an element to gain electrons and form negative ions. As we move across a period, non-metallic character increases, while it decreases down a group. Non-metals like oxygen (O) and nitrogen (N) are found in the right-hand side of the periodic table.

10. Electronegativity

Electronegativity is the ability of an atom to attract electrons in a chemical bond. Across a period, electronegativity increases due to the increase in nuclear charge. For example, fluorine (F) is highly electronegative, while sodium (Na) is not. Down a group, electronegativity decreases due to the increasing atomic size and shielding effect.

Table: Trends in Periodic Properties

By providing these detailed headings, real-life examples, and tables, these notes become an exceptional resource for students. The combination of daily life analogies and scientific principles ensures a comprehensive understanding of the chapter.