NCERT Science Notes - Class 9
Chapter 4 - Structure of Atom

Welcome to AJs Chalo Seekhen. This webpage is dedicated to Class 9 | Science | Chapter 4 - Structure of Atom. In this chapter, students dive into the intricacies of atomic structure. They learn about the discovery of subatomic particles—electrons, protons, and neutrons. Key concepts include Rutherford’s model, Bohr’s model of the atom, and the arrangement of electrons in different energy levels or shells. The chapter also covers the concept of valency, isotopes, and isobars. Understanding these fundamental ideas is crucial for grasping how atoms combine to form molecules and engage in chemical reactions. This knowledge forms the cornerstone of modern chemistry, helping students appreciate the behavior of matter at an atomic level. 🧬📘

NCERT Science Notes - Class 9 | Science | Chapter 4 - Structure of Atomnnotes ajs, cbse notes class 10 ajslearning, ajs notes class 10, ajslearning, ajs chalo seekhen

NCERT Science Notes - Class 9
Chapter 4 - Structure of Atom

Fundamental Concepts

• Atoms and Molecules:
  • Atoms are the smallest units of matter, while molecules are formed by two or more atoms bonded together.
  • The diversity of matter arises from the different types of atoms that combine to form molecules.

1. Differences Between Atoms of Different Elements:
   • Atoms of different elements differ in the number of protons, neutrons, and electrons they contain.
   • This difference is what defines the element and its unique properties.
2. Indivisibility of Atoms:
   • Historical View: Dalton proposed that atoms are indivisible and the basic building blocks of matter.
   • Modern Understanding: Atoms are not indivisible; they are composed of smaller particles.

Sub-Atomic Particles

• Definition: Particles that make up an atom.
  • Protons: Positively charged particles found in the nucleus.
  • Neutrons: Neutral particles also located in the nucleus.
  • Electrons: Negatively charged particles that orbit the nucleus.

• At the end of the 19th century, scientists were focused on uncovering the structure of the atom and understanding its properties.
• A major breakthrough came from studying static electricity and the conditions under which different substances conduct electricity.

• The elucidation of atomic structure is based on a series of experiments that revealed:
  • The existence of sub-atomic particles.
  • The arrangement of these particles within the atom.

4.1 - Charged Particles in Matter

Key Concepts

• Charged Particles: Atoms consist of charged particles that play a crucial role in the behavior of matter.
• Electric Charge: A property of particles that causes them to experience a force when placed in an electromagnetic field.

• Activity 4.1:
  • A. Comb Dry Hair:
    • Observation: After combing, the comb attracts small pieces of paper.
    • Conclusion: The comb becomes electrically charged, indicating that rubbing can transfer charge.
  • B. Rub a Glass Rod with a Silk Cloth:
    • Observation: The rubbed glass rod attracts an inflated balloon when brought close.
    • Conclusion: This further demonstrates that rubbing two objects together generates an electric charge.
      
      

1. Can we conclude that on rubbing two objects together, they become electrically charged?
   • Answer: Yes, rubbing two objects together results in the transfer of charge, causing both objects to become electrically charged.
2. Where does this charge come from?
   • Answer: The charge comes from the movement of electrons within atoms. When two objects are rubbed together, electrons can be transferred from one object to another.

• Electrons:
  • Discovered by J.J. Thomson.
  • Symbol: e–.
  • Charge: -1 (negative).
  • Mass: Considered negligible compared to protons.
• Protons:
  • Discovered through the study of canal rays by E. Goldstein in 1886.
  • Symbol: p+.
  • Charge: +1 (positive).
  • Mass: Approximately 2000 times that of an electron; taken as one unit.

• Atoms are composed of protons and electrons:
  • Charge Balance: The positive charge of protons balances the negative charge of electrons.
  • Location:
    • Protons are located in the nucleus (interior of the atom).
    • Electrons are found in orbits around the nucleus and can be easily removed.

