Grade 10 Chapter 2: Periodic Classification of Elements

Introduction

The periodic table is one of the most significant achievements in the field of chemistry. It organizes all known elements in a systematic way that reveals patterns in their properties. This chapter explores the development of the periodic table, the modern periodic law, and the trends in properties of elements across periods and groups.

Historical Development of the Periodic Table

Early Attempts at Classification

The journey to organize elements began in the early 19th century when scientists noticed patterns in the properties of elements:

1. Döbereiner's Triads (1829): Johann Döbereiner arranged elements with similar properties in groups of three (triads). He observed that the atomic mass of the middle element was approximately the average of the other two elements.

2. Newlands' Law of Octaves (1864): John Newlands arranged elements in order of increasing atomic masses and noted that every eighth element had similar properties, similar to musical octaves.

3. Lothar Meyer's Work (1869): Lothar Meyer independently developed a periodic table based on the atomic volume of elements, showing periodicity in physical properties.

Mendeleev's Periodic Table

In 1869, Russian chemist Dmitri Mendeleev published his periodic table, which became the foundation for the modern periodic table:

1. Arrangement: Mendeleev arranged elements in order of increasing atomic mass.

2. Periodicity: He observed that elements with similar properties occurred at regular intervals.

3. Gaps: Mendeleev left gaps in his table for undiscovered elements and predicted their properties.

4. Corrections: He even placed some elements out of atomic mass order to better match their chemical properties.

Achievements of Mendeleev's Periodic Table

1. Prediction of New Elements: Mendeleev predicted the existence and properties of several undiscovered elements, such as gallium (eka-aluminum), scandium (eka-boron), and germanium (eka-silicon).

2. Correction of Atomic Masses: He suggested that the atomic masses of some elements were incorrect and needed revision.

3. Accommodation of Noble Gases: When noble gases were discovered later, they fit perfectly into Mendeleev's table as a new group.

Modern Periodic Table

Modern Periodic Law

The modern periodic law, proposed by Henry Moseley in 1913, states that:

"The physical and chemical properties of elements are periodic functions of their atomic numbers."

This replaced Mendeleev's law, which was based on atomic mass. Moseley showed that atomic number (the number of protons in the nucleus) was more fundamental than atomic mass for determining an element's position in the periodic table.

Structure of the Modern Periodic Table

The modern periodic table is organized as follows:

1. Periods: Horizontal rows numbered from 1 to 7. Elements in the same period have the same number of electron shells.

2. Groups: Vertical columns numbered from 1 to 18. Elements in the same group have similar chemical properties due to the same number of valence electrons.

3. Blocks: The periodic table is divided into four blocks based on the type of orbital being filled:

4. s-block (Groups 1-2)
5. p-block (Groups 13-18)
6. d-block (Groups 3-12)
7. f-block (Lanthanides and Actinides)

Classification of Elements

1. Metals: Located on the left and center of the periodic table. They are good conductors of heat and electricity, malleable, ductile, and typically have 1-3 valence electrons.

2. Non-metals: Located on the right side of the periodic table. They are poor conductors of heat and electricity, brittle, and typically have 4-8 valence electrons.

3. Metalloids: Elements that show properties of both metals and non-metals. They form a diagonal border between metals and non-metals (B, Si, Ge, As, Sb, Te, Po).

Periodic Trends

Atomic Size (Atomic Radius)

1. Across a Period: Atomic radius decreases from left to right due to increasing nuclear charge and the same number of electron shells.

2. Down a Group: Atomic radius increases from top to bottom due to the addition of new electron shells.

Ionization Energy

Ionization energy is the energy required to remove an electron from a gaseous atom in its ground state.

1. Across a Period: Ionization energy generally increases from left to right due to increasing nuclear charge and decreasing atomic size.

2. Down a Group: Ionization energy decreases from top to bottom due to increasing atomic size and greater shielding effect.

Electron Affinity

Electron affinity is the energy released when a gaseous atom accepts an electron.

1. Across a Period: Electron affinity generally increases from left to right (becomes more negative).

2. Down a Group: Electron affinity generally decreases from top to bottom (becomes less negative).

Electronegativity

Electronegativity is the tendency of an atom to attract shared electrons in a chemical bond.

