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STRUCTURE OF ATOM in One Shot - From Zero to Hero || Class 9th

Physics Wallah Foundation・2 minutes read

Chemistry teacher Senanayak introduces students to the Structure of Atom chapter, covering Dalton's Atomic Theory, Thomson's experiment, Rutherford's model, Bohr's model, electron configuration rules, valency, isotopes, and practical applications. The text also explores the concept of atomic mass, isotopes with varying mass numbers like chlorine 35 and 37, fractional atomic mass calculations, and the differences between isotopes and isobars, providing insights into their chemical and physical properties.

Insights

Dalton's Atomic Theory initially proposed indivisible atoms but later research revealed atoms as divisible, composed of electrons, protons, and neutrons with specific masses and charges.
Rutherford's Alpha Particle Scattering Experiment showed atoms as mostly empty space with a dense, positively charged nucleus, leading to the model of electrons orbiting the nucleus like planets around the sun.
Bohr's Model of the Atom introduced stable electron orbits with fixed energy levels, emphasizing that electrons in discrete orbits do not lose energy, contributing to atomic stability and electron arrangement in shells for stability and valency determination.

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Recent questions

What did Rutherford's experiment reveal about the atom?
Most of the atom is empty space.

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Summary

00:00

Atomic Structure: From Indivisible to Divisible

Chemistry teacher Senanayak welcomes students to an online platform discussing the ninth chapter on the Structure of Atom.
Dalton's Atomic Theory, foundational to modern chemistry, initially posited atoms as indivisible.
Further research revealed atoms as divisible, with electrons, protons, and neutrons as atomic particles.
J. Thomson's 1897 experiment with cathode rays led to the discovery of electrons, negatively charged particles.
Electrons have a mass of 9*10^-31 kg and a charge of 1.6*10^-19 C.
Protons, positively charged particles, were discovered by Goldstein and have a mass of 1.6*10^-27 kg and a charge of 1.6*10^-19 C.
Neutrons, neutral particles, were discovered by Chadwick in 1932 and have a mass of 1.6*10^-27 kg.
Thomson's model of the atom likened it to a watermelon, with positive charge uniformly distributed and electrons embedded like seeds.
Thomson's model was also compared to a plum pudding, with positive charge spread like pudding and electrons as dry fruits.
The model emphasized the equal distribution of positive and negative charges within the atom.

16:10

Rutherford's Atom: Empty Space, Positive Nucleus

If the charge is plus 25, then the negative charge will be -25 to cancel each other out, resulting in an electrically neutral atom.
The experiment conducted by respected scientist Ernest Rutherford involved alpha particle scattering, where fast-moving alpha particles with a considerable amount of energy were used.
Rutherford's experiment observed that most alpha particles passed straight through the gold foil, indicating that most of the space inside the atom is empty.
A small fraction of alpha particles were deflected by the gold foil at small angles, suggesting the presence of positive charge concentrated in a small volume within the atom.
Out of every 12,000 alpha particles, one was observed to rebound at 180 degrees, indicating that the positive charge and mass of the atom are concentrated in a very small space.
Rutherford concluded that the space inside the atom is mostly empty, with very few alpha particles being deflected, indicating a small positive charge occupying a small space.
The nucleus of an atom was discovered to be positively charged, very dense, and significantly smaller in size compared to the atom itself.
Rutherford's model of the atom included a positively charged nucleus at the center, with electrons revolving around it in circular paths, similar to planets orbiting the sun.
The drawback of Rutherford's model was that electrons moving in circular orbits would undergo acceleration, lose energy, and eventually fall into the nucleus, making the atom unstable.
The drawback stemmed from the fact that charged particles, like electrons, radiate energy when undergoing acceleration, leading to a loss of energy and potential collapse of the atom.