4.2 - The Structure of an Atom

Overview

• Dalton’s Atomic Theory: Proposed that atoms are indivisible and indestructible.
  • Key Limitation: The discovery of sub-atomic particles (electrons and protons) invalidated the idea that atoms are indivisible.
    
    

• With the identification of electrons and protons, scientists needed to understand:
  • Arrangement of Particles: How these charged particles are structured within the atom.
  • Behavior of Atoms: The implications of this structure on the properties of elements and compounds.
    
    

• J.J. Thomson:
  • Contribution: Proposed the first model for the structure of an atom, known as the Plum Pudding Model.
  • Description of the Model:
    • Visualized the atom as a positively charged "soup" with negatively charged electrons embedded within it, similar to plums in a pudding.
    • This model was an early attempt to explain the distribution of electrons in an atom.

4.2.1 - Thomson’s Model of an Atom

Description of the Model

• Analogy: Thomson compared the atom to a Christmas pudding or a watermelon:
  • Christmas Pudding Analogy:
    • Electrons are like currants (dried fruits) embedded in a positively charged spherical "pudding."
  • Watermelon Analogy:
    • The positive charge represents the red edible part of the watermelon.
    • Electrons are likened to the seeds scattered throughout the fruit.

• Biography:
  • Born: December 18, 1856, in Cheetham Hill, Manchester, England.
  • Nobel Prize: Awarded the Nobel Prize in Physics in 1906 for his work on the discovery of electrons.
  • Leadership: Directed the Cavendish Laboratory at Cambridge for 35 years, during which seven of his research assistants won Nobel prizes.

1. Structure:
   • An atom consists of a positively charged sphere with electrons embedded within it.
2. Charge Neutrality:
   • The negative charge of electrons is equal in magnitude to the positive charge of the sphere, making the atom as a whole electrically neutral.

Limitations of Thomson’s Model

• Although Thomson’s model effectively explained the neutrality of atoms, it could not account for the results of subsequent experiments conducted by other scientists. These experiments revealed inconsistencies that prompted further investigation into atomic structure.

4.2.2 - Rutherford’s Model of an Atom

Background

• Ernest Rutherford: A physicist who aimed to understand the arrangement of electrons within an atom.

• Setup: Rutherford conducted an experiment using fast-moving alpha (α) particles directed at a thin gold foil.
• Choice of Gold Foil:
  • Selected for its thinness; the gold foil was approximately 1000 atoms thick. This allowed for minimal interference with the α-particles.
• Alpha Particles:
  • Defined as doubly charged helium ions.
  • Mass: Each α-particle has a mass of 4 u (atomic mass units), giving them significant energy due to their speed.
• Expectations:
  • Rutherford anticipated that the α-particles would be deflected by the sub-atomic particles in the gold atoms.
  • Since α-particles are much heavier than protons, he did not expect to see large deflections; he believed most particles would pass through with only minor deflections.

Observations from Rutherford’s Experiment

Rutherford’s α-particle scattering experiment yielded unexpected results, which led to significant conclusions about atomic structure:

1. Most α-Particles Passed Through:
   • The majority of the fast-moving α-particles went straight through the gold foil without any deflection.
2. Small Angle Deflections:
   • A portion of the α-particles were deflected at small angles, suggesting some interaction with the atoms in the foil.
3. Significant Rebounds:
   • Surprisingly, approximately 1 in every 12,000 α-particles was observed to rebound directly back toward the source.
   • Rutherford’s Remark: He likened this to firing a 15-inch shell at a piece of tissue paper and having it bounce back, emphasizing the improbability of such an occurrence.

• The results indicated that the atom is mostly empty space, with a small, dense, positively charged center (the nucleus) that causes the significant deflection of some α-particles.
• This led to the proposal of a new atomic model, where electrons orbit a dense nucleus, fundamentally changing the understanding of atomic structure.