1. Across a Period: Electronegativity increases from left to right.

2. Down a Group: Electronegativity decreases from top to bottom.

Metallic Character

1. Across a Period: Metallic character decreases from left to right.

2. Down a Group: Metallic character increases from top to bottom.

Properties of Elements Based on Groups

Group 1: Alkali Metals (Li, Na, K, Rb, Cs, Fr)

1. Electronic Configuration: ns¹ (one valence electron)
2. Properties:
3. Soft, silvery metals with low melting and boiling points
4. Highly reactive, stored in oil to prevent reaction with air
5. Form ionic compounds with non-metals
6. React with water to form hydroxides and hydrogen gas
7. Strong reducing agents

Group 2: Alkaline Earth Metals (Be, Mg, Ca, Sr, Ba, Ra)

1. Electronic Configuration: ns² (two valence electrons)
2. Properties:
3. Harder and denser than alkali metals
4. Less reactive than alkali metals
5. Form ionic compounds with non-metals
6. React with water to form hydroxides and hydrogen gas (less vigorously than alkali metals)
7. Strong reducing agents

Group 13-16: p-Block Elements

These groups show a gradual transition from metallic to non-metallic character:

1. Group 13 (B, Al, Ga, In, Tl):
2. Electronic Configuration: ns²np¹
3. Boron is a metalloid; others are metals

4. Show both +1 and +3 oxidation states

5. Group 14 (C, Si, Ge, Sn, Pb):

6. Electronic Configuration: ns²np²
7. Carbon is a non-metal, silicon and germanium are metalloids, tin and lead are metals

8. Show +2 and +4 oxidation states

9. Group 15 (N, P, As, Sb, Bi):

10. Electronic Configuration: ns²np³
11. Nitrogen and phosphorus are non-metals, arsenic and antimony are metalloids, bismuth is a metal

12. Show -3, +3, and +5 oxidation states

13. Group 16 (O, S, Se, Te, Po):

14. Electronic Configuration: ns²np⁴
15. Oxygen and sulfur are non-metals, selenium and tellurium are metalloids, polonium is a metal
16. Show -2, +2, +4, and +6 oxidation states

Group 17: Halogens (F, Cl, Br, I, At)

1. Electronic Configuration: ns²np⁵ (seven valence electrons)
2. Properties:
3. Highly reactive non-metals
4. Exist as diatomic molecules (F₂, Cl₂, Br₂, I₂)
5. Form salts with metals
6. Strong oxidizing agents
7. Reactivity decreases down the group

Group 18: Noble Gases (He, Ne, Ar, Kr, Xe, Rn)

1. Electronic Configuration: ns²np⁶ (eight valence electrons, except helium with 1s²)
2. Properties:
3. Extremely stable and unreactive due to complete outer electron shells
4. Exist as monatomic gases
5. Low melting and boiling points
6. Used in lighting, welding, and as inert atmospheres

Anomalous Properties of Second Period Elements

The first elements of each group (Li, Be, B, C, N, O, F) often show properties that differ from the general trends of their groups. This is due to:

1. Small Size: These elements have much smaller atomic radii.
2. High Electronegativity: They are generally more electronegative than other members of their groups.
3. Absence of d-orbitals: They cannot expand their octet, limiting their bonding patterns.
4. High Charge Density: Their ions have high charge-to-radius ratios.

Examples of anomalous behavior: - Lithium forms covalent compounds more readily than other alkali metals - Beryllium forms covalent compounds, while other alkaline earth metals form ionic compounds - Boron forms electron-deficient compounds (e.g., BF₃) - Carbon has a unique ability to form long chains (catenation)

Significance of the Periodic Table

The periodic table is not just a classification system but a powerful predictive tool:

1. Predicting Properties: The position of an element can help predict its physical and chemical properties.

2. Understanding Bonding: The periodic table helps explain why certain elements form particular types of bonds.

3. Discovering New Elements: The periodic table has guided the search for new elements by predicting their properties.

4. Understanding Reactions: Knowledge of an element's group can help predict its reactions with other elements.

Applications in Modern Chemistry

1. Materials Science: The periodic table guides the development of new materials with specific properties.

2. Medicinal Chemistry: Understanding element properties helps in designing drugs and medical treatments.

3. Environmental Chemistry: Knowledge of element behavior helps in addressing pollution and environmental issues.

4. Nuclear Chemistry: The periodic table is essential for understanding radioactive decay and nuclear reactions.

Summary

The periodic table is a masterpiece of scientific classification that organizes elements based on their atomic numbers and reveals patterns in their properties. It has evolved from early attempts at classification to the modern periodic table we use today. Understanding the periodic trends and group properties helps predict the behavior of elements and their compounds, making the periodic table an indispensable tool in chemistry.

Practice Questions

1. Explain the difference between Mendeleev's periodic law and the modern periodic law.
2. How does atomic radius change across a period and down a group? Explain the reasons for these trends.
3. Why do elements in the same group have similar chemical properties?
4. Compare and contrast the properties of alkali metals and halogens.
5. Explain why the first element of each group often shows anomalous behavior compared to other elements in the same group.
6. How does electronegativity vary across the periodic table, and why?
7. Describe the relationship between an element's position in the periodic table and its electron configuration.
8. How did Mendeleev's periodic table predict the properties of undiscovered elements?
9. Explain the significance of the s, p, d, and f blocks in the periodic table.
10. How has the periodic table contributed to the development of modern chemistry?

References: 1. Maharashtra State Board 10th Standard Science Syllabus 2025-26 2. NCERT Science Textbook for Class 10 3. Royal Society of Chemistry - Periodic Table Resources 4. American Chemical Society - Periodic Table Information