Structure of Atom NEET Questions: Answer Key, pdf

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Structure of Atom topic is important for NEET instruction, that specialize in atomic models, subatomic debris, and electron configuration. Key concepts consist of the historical improvement of atomic theory, quantum mechanics, and the association of electrons in shells and subshells. Understanding the periodic desk’s dating to atomic structure is crucial for tackling related questions. Mastery of those standards enhances trouble-fixing capabilities and boosts self belief for the NEET examination, permitting college students to correctly analyze and interpret atomic conduct in various chemical contexts.

Introduction to the Structure of Atom
Download: Structure of Atom
Atomic Models: Structure of Atom
Subatomic Particles: Structure of Atom
Atomic Number and Mass Number
Electron Configuration: Structure of Atom
Periodic Table and Atomic Structure
Chemical Bonding and Atomic Structure
Quantum Numbers: Structure of Atom
FAQs about Structure of Atom

Introduction to the Structure of Atom

The Structure of Atom is a fundamental subject matter within the NEET (National Eligibility cum Entrance Test) syllabus, essential for information the basics of chemistry and physics. This phase explores the arrangement and behavior of subatomic debris—protons, neutrons, and electrons—within an atom. It covers ideas together with atomic fashions, quantum numbers, and electron configurations, which can be critical for greedy chemical bonding and reactions. NEET questions on this subject matter frequently require college students to use theoretical information to practical eventualities, reinforcing their comprehension and analytical skills. Mastering the shape of the atom lays the groundwork for advanced studies in chemistry, making it a vital area for NEET aspirants aiming for success inside the examination.

Historical Background

The concept of the atom has evolved over centuries. Key ancient figures and their contributions consist of:

Democritus: A Greek truth seeker who proposed the concept of indivisible debris referred to as “atomos.”
John Dalton: An English chemist who revived the atomic concept, suggesting that atoms are solid spheres.
J.J. Thomson: A British physicist who observed the electron, a negatively charged particle within the atom.
Ernest Rutherford: A New Zealand-born physicist who proposed the nuclear version of the atom, with a dense, definitely charged nucleus surrounded through negatively charged electrons.
Niels Bohr: A Danish physicist who subtle Rutherford’s model, introducing the idea of quantized power ranges for electrons.
Erwin Schrödinger and Werner Heisenberg: Scientists who advanced quantum mechanics, offering a more complex and correct description of the atom, consisting of the wave-like nature of electrons.

Fundamental Concepts

Subatomic Particles:

Protons: Positively charged particles discovered inside the nucleus.
Neutrons: Neutral debris found within the nucleus.
Electrons: Negatively charged particles that orbit the nucleus.

Atomic Number: The wide variety of protons in an atom’s nucleus, which determines its element.

Mass Number: The total variety of protons and neutrons in an atom’s nucleus.

Isotopes: Atoms of the identical element with one of a kind numbers of neutrons.

Atomic Orbitals: Regions of area around the nucleus in which electrons are probable to be located.

Electron Configuration: The arrangement of electrons in an atom’s orbitals.

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Atomic Models: Structure of Atom

Model	Description
Dalton’s Atomic Theory	– Atoms are indivisible and indestructible spheres. – Atoms of the same element are identical. – Atoms of different elements combine in simple whole-number ratios to form compounds.
Thomson’s Model (Plum Pudding Model)	– Atoms are positively charged spheres with negatively charged electrons embedded in them, like plums in a pudding.
Rutherford’s Model (Nuclear Model)	– Atoms have a small, dense, positively charged nucleus at the center. – Negatively charged electrons orbit the nucleus in circular paths.
Bohr’s Model of the Hydrogen Atom	– Electrons orbit the nucleus in fixed energy levels or shells. – Electrons can jump between energy levels by absorbing or emitting energy.
Quantum Mechanical Model	– Electrons exist in orbitals, regions of space where there is a high probability of finding an electron. – The exact position of an electron cannot be determined precisely.

