Classification Of Elements- The Periodic Table (LAQs)

Physical Science | 7. Classification Of Elements- The Periodic Table – LAQs:
Welcome to LAQs in Chapter 7: Classification Of Elements- The Periodic Table. This page contains the most Important FAQs for Long Answer Questions in this Chapter. Each answer is provided in simple English, with a Telugu explanation, and formatted according to the exam style. This will support your preparation and help you secure top marks in your exams.

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LAQ-1 : Explain how does the atomic size and ionization energy changes in groups and periods in the Modern periodic table.

Introduction

The modern periodic table exhibits distinct patterns in atomic size and ionization energy as one moves across periods and down groups. Understanding these trends is key to grasping the chemical properties and reactivity of elements.

Change in Atomic Size

1. Within Groups (Top to Bottom):
   • Trend: Increase in atomic size.
   • Reason: Addition of electron shells increases the distance between the nucleus and the outermost electron.
2. Across Periods (Left to Right):
   • Trend: Decrease in atomic size.
   • Reason: Increase in nuclear charge pulls electrons closer to the nucleus, reducing the atomic size.

Change in Ionization Energy

1. Within Groups (Top to Bottom):
   • Trend: Decrease in ionization energy.
   • Reason: Larger atomic size makes it easier to remove an electron as the outermost electron is farther from the nucleus.
2. Across Periods (Left to Right):
   • Trend: Increase in ionization energy.
   • Reason: Smaller atomic size and higher nuclear charge mean electrons are more tightly bound to the nucleus, requiring more energy to remove.

Summary

In summary, atomic size increases down a group and decreases across a period, while ionization energy decreases down a group and increases across a period. These trends are influenced by the addition of electron shells and the increase in nuclear charge, which are fundamental in predicting the behavior and reactivity of elements in chemical reactions.

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LAQ-2 : Define the modern  periodic law. Discuss the construction of the long form of the periodic table.

Introduction

The Modern Periodic Law states that the physical and chemical properties of elements are periodic functions of their atomic numbers, not atomic weights. This principle, proposed by Moseley, represents a modification to Mendeleev’s original periodic law.

Construction of the Long Form of the Periodic Table

The Long Form of the Periodic Table organizes all known elements based on the Modern Periodic Law.

Periods

1. Consists of seven horizontal rows called periods, numbered from 1 to 7.
2. The period number corresponds to the principal quantum number (n).
3. The number of elements in each period varies, increasing with the period number.

Groups

1. Features 18 vertical columns known as groups, numbered from 1 to 18.
2. Elements within a group exhibit similar chemical properties and reactivity.
3. Groups are further divided into blocks (s, p, d, and f) based on the last electron’s orbital type.

Blocks

1. s-block: Groups 1 and 2, excluding Hydrogen in Group 1, and Helium.
2. p-block: Groups 13 to 18.
3. d-block: Transition metals, Groups 3 to 12.
4. f-block: Lanthanoids and Actinoids, placed separately for compactness.

Element Placement

1. Each element is positioned according to its atomic number, electron configuration, and chemical properties.
2. Periods fill sequentially, reflecting the filling of s, p, d, and f orbitals.

Special Groups

1. Lanthanoids: Elements 57 (La) to 71 (Lu) in the f-block.
2. Actinoids: Elements 89 (Ac) to 103 (Lr) in the f-block.

Summary

The modern periodic table’s structure, with its division into groups, periods, and blocks, offers a comprehensive representation of all known elements. This arrangement effectively correlates elements’ atomic numbers with their electron configurations, providing valuable insights into their physical and chemical properties. It facilitates a systematic and logical approach to studying and applying elements in various scientific fields.

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LAQ-3 : Write down the characteristics of the elements having atomic number 17.

Electronic configuration——-Period number——

Group number——                        Element family——

No. of valency electrons—–Valency——-

Metal or non-metal—–

Characteristics of the Element with Atomic Number 17 (Chlorine, Cl)

1. Electronic Configuration
   • Configuration:
     $$1s^2 2s^2 2p^6 3s^2 3p^5$$
   • Indicates 17 electrons, filling the energy levels according to the Aufbau principle.
2. Period Number
   • Period: 3
   • Chlorine’s electron configuration places it in the third period of the periodic table.
3. Group Number
   • Group: 17 (VIIA in older notation)
   • The five electrons in the 3p orbital signify placement in Group 17.
4. Element Family
   • Family: Halogen
   • As a member of Group 17, chlorine is classified in the Halogen family, known for its high reactivity and ability to gain one electron for stability.
5. Number of Valency Electrons
   • Valence Electrons: 7
   • The 3s²3p⁵ configuration shows seven electrons in the outermost shell.
6. Valency
   • Valency: 1
   • Chlorine typically gains one electron to achieve a stable noble gas configuration (3s²3p⁶), giving it a valency of 1.
7. Metal or Non-Metal
   • Type: Non-metal
   • Exhibits non-metallic properties like high electronegativity and forming covalent bonds by gaining electrons.