• Birth: Born on August 30, 1871, in Spring Grove, New Zealand.
• Contributions:
  • Known as the ‘Father of Nuclear Physics’.
  • Famous for his groundbreaking work on radioactivity and the discovery of the atomic nucleus through the gold foil experiment.
• Awards:
  • Awarded the Nobel Prize in Chemistry in 1908 for his research contributions.

Rutherford's α-particle Scattering Experiment

1. Activity Example:

• Imagine a child standing in front of a wall with his eyes closed, throwing stones at the wall.
• Each stone that hits the wall makes a sound.
• When stones are thrown ten times, ten sounds are heard.
• However, if the child throws stones at a barbed-wire fence (with gaps), many stones pass through without hitting the wire, and no sound is heard.
• Purpose of the analogy: This illustrates that when there are gaps (like in the barbed-wire fence), objects can pass through without interaction, similar to particles in the atomic model.

• Rutherford conducted an experiment by bombarding a thin gold foil with α-particles.
• He observed how these particles scattered to understand the structure of the atom.

1. Most of the space inside the atom is empty:
   • Reasoning: Most α-particles passed through the gold foil without deflection.
   • This shows that atoms are largely empty space.
2. Positive charge occupies a small space:
   • Observation: Only a few α-particles were deflected from their original path.
   • This suggests that the positive charge is concentrated in a very small region of the atom.
3. Concentration of mass and positive charge in the nucleus:
   • A tiny fraction of α-particles were deflected by 180°, meaning that the mass and positive charge are located in an extremely small, dense region.
   • This dense region is the nucleus.
   • Rutherford calculated that the radius of the nucleus is about 100,000 times smaller than the radius of the atom.

Rutherford’s Nuclear Model of the Atom

1. Nucleus:
   • There is a positively charged center in the atom called the nucleus.
   • Nearly all the mass of the atom is concentrated in the nucleus.
2. Electrons:
   • Electrons revolve around the nucleus in circular orbits.
3. Size of the nucleus:
   • The nucleus is extremely small compared to the overall size of the atom.

• α-particles (alpha particles): These are positively charged particles used in the experiment to probe the structure of the atom.
• Deflection: The change in direction of α-particles after encountering the positive charge in the nucleus helped to map out the distribution of mass and charge within the atom.

Questions and Answers:

Q: Why did Rutherford conclude that most of the space inside the atom is empty?

• Because most of the α-particles passed through the gold foil without being deflected, showing that they encountered no obstacles, meaning that the atom is mostly empty space.

• It indicates that a very small, dense, and positively charged region (the nucleus) is present in the atom, which contains most of the atom's mass.

• Rutherford’s calculations suggest that the nucleus is about 100,000 times smaller than the radius of the atom.

• It was a groundbreaking model that introduced the idea that atoms consist of a small, dense nucleus surrounded by orbiting electrons, a concept that remains fundamental to atomic theory today.

Drawbacks of Rutherford’s Model of the Atom

1. Instability of Electron Orbits:
   • According to classical electrodynamics, any charged particle that moves in a circular orbit undergoes acceleration.
   • Issue: An accelerating charged particle (like an electron) would continuously radiate energy.
   • As the electron loses energy, it would spiral inward and eventually collapse into the nucleus, leading to the collapse of the atom.
2. Contradiction with Atomic Stability:
   • If Rutherford’s model were correct, atoms would be unstable and would collapse, but in reality, atoms are stable.
   • This implies that electrons cannot follow the circular orbits predicted by Rutherford’s model, as that would lead to the disintegration of atoms.
3. Failure to Explain Atomic Spectra:
   • Rutherford’s model could not explain the discrete line spectra observed in the emission of light by atoms.
   • Observation: Atoms emit light at specific wavelengths, leading to distinct spectral lines.
   • If electrons were losing energy continuously, the spectrum should be continuous, not discrete, which contradicts experimental evidence.

