Structure of the Atom Class 9 Notes: CBSE Science Chapter 4

Atoms

• Atoms are the basic units of matter and the defining structure of elements.

• It consists of three basic particles, i.e., protons, electrons, and neutrons, that build the structure of an atom.

• Protons are positively charged particles and were founded by E. Goldstein.

• Electrons are negatively charged particles and were founded by J.J. Thomson.

• Neutrons have no charge and were founded by Chadwick.

• The nucleus is in the centre of the atom and contains protons and neutrons.

• The outer region of the atom which holds electrons in orbit around the nucleus is known as shell/energy level/orbits. 

• These shells are further divided into subshells.

• The electrons present in the outermost shell of an atom are known as the valence electrons.

Thomson’s Model of an Atom

• Proposed by J.J. Thomson in 1897: He discovered the electron and suggested a new model of the atom.

• Plum Pudding Model: Thomson compared the atom to a pudding or a sphere of positive charge, with negatively charged electrons embedded in it, like plums in a pudding.

• No Nucleus: According to this model, there was no central nucleus. The positive and negative charges were spread throughout the atom.

• Neutral Atom: The overall charge of the atom was neutral, as the positive charge balanced out the negative charge of electrons.

• Failed Model: This model was later rejected after Ernest Rutherford discovered the nucleus.

Rutherford’s Model of an Atom

• Proposed by Ernest Rutherford in 1911: Based on his famous gold foil experiment.

• Nucleus: He discovered that the atom has a small, dense, positively charged centre called the nucleus.

• Electrons Orbit: Electrons revolve around the nucleus in circular paths, similar to planets orbiting the sun.

• Mostly Empty Space: Most of the atom’s volume is empty space, where electrons move.

• Positive Charge in Nucleus: The nucleus contains positively charged protons, which hold most of the atom's mass.

• Failed to Explain Stability: While this model explained the atom's structure, it couldn’t explain why electrons don’t spiral into the nucleus due to attraction, leading to the development of newer models.

Rutherford’s Experiment: Scattering of α-particles by a Gold Foil

This scattering experiment was crucial in proposing the nuclear model of the atom.

1. Conducted by Rutherford in 1909: Rutherford and his team used a thin sheet of gold foil to observe how alpha (α) particles behave when they hit it.

2. Alpha Particles: These are positively charged particles (helium nuclei) emitted from a radioactive source.

3. Most Passed Through: A majority of the α-particles passed straight through the gold foil without deflection, indicating that most of the atom is empty space.

4. Some Deflected: A few α-particles were deflected at small angles, suggesting that they encountered a small, positively charged region.

5. Few Rebounded: A very small number of particles bounced back, indicating they hit a dense, positively charged centre (nucleus).

6. Discovery of the Nucleus: This experiment led to the conclusion that the atom has a small, dense nucleus where most of its positive charge and mass are concentrated.

Drawbacks of Rutherford’s Model of the Atom

1. Electrons in Orbit: According to Rutherford’s model, electrons revolve around the nucleus, which should cause them to lose energy and collapse into the nucleus.

2. Problem with Classical Physics: Classical physics predicts that electrons in motion would radiate energy and eventually spiral into the nucleus, leading to atom instability.

3. Bohr’s Explanation: Niels Bohr solved this by proposing that electrons can only exist in specific, quantized orbits without radiating energy.

4. Stable Energy Levels: In Bohr’s model, as long as electrons remain in these specific orbits (energy levels), they don’t lose energy and the atom stays stable.

5. Quantum Mechanics: Modern atomic theory based on quantum mechanics explains that electrons have stable energy states, which prevents them from collapsing into the nucleus.

Bohr’s Model of Atom

• Proposed by Niels Bohr in 1913: Bohr modified Rutherford’s model to explain atomic stability and the line spectra of elements.

• Quantised Orbits: Electrons revolve around the nucleus in specific, fixed orbits or energy levels without radiating energy. These orbits are called quantized orbits.

• Energy Levels: The energy of an electron is fixed for each orbit. The closer the orbit to the nucleus, the lower the energy, and vice versa.

• Electron Transitions: Electrons can jump between energy levels by absorbing or emitting specific amounts of energy (quanta). When an electron jumps from a higher to a lower energy level, it emits energy in the form of light.

• Explains Atomic Spectra: Bohr’s model successfully explained the discrete spectral lines seen in the hydrogen atom’s emission and absorption spectra.

• Stable Orbits: Electrons in a particular orbit do not lose energy, which explains the stability of atoms.

Neutrons:

1. Discovery in 1932: J. Chadwick discovered a new subatomic particle called the neutron.

2. No Charge: Neutrons have no electric charge, unlike protons, which are positively charged.

3. Mass of a Neutron: The mass of a neutron is nearly equal to that of a proton.

4. Presence in Nucleus: Neutrons are found in the nucleus of all atoms except hydrogen, which has only one proton and no neutron.

5. Representation: Neutrons are represented by the symbol 'n'.

6. Mass of Atom: The mass of an atom is determined by the sum of the masses of the protons and neutrons in its nucleus.

How are Electrons Distributed in Different Orbits (Shells)?