Subatomic Particles: Structure of Atom

Subatomic Particle	Charge	Mass (amu)	Location in Atom	Properties and Characteristics
Proton	+1	1	Nucleus	Positively charged, determines the element’s identity, found in the nucleus with neutrons
Neutron	0	1	Nucleus	Neutral, contributes to the atom’s mass, found in the nucleus with protons
Electron	-1	Negligible (1/1836 amu)	Electron cloud (orbitals)	Negatively charged, involved in chemical reactions, orbits the nucleus in energy levels

Atomic Number and Mass Number

Definition and Significance

Atomic Number (Z):

The number of protons in an atom’s nucleus.
It defines the element’s identity.
All atoms of a particular element have the same atomic number.

Mass Number (A):

The total number of protons and neutrons in an atom’s nucleus.
It approximates the atomic mass of an atom.

Isotopes and Isobars

Isotopes:

Atoms of the same element with the same atomic number but different mass numbers.
They have the same number of protons but different numbers of neutrons.
Examples: Carbon-12, Carbon-13, Carbon-14.

Isobars:

Atoms of different elements with different atomic numbers but the same mass number.
They have different numbers of protons and neutrons but the same total number of nucleons.
Example: Calcium-40 and Argon-40.

Calculation of Average Atomic Mass

The average atomic mass of an element is the weighted average of the masses of its isotopes. It is calculated using the following formula:

Average Atomic Mass = Σ (fractional abundance × isotopic mass)

Steps:

Determine the fractional abundance of each isotope: This is the percentage of each isotope in a natural sample of the element.
Multiply the fractional abundance of each isotope by its isotopic mass.
Sum up the products from step 2.6

Example:

Consider an element with two isotopes:

Isotope 1: Mass = 10 amu, Abundance = 20%
Isotope 2: Mass = 11 amu, Abundance = 80%

Average Atomic Mass = (0.20 × 10 amu) + (0.80 × 11 amu) = 10.8 amu

Electron Configuration: Structure of Atom

Principles of Electron Configuration

Electron configuration describes the arrangement of electrons in an atom’s orbitals. Three fundamental standards govern this arrangement:

Aufbau Principle:

Electrons fill orbitals starting from the lowest strength level to higher energy tiers.
The order of filling is: 1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p, 5s, 4d, 5p, 6s, 4f, 5d, 6p, 7s, 5f, 6d, 7p.

Pauli Exclusion Principle:

No two electrons in an atom may have the same set of 4 quantum numbers.
This approach that an orbital can hold a maximum of 2 electrons with opposite spins.

Hund’s Rule:

Electrons will fill orbitals of the identical energy level with parallel spins earlier than pairing up.
This results in the most stable electron configuration.

Example:

Let’s consider the electron configuration of carbon (atomic number 6).

Fill the 1s orbital: 1s²
Fill the 2s orbital: 1s² 2s²
Fill the 2p orbitals: 1s² 2s² 2p²

The electron configuration of carbon is 1s² 2s² 2p².

Periodic Table and Atomic Structure

Trend	Description	Relation to Atomic Structure
Atomic Radius	Size of an atom	Increases down a group (more energy levels) and decreases across a period (increased nuclear charge pulls electrons closer)
Ionic Radius	Size of an ion	Cations are smaller than their parent atoms, anions are larger. Trends are similar to atomic radius.
Ionization Energy	Energy required to remove an electron	Increases across a period (greater nuclear charge holds electrons more tightly) and decreases down a group (outer electrons are further from the nucleus)
Electron Affinity	Energy change when an atom gains an electron	Generally increases across a period and decreases down a group, with exceptions.
Electronegativity	Ability of an atom to attract electrons in a bond	Increases across a period and decreases down a group.