4.2.3 - Bohr’s Model of the Atom

4.2.4 - NEUTRONS

Discovery:

• James Chadwick discovered the neutron in 1932.
• Neutrons are neutral subatomic particles (i.e., they carry no charge).
• The mass of a neutron is nearly equal to that of a proton.

• Neutrons are found in the nucleus of almost all atoms, except for hydrogen, which has no neutron in its most common form.
• Neutrons, along with protons, make up the atomic nucleus.

• Neutrons are generally represented as ‘n’ in atomic symbols and equations.

• The mass of an atom is primarily determined by the sum of the masses of protons and neutrons in the nucleus, since electrons have negligible mass in comparison.

• Neutrons play a critical role in nuclear stability. The balance between protons and neutrons in a nucleus influences whether an atom is stable or radioactive.
• Neutrons also participate in nuclear reactions, such as nuclear fission and fusion, making them essential in nuclear physics and energy production.

4.3 - Distribution of Electrons in Different Orbits (Shells)?

4.4 - Valency

Definition:

• Valency is the combining capacity of an atom, which reflects its ability to gain, lose, or share electrons to achieve a fully-filled outermost shell (usually an octet of 8 electrons, except for helium, which has 2 electrons in its outermost shell).

1. Valence Electrons:
   • The electrons in the outermost shell of an atom are called valence electrons.
   • These electrons determine the atom's chemical reactivity and its valency.
2. Octet Rule:
   • Atoms tend to react to achieve a stable configuration, often an octet (8 electrons in the outermost shell).
   • Atoms with a complete octet (like the noble gases) are chemically inert and have a valency of zero.
   • For instance, helium has 2 electrons (a complete outer shell for helium), and other noble gases (like neon and argon) have 8 electrons in their outermost shell, making them stable.
3. Determining Valency:
   • If an atom has less than 4 electrons in its outermost shell, it tends to lose electrons, and its valency is equal to the number of electrons lost.
     • Example: Sodium (Na) has 1 valence electron, so it loses 1 electron and has a valency of 1.
     • Magnesium (Mg) has 2 valence electrons, so it loses 2 electrons, giving it a valency of 2.
   • If an atom has more than 4 electrons in its outermost shell, it tends to gain electrons to complete the octet, and its valency is equal to 8 minus the number of valence electrons.
     • Example: Fluorine (F) has 7 valence electrons, and it gains 1 electron to complete the octet, giving it a valency of 1.
     • Oxygen (O) has 6 valence electrons, so it gains 2 electrons, giving it a valency of 2.
4. Examples of Valency:
   • Hydrogen (H), Lithium (Li), Sodium (Na): 1 electron in their outermost shell; Valency = 1 (since they lose 1 electron).
   • Magnesium (Mg): 2 electrons in the outermost shell; Valency = 2.
   • Aluminium (Al): 3 electrons in the outermost shell; Valency = 3.
   • Fluorine (F): 7 electrons in the outermost shell; Valency = 1 (gains 1 electron to complete the octet).
   • Oxygen (O): 6 electrons in the outermost shell; Valency = 2 (gains 2 electrons).

• Valency is the number of electrons an atom loses, gains, or shares to complete its outermost shell (usually an octet).
• Elements with 1 to 3 valence electrons tend to lose electrons, while elements with 5 to 7 valence electrons tend to gain electrons.
• Valency helps in predicting how elements will combine to form compounds, such as in ionic or covalent bonding.

4.5 - Atomic Number and Mass Number

4.5.1 - Atomic Number (Z)

• Definition: The atomic number (Z) is the number of protons present in the nucleus of an atom. It is the defining property of an element.
• Importance:
  • The atomic number uniquely identifies each element.
  • All atoms of an element have the same atomic number.
  • The atomic number determines the chemical behavior of an element because it also defines the number of electrons (in a neutral atom).
• Example:
  • For Hydrogen, Z = 1 (it has 1 proton).
  • For Carbon, Z = 6 (it has 6 protons).