Here’s a simplified distribution of electrons in an atom based on Bohr and Bury's rules:

• Electron Distribution Rules: Bohr and Bury suggested rules for how electrons are arranged in different orbits or energy levels in an atom.

• Maximum Electrons in a Shell: The maximum number of electrons in a shell is calculated using the formula 2n², where n is the orbit number (1, 2, 3,...).

• Examples of Shells:

• K-shell (first orbit): Can hold a maximum of 2 electrons (2 × 1² = 2).

• L-shell (second orbit): Can hold up to 8 electrons (2 × 2² = 8).

• M-shell (third orbit): Can hold up to 18 electrons (2 × 3² = 18).

• N-shell (fourth orbit): Can hold up to 32 electrons (2 × 4² = 32), and so on.

• Outermost Shell Rule: The outermost shell can hold a maximum of 8 electrons, regardless of the shell’s total capacity.

• Stepwise Filling of Shells: Electrons fill the shells in a step-by-step manner, meaning inner shells must be filled first before electrons occupy the outer shells.

• Atomic Structure Example: The electron configuration of the first 18 elements follows these rules, with their atomic structures shown in diagrams or tables.

Valency:

1. Definition: Valency is the ability of an atom to combine with other atoms, determined by the number of electrons in the outermost shell.

2. Valency Rule: It depends on the number of electrons in the outermost shell:

• If the outer shell has fewer than 4 electrons, valency = number of outer electrons.

• If the outer shell has more than 4 electrons, valency = 8 - number of outer electrons.

3. Stable Configuration: Atoms try to achieve 8 electrons (octet rule) in their outermost shell by gaining, losing, or sharing electrons.

Schematic Atomic Structure of the First Eighteen Elements:

The atomic structure for each of the first 18 elements is based on the arrangement of electrons in shells.

1. Hydrogen (H): Atomic number 1; Valency = 1; Electron distribution: 1 (K-shell).

2. Helium (He): Atomic number 2; Valency = 0; Electron distribution: 2 (K-shell).

3. Lithium (Li): Atomic number 3; Valency = 1; Electron distribution: 2, 1 (K, L).

4. Beryllium (Be): Atomic number 4; Valency = 2; Electron distribution: 2, 2 (K, L).

5. Boron (B): Atomic number 5; Valency = 3; Electron distribution: 2, 3 (K, L).

6. Carbon (C): Atomic number 6; Valency = 4; Electron distribution: 2, 4 (K, L).

7. Nitrogen (N): Atomic number 7; Valency = 3; Electron distribution: 2, 5 (K, L).

8. Oxygen (O): Atomic number 8; Valency = 2; Electron distribution: 2, 6 (K, L).

9. Fluorine (F): Atomic number 9; Valency = 1; Electron distribution: 2, 7 (K, L).

10. Neon (Ne): Atomic number 10; Valency = 0; Electron distribution: 2, 8 (K, L).

11. Sodium (Na): Atomic number 11; Valency = 1; Electron distribution: 2, 8, 1 (K, L, M).

12. Magnesium (Mg): Atomic number 12; Valency = 2; Electron distribution: 2, 8, 2 (K, L, M).

13. Aluminium (Al): Atomic number 13; Valency = 3; Electron distribution: 2, 8, 3 (K, L, M).

14. Silicon (Si): Atomic number 14; Valency = 4; Electron distribution: 2, 8, 4 (K, L, M).

15. Phosphorus (P): Atomic number 15; Valency = 3; Electron distribution: 2, 8, 5 (K, L, M).

16. Sulfur (S): Atomic number 16; Valency = 2; Electron distribution: 2, 8, 6 (K, L, M).

17. Chlorine (Cl): Atomic number 17; Valency = 1; Electron distribution: 2, 8, 7 (K, L, M).

18. Argon (Ar): Atomic number 18; Valency = 0; Electron distribution: 2, 8, 8 (K, L, M).

Atomic Number

• The atomic number of an element is the same as the number of protons in the nucleus of its atom. As atoms are electrically neutral, an atom contains as many electrons as it has protons. The atomic number is denoted by Z.

Mass Number

• The mass number of an atom is equal to the number of nucleons in its nucleus. Nucleons are the collective term for protons and neutrons. The mass number is denoted by A.

• In the notation of an atom, the atomic number is written as a subscript on the left of the element symbol and the mass number is written as a superscript on the left of the element symbol.

Isotopes

• Isotopes can be defined as elements that possess the same number of protons and electrons, but a different number of neutrons.

• For example, Protium, deuterium, and tritium are the isotopes of hydrogen. They each have one single proton Z=1 and a single electron but differ in the number of their neutrons. Hydrogen has no neutrons, deuterium has one, and tritium has two neutrons.

Isobars and Isotones

• Isobar is an element that differs in chemical property but has the same physical property which means isobars are those elements that have a different atomic number but the same mass number. 

• For example, Calcium and chlorine are isobars since both have a mass number of 40 but calcium has an atomic number of 20 and chlorine has an atomic number of 17.

• Isotones are elements that have the same number of neutrons but different atomic numbers.