Summary

In summary, chlorine, with an atomic number of 17, is a highly reactive non-metal located in Group 17 (Halogen Family) and Period 3 of the periodic table. It has a valency of 1 and seven valence electrons, aligning with its characteristic properties as a halogen.

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LAQ-4 : Observe the information and answer the following questions.

Name of the Element	Atomic Number	Electronic Configuration
Sodium	11	[Ne] 3s1
Magnesium	12	[Ne] 3s2
Potassium	19	[Ar] 4s1
Calcium	20	[Ar] 4s2

1.What is valency of Magnesium?

2.Which element has more electropositivity?

3.Write the elements which belongs to 3rd period.

4.Write the elements which belongs to 1st group.

Introduction

This guide focuses on topics like valency and electropositivity, specifically for elements such as Magnesium and Potassium, and their classification in the periodic table. The language used is simple to assist students and non-native English speakers in understanding these chemical concepts.

1. Valency of Magnesium:
   • Valency: 2, indicating Magnesium can lose two electrons during chemical reactions.
2. Element with More Electropositivity:
   • Potassium (K) is more electropositive, tending to lose electrons more easily.
3. Elements in the 3rd Period:
   • 3rd Period Elements: Includes Sodium (Na) and Magnesium (Mg), extending from Sodium to Argon.
4. Elements in the 1st Group:
   • 1st Group Elements (Alkali Metals): Includes Sodium (Na) and Potassium (K).

A Deeper Look

1. Valency Explanation: Reflects an atom’s capability to form bonds. Magnesium’s valency of 2 means it can form bonds by losing two electrons.
2. Electropositivity Concept: Describes an atom’s ability to donate electrons. Potassium is highly electropositive, readily giving away electrons.
3. Periods in the Periodic Table: Periods are the rows of the table, with the 3rd period comprising elements from Sodium to Argon.
4. Groups in the Periodic Table: Groups are columns in the table. The 1st group includes alkali metals like Sodium and Potassium.

Summary

1. The valency of Magnesium (2).
2. Potassium’s higher electropositivity.
3. The placement of Sodium and Magnesium in the 3rd period.
4. The inclusion of Sodium and Potassium in the 1st group (alkali metals). These points are crucial for exam preparation and enhancing understanding of chemical properties and periodic table organization.

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LAQ-5 : Why Mendeleeff had to leave certain blank spaces in his periodic table? What is your explanation for this?

Introduction

This discussion focuses on Dmitri Mendeleev, the creator of the Periodic Table, and his rationale for leaving blank spaces in the original table. We will delve into his emphasis on the chemical properties of elements and his foresight regarding undiscovered elements.

Why Mendeleev Left Blank Spaces

1. Emphasis on Chemical Properties:
   Mendeleev organized elements with similar properties vertically, ensuring that elements with shared characteristics were grouped together.
2. Prediction of Missing Elements:
   Mendeleev noticed gaps where elements, according to the patterns in the table, had not yet been discovered and predicted their existence, leaving blank spaces for them.
3. Discovery of New Elements:
   New elements discovered post-publication fit perfectly into the blank spaces, validating Mendeleev’s predictions and the accuracy of his table.

Explanation and Further Analysis

1. Mendeleev’s method showed remarkable foresight and understanding of elemental properties and trends.
2. The discovery of new elements that fit the blank spaces underscored the accuracy of his Periodic Table, contributing to its widespread acceptance.
3. His approach ensured that elements with similar characteristics and behaviors were grouped together, aiding the study of elements and their interactions.

Summary

In conclusion, Mendeleev’s strategic choice to leave blank spaces in his Periodic Table was based on his focus on chemical properties and his predictions of undiscovered elements. The subsequent discoveries that aligned with his predictions highlighted the precision of his work and established the Periodic Table as an essential tool in chemistry. Understanding this aspect of the table’s history enhances comprehension of element organization and classification.

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LAQ-6 : Explain how the elements are classified into s, p, d and f block elements in the periodic table and give the advantage of this kind of classification.

Introduction

This discussion explores the classification of elements in the periodic table into s, p, d, and f block elements based on their valence shell electronic configuration. We’ll examine the characteristics of each block and the advantages of this classification system.

Classification of Elements

1. s-block Elements:
   • Configuration: ns¹ or ns²
   • Groups: IA and IIA
   • Characteristics: Differentiating electron enters the s-orbital.
2. p-block Elements:
   • Configuration: ns² np^1-6 (except Helium)
   • Groups: IIIA to VIIIA
   • Characteristics: Differentiating electron enters the p-orbital.
   • Also known as representative elements.
3. d-block Elements:
   • Configuration: ns_² np (n-1)d1 to ns² np (n-1)d10
   • Groups: IB to VIIIB
   • Characteristics: Differentiating electron enters the (n-1) d orbital.
   • Also known as transition elements.
4. f-block Elements:
   • Characteristics: Differentiating electron enters the (n-2)f orbital.
   • Contains Lanthanoids and Actinoids.
   • Also known as inner transition elements.