4.5.2 - Mass Number (A)

• Definition: The mass number (A) is the total number of protons and neutrons present in the nucleus of an atom.
  • Since electrons have negligible mass, the atomic mass is primarily due to protons and neutrons, which are collectively referred to as nucleons.
• Formula: $$\text{Mass Number (A)} = \text{Number of Protons (Z)} + \text{Number of Neutrons}$$
• Example:
  • For Carbon, A = 12 (6 protons + 6 neutrons).
  • For Aluminium, A = 27 (13 protons + 14 neutrons).

• An element is represented by its symbol, along with its atomic number (Z) and mass number (A). This is written in the following form: $$^{A}_{Z}\text{Element Symbol}$$
• Example:
  • Nitrogen can be written as: $$^{14}_{7}\text{N}$$This means nitrogen has 7 protons (Z = 7) and a total of 14 nucleons (A = 14), so it has 7 neutrons (14 - 7 = 7).

• The atomic number (Z) represents the number of protons in an atom, which defines the element.
• The mass number (A) is the sum of the protons and neutrons, determining the total mass of the atom.
• Elements are represented using their atomic number and mass number, such as $^A_Z\text{Element Symbol}$.

4.6 - Isotopes

Definition:

• Isotopes are atoms of the same element that have the same atomic number (i.e., the same number of protons) but different mass numbers (i.e., different numbers of neutrons).

1. Hydrogen has three isotopes:
   • Protium (  $^1_1\text{H}$): 1 proton, 0 neutrons, mass number = 1.
   • Deuterium (  $^2_1\text{H}$ or D): 1 proton, 1 neutron, mass number = 2.
   • Tritium (  ${}_1^3\text{H or T): 1 proton, 2 neutrons, mass number = 3. <span style="font-size: 1.07rem;"><br></span>}$
2. Carbon has two isotopes:
   • Carbon-12 ( $^12_6\text{C}$): 6 protons, 6 neutrons, mass number = 12.
   • Carbon-14 ( $^14_6\text{C}$): 6 protons, 8 neutrons, mass number = 14.
3. Chlorine has two isotopes:
   • Chlorine-35 ( $^35_{17}\text{Cl}$): 17 protons, 18 neutrons, mass number = 35.
   • Chlorine-37 ( $^37_{17}\text{Cl}$): 17 protons, 20 neutrons, mass number = 37.

• Same Atomic Number (Z): Isotopes have the same number of protons and therefore belong to the same element.
• Different Mass Numbers (A): Isotopes differ in their number of neutrons, which results in different mass numbers.
• Similar Chemical Properties: Since the chemical behavior of an atom is largely determined by the number of protons and electrons, isotopes of the same element have similar chemical properties.
• Different Physical Properties: The physical properties of isotopes, such as mass, can differ due to the different numbers of neutrons.

4.6.1 - Isobars

Definition:

• Isobars are atoms of different elements that have different atomic numbers (i.e., different numbers of protons) but the same mass number (i.e., the same total number of nucleons—protons and neutrons combined).
  
  

1. Calcium (Ca) and Argon (Ar):
   • Calcium has an atomic number of 20 (20 protons) and a mass number of 40.
   • Argon has an atomic number of 18 (18 protons) and a mass number of 40.
   • Both calcium and argon have the same mass number, but their atomic numbers are different. Therefore, they are isobars.
     
     

• Different Atomic Numbers (Z): Isobars have different numbers of protons, meaning they belong to different elements.
• Same Mass Number (A): Despite having different atomic numbers, isobars have the same total number of protons and neutrons, leading to the same mass number.
• Different Chemical Properties: Since they belong to different elements, isobars exhibit different chemical properties.
  
  

• Isobars are atoms from different elements that have the same mass number but different atomic numbers. An example is calcium and argon, both of which have a mass number of 40 but different atomic numbers.

NCERT Science Notes - Class 9 | Science | Chapter 4 - Structure of Atom

NCERT Science Notes - Class 9 | Science | Chapter 4 - Structure of Atom