• For example, Chlorine with atomic number 37 and potassium with atomic number 39 is isotones because both chlorine and potassium have the same number of neutrons i.e., 20.

Chapter 4 Science Class 9 Notes: Atomic Number and Mass Number

Chapter 4 Science Class 9 notes tell us that the total number of protons present in an element is called the atomic number while the mass number is both the number of protons plus the number of neutrons.

Atomic Number

According to Class 9 Chapter 4 Science notes, an atomic number is the number of the chemical elements present in the periodic system so that the elements are organised in order of the increasing number of protons present in the nucleus. Also, the total number of protons that are equal to the electrons in a neutral atom is called the atomic number. For example, iron has 26 protons in its nucleus; hence, the atomic number of the matter is 26.  An equal number of electrons and protons are presented in a neutral atom.

Mass Number

According to Class 9th Science Chapter 4 notes, the total number of protons and neutrons is combined together and forms an element’s mass number. The contribution of the mass from the electron is disdained and then calculated as the mass number. Hence, the approximation of the mass is used to calculate the number of neutrons that an element has by subtracting the total number of protons from the mass number. Both neutrons and protons weigh about 1 atomic mass unit or amu. Isotopes of this same element will have the same number of atoms but a different mass.

Class 9 Science Chapter 4 Notes: Details of Isotopes

Notes of Chapter 4 Science Class 9 also give us a brief description of isotopes. It is the variant of a specific chemical element that differs in its number of neutrons and also in the number of nucleons. The isotopes have the same number of protons and a different number of neutrons. Thus, they have the same atomic number but a different mass number. All elements are a mixture of isotopes. Each of them is a pure substance. They have the same chemical properties but different physical properties. If there is an absence of isotope in an element, then its mass would be the same as the total number of neutrons and protons. Chapter 4 Science Class 9 notes also tell us that if there is a presence of isotope in any element, then the percentage of all the isotope forms should be known and then the average mass should be calculated.

Science Class 9 Chapter 4 Notes: Details of Isobars

The nuclides of various chemical elements are called isobars. They have the same number of nucleons which means that they have the same mass number. But they differ in atomic number. Alfred Walter Stewart suggested the name isobars. It has been derived from the Greek word isos which means equal and bars, which means weight.

Important Topics of Class 9 Chapter 4 Science Structure of the Atom You Shouldn’t Miss!

Here are the important topics of Class 9 Science Chapter 4 Structure of the Atom Notes that students should focus on:

1. Discovery of Subatomic Particles: Learn about the discovery of electrons, protons, and neutrons.

2. Thomson's Model of the Atom: Understand the key points of Thomson’s model and its limitations.

3. Rutherford’s Model of the Atom: Explore Rutherford’s famous gold foil experiment and conclusions.

4. Bohr’s Model of the Atom: Grasp the key postulates of Bohr’s atomic model.

5. Isotopes and Isobars: Study the differences and significance of isotopes and isobars.

6. Atomic Number and Mass Number: Understand the importance of these properties in determining atomic structure.

Importance of Class 9 Science Notes Chapter 4 Structure of the Atom PDF

Class 9 Science Notes Chapter 4 Structure of the Atom in PDF format is crucial for several reasons:

1. Comprehensive Understanding: The notes provide a detailed explanation of atomic structure, helping students grasp fundamental concepts such as subatomic particles and atomic models.

2. Structured Learning: Well-organised notes break down complex topics into manageable sections, making it easier to study and retain information.

3. Quick Revision: Having a PDF version allows for easy access and quick revision, which is essential during exam preparation.

4. Visual Aids: Many PDFs include diagrams and models that visually represent atomic structure, enhancing comprehension.

5. Exam Readiness: Concise and focused notes help in efficient revision, aiding students in better performance in tests and exams.

Tips for Learning the Class 9 Science Chapter 4 Notes Structure of the Atom

Here are some tips for learning Class 9 Science Chapter 4: Structure of the Atom:

• Understand Key Concepts: Start by thoroughly understanding the fundamental concepts such as subatomic particles (electrons, protons, and neutrons) and their roles in atomic structure.

• Study Models: Pay attention to different atomic models (Thomson’s, Rutherford’s, and Bohr’s models). Understand the significance of each model and how they contributed to the current understanding of the atom.

• Use Diagrams: Utilise diagrams and visual aids to grasp the arrangement of subatomic particles and atomic structure. This can help in visualizing concepts better.

• Create Summaries: Summarize each section of the notes in your own words to reinforce your understanding and make revision easier.

• Practice Questions: Work on practice questions and problems related to atomic structure to test your understanding and application of the concepts.

Conclusion

Mastering Class 9 Science Chapter 4 Structure of the Atom is essential for a solid foundation in chemistry. By focusing on key concepts such as the discovery of subatomic particles, understanding various atomic models, and the significance of isotopes and isobars, students can gain a comprehensive understanding of atomic structure. Utilizing structured notes, visual aids, and regular revision can significantly enhance learning and retention. Engaging with practice questions and discussing topics with peers further reinforces comprehension. Following these strategies will not only aid in exam preparation but also build a strong base for future scientific studies.

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