31:27

Atomic Models and Electron Stability Rules

Rutherford's model of the atom was well understood after his work.
Bohr's model of the atom, introduced in 1913, focused on stable electron orbits.
Electrons in an atom revolve in discrete orbits with fixed energy levels.
Electrons in discrete orbits do not lose energy and remain stable.
Energy levels in an atom are represented by letters (K, L, M, N, etc.) or numbers (1, 2, 3, 4, etc.).
Energy levels closer to the nucleus have minimum energy, while those farther have maximum energy.
The atomic number of an element equals the number of protons in its atom.
In a neutral atom, the atomic number is equal to the number of protons and electrons.
Atoms are stable when they have minimum energy, occupying low energy levels first.
Electrons are filled in orbits following specific rules to achieve stability.

47:08

"Atomic Number, Electron Arrangement, and Valency"

The atomic number of neutral atoms is equal to the number of electrons they have.
The number of electrons in an atom and how to fill them in orbits is determined by the atomic number.
Electrons are filled in shells following specific rules, with the maximum number of electrons in a shell given by 2n².
The maximum number of electrons in different shells is determined by the formula 2n².
The third rule states that the maximum number of electrons in the outermost orbit is determined by the outermost shell, with a maximum of 8 electrons.
The maximum number of electrons in the outermost orbit is crucial for stability, with 8 electrons providing stability.
The fourth rule emphasizes that the maximum number of electrons in the outermost shell cannot exceed 8 for stability.
The arrangement of electrons in different shells is called electronic configuration, following the rules of filling electrons.
The electronic configuration of elements like hydrogen, helium, lithium, magnesium, and potassium is explained based on the rules.
Valency is determined by the number of valence electrons in the outermost shell of an atom, crucial for understanding chemical bonding.

01:03:53

Valence electrons determine atom's chemical properties

Valence electrons are crucial for determining an atom's chemical properties.
The number of valence electrons in an atom dictates its combining capacity.
Elements with a similar number of valence electrons exhibit comparable chemical properties.
Lithium, sodium, and potassium exemplify elements with distinct valence electron numbers.
Achieving stability is the primary goal in the world of atoms.
Stability is attained by completing octants or achieving duplets in the outermost shell.
Sodium, magnesium, and aluminum showcase different methods to achieve stability through electron manipulation.
Valency is determined by the number of electrons gained, lost, or shared by an element.
The valency of an element is crucial in understanding its combining capacity.
The mass number of an atom is the sum of protons and neutrons in its nucleus, defining its overall mass.

01:20:34

Understanding Isotopes and Atomic Masses

Mass number is equal to the atomic number plus the number of neutrons, represented as mass number = atomic number + number of neutrons.
Protons are collectively called the nucleus, while the number of protons is equal to the atomic number.
Isotopes are atoms of an element with the same atomic number but different mass numbers due to varying numbers of neutrons.
The number of neutrons in isotopes is determined by subtracting the number of protons from the mass number.
Chlorine has isotopes with atomic number 17 but different mass numbers, resulting in varying numbers of neutrons.
Hydrogen has three isotopes with the same atomic number but different mass numbers.
Isotopes have identical chemical properties due to the same number of protons and electrons, influencing valence electrons and chemical properties.
Physical properties of isotopes differ based on their mass numbers, affecting properties like density, melting point, and boiling point.
The concept of fractional atomic mass or average atomic mass is exemplified by chlorine's atomic mass of 35.5, indicating isotopes with varying mass numbers.
Fractional atomic mass is represented by decimal values, reflecting the average mass of isotopes of an element.

01:36:14

Atomic Mass and Isotopes: Chlorine Example

The text discusses the concept of atomic mass and isotopes, focusing on chlorine isotopes 35 and 37.
It explains how to calculate the fractional atomic mass based on the natural abundance of each isotope.
The average atomic mass of chlorine is determined to be 35.5 u, representing a mix of isotopes.
The text then delves into the practical applications of isotopes, such as in nuclear reactions, cancer treatment, and goiter treatment.
It contrasts isotopes with isobars, which are atoms of different elements with the same mass number but different atomic numbers.
Isobars exhibit different chemical and physical properties due to variations in electron configuration and valence electrons.

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