Group/Period	Characteristics
Group 1 (Alkali Metals)	Soft, reactive metals, easily lose one electron to form +1 ions.
Group 2 (Alkaline Earth Metals)	Reactive metals, easily lose two electrons to form +2 ions.
Group 17 (Halogens)	Nonmetals, highly reactive, gain one electron to form -1 ions.
Group 18 (Noble Gases)	Nonreactive gases, have complete valence electron shells.
Period 1	Smallest atoms, few electrons.
Period 2	Second smallest atoms, more complex electron configurations.
Period 3	Larger atoms than periods 1 and 2, more complex electron configurations.
Period 4 and beyond	Increasingly complex electron configurations, introduction of d and f orbitals.

Chemical Bonding and Atomic Structure

Ionic and Covalent Bonds

Ionic Bonds

Formed among a metal and a nonmetal.
Involve the transfer of electrons from a metal atom to a nonmetal atom.
Result in the formation of ions: cations (positive ions) and anions (negative ions).
Strong electrostatic forces keep the ions together.
Example: NaCl (sodium chloride)

Covalent Bonds

Formed between nonmetal atoms.
Involve the sharing of electron pairs among atoms.
Can be polar or nonpolar, depending on the electronegativity difference between the atoms.
Polar covalent bonds: Unequal sharing of electrons.
Nonpolar covalent bonds: Equal sharing of electrons.
Example: H₂O (water)

Lewis Structures

Lewis structures are diagrams that show the bonding among atoms of a molecule and the lone pairs of electrons that can exist in the molecule.

Steps to draw a Lewis structure:

Count the valence electrons of all atoms within the molecule.
Determine the central atom: Usually the least electronegative atom.
Connect the atoms to the central atom with single bonds.
Distribute the remaining electrons as lone pairs to complete octets for each atom (except hydrogen, which needs only 2 electrons).
If necessary, form double or triple bonds to complete octets.

VSEPR Theory

VSEPR (Valence Shell Electron Pair Repulsion) theory predicts the geometry of molecules based on the repulsion between electron pairs around a central atom.

Basic Shapes:

Linear: Two electron pairs, bond angle 180°.
Trigonal planar: Three electron pairs, bond angle 120°.
Tetrahedral: Four electron pairs, bond angle 109.5°.
Trigonal bipyramidal: Five electron pairs.
Octahedral: Six electron pairs.

Factors affecting molecular geometry:

Lone pairs: Lone pairs exert more repulsion than bonding pairs, influencing bond angles.
Multiple bonds: Multiple bonds (double or triple bonds) count as one electron pair for VSEPR purposes.

Quantum Numbers: Structure of Atom

Types of Quantum Numbers

Type of Quantum Number	Symbol	Significance	Allowed Values
Principal Quantum Number	n	Energy level, size of orbital	Positive integers (1, 2, 3, …)
Azimuthal (Angular Momentum) Quantum Number	l	Shape of orbital	Integers from 0 to n-1
Magnetic Quantum Number	m_l	Orientation of orbital	Integers from -l to +l
Spin Quantum Number	m_s	Spin of electron	+1/2 or -1/2

Determining Electron Configurations

Aufbau Principle: Electrons fill orbitals in order of increasing energy levels.
Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers.
Hund’s Rule: Orbitals of the same energy level are filled with one electron each before any orbital receives a second electron.

Example: Electron Configuration of Carbon (C)

Atomic number of C = 6, so it has 6 electrons.
Configuration: 1s² 2s² 2p²

Explanation:

1s²: The first energy level (n=1) has one subshell (l=0), which is an s orbital. It holds 2 electrons.
2s²: The second energy level (n=2) has an s subshell (l=0). It holds 2 electrons.
2p²: The second energy level also has a p subshell (l=1). It can hold 6 electrons, but in this case, it only has 2.

FAQs about Structure of Atom

Q. What is the structure of an atom?

Ans: An atom consists of a nucleus (containing protons and neutrons) surrounded by means of electrons in power tiers or shells.

Q. What are protons and neutrons?

Ans: Protons are positively charged debris positioned inside the nucleus, whilst neutrons are neutral particles that still live within the nucleus.

Q. What is the significance of electrons in an atom?

Ans: Electrons are negatively charged and orbit the nucleus, determining the atom’s chemical properties and its capacity to bond with different atoms.