Advantages of Block Classification

1. This classification is an efficient method to organize elements by similar properties and electronic configurations.
2. It aids in easy and systematic study of elements.
3. Helps in understanding the behavior and characteristics of elements.
4. Useful in identifying groups and periods of elements, simplifying the study of elements.

Explanation and Further Analysis

1. The block classification system is based on the electronic configuration of elements.
2. It groups elements with similar properties and behavior, facilitating study and understanding.
3. This classification is crucial in the study of chemistry, offering a structured and organized way to explore elements.

Summary

In summary, elements are classified into s, p, d, and f blocks in the periodic table based on their valence shell electronic configuration. Each block contains specific groups of elements with unique characteristics. This classification system provides clear, organized, and systematic insights into studying elements, understanding their properties, and exploring their interactions. It is a fundamental aspect of modern chemistry, offering a structured approach to elemental exploration.

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LAQ-7 : What is a periodic property? How do the following properties change in a group and period? Explain?

i.a) Atomic radius b) Ionization energy c) Electron affinity d) Electronegativity

ii.Explain the ionization order in the following sets of elements.

a) Na, Al, Cl b) Li, Be, B c) C, N, O d) F, Ne, Na e) Be, Mg, Ca

Introduction

This discussion focuses on periodic properties of elements and examines how properties like atomic radius, ionization energy, electron affinity, and electronegativity change across groups and periods. Additionally, it explores the ionization energy order for certain sets of elements.

What is a Periodic Property?

1. Periodic Property refers to the recurring trends in properties of elements as you move through the periodic table.
2. These changes are consistent due to the arrangement of electrons in atoms.

Changes in Periodic Properties

1. Atomic Radius:
   • Increases from top to bottom within a group (addition of electron shells).
   • Decreases from left to right across a period (electrons pulled closer to the nucleus).
2. Ionization Energy:
   • Decreases from top to bottom within a group (increased atomic size).
   • Increases from left to right across a period (increased nuclear charge).
3. Electron Affinity:
   • Decreases from top to bottom within a group (increased atomic size).
   • Increases from left to right across a period (decreased atomic size).
4. Electronegativity:
   • Decreases from top to bottom within a group.
   • Increases from left to right across a period.

Ionization Energy Order in Sets of Elements

1. Na, Al, Cl:
   • Order: Cl > Al > Na.
   • Reason: Increasing ionization energy from left to right in a period.
2. Li, Be, B:
   • Order: Be > Li > B.
   • Reason: 2p electrons are easier to remove compared to 2s.
3. C, N, O:
   • Order: N > C > O.
   • Reason: Nitrogen’s half-filled degenerate orbital configuration.
4. F, Ne, Na:
   • Order: Ne > F > Na.
   • Reason: Neon’s high ionization energy as an inert gas.
5. Be, Mg, Ca:
   • Order: Be > Mg > Ca.
   • Reason: Decrease in ionization energy with increasing atomic size.

Summary

In conclusion, periodic properties like atomic radius, ionization energy, electron affinity, and electronegativity exhibit specific trends across the periodic table. Understanding these trends and ionization energy orders provides crucial insights into element behavior and reactivity. This knowledge is integral to predicting element interactions and enhances the study and application of chemistry.

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LAQ-8 : Write the factors that influence ionization energy and explain any three of them.

Introduction

Ionization energy is the energy required to remove an electron from a gaseous atom or ion. Various factors affect ionization energy, which can provide insights into elemental properties and reactivities.

Factors Influencing Ionization Energy

1. Nuclear Charge:
   • Involves the charge within the atom’s nucleus.
   • A greater nuclear charge typically means higher ionization energy.
2. Screening or Shielding Effect:
   • Refers to the reduction in effective nuclear charge on the electron cloud, influenced by the attraction forces between electrons and the nucleus.
   • A higher screening effect often results in lower ionization energy.
3. Penetration Power of Orbitals:
   • Different orbitals have varying abilities to approach the nucleus within the same main shell.
   • Orbitals with higher penetration power usually exhibit higher ionization energy.

Explanation of Selected Factors

1. Nuclear Charge:
   • A higher nuclear charge results in a stronger attraction to electrons, increasing ionization energy.
   • Example: Chlorine has a higher nuclear charge than Sodium, leading to more ionization energy.
2. Screening or Shielding Effect:
   • Inner electrons repel outer electrons, reducing the nucleus’s effective charge on outer electrons and lowering ionization energy.
   • Example: Cesium has a higher shielding effect compared to Lithium, resulting in lower ionization energy.
3. Penetration Power of Orbitals:
   • Orbitals closer to the nucleus are more strongly attracted, leading to higher ionization energy.
   • Example: Beryllium’s 2s electrons have more penetration power compared to Boron’s 2p electrons, resulting in higher ionization energy for Beryllium.

Summary

In conclusion, factors such as nuclear charge, screening or shielding effect, and penetration power of orbitals significantly influence an element’s ionization energy. Understanding these factors enriches our knowledge of atomic structure and behavior, impacting how elements interact and react in various chemical contexts.