Welcome to iPrep, your Learning Super App. Our learning resources for Chapter 2, “Structure of Atom,” in Class 11 Chemistry are meticulously designed to ensure students gain a comprehensive understanding of this essential topic. These resources include detailed notes on atomic theory, the arrangement of electrons in shells, and the quantum mechanical model of the atom. Students will also explore the concepts of atomic orbitals, electron configuration, and periodic trends. The materials provide visual aids and diagrams to illustrate the structure of atoms and include practice questions to solidify understanding. These resources are crafted to help students grasp the fundamental concepts necessary for studying the chemistry of biological systems.

Meaning of The Structure Of Atom

The concept of “Structure of Atom” in Class 11 Chemistry delves into the foundational principles of life by exploring the basic building blocks of matter. It examines the atomic theory, crucial for understanding how elements and compounds interact to form biological molecules. The chapter covers the arrangement of electrons in atoms and how this arrangement influences chemical bonding and molecular structure. This understanding is essential for comprehending how atoms combine to form complex biomolecules such as proteins, nucleic acids, and lipids. By linking the structure of atoms to chemical functions, students gain insight into the chemical basis of life processes.

Discovery of Subatomic Particles

The atom, once thought to be indivisible, was found to consist of smaller particles. The discovery of these subatomic particles electrons, protons, and neutrons revolutionized the understanding of the structure of the atom.

Properties of Cathode Rays

The chapter “Structure Of Atom” significantly covers the properties of cathode rays. These Cathode rays, discovered in the late 19th century, are streams of electrons emitted from the cathode in a vacuum tube. They exhibit several properties:

1. Travel in straight lines.
2. It causes fluorescence when it strikes certain materials.
3. Are deflected by electric and magnetic fields, indicating they carry a negative charge.
4. Possess energy and momentum.

Charge to Mass Ratio of Electron

J.J. Thomson, through his experiments with cathode rays, calculated the charge-to-mass ratio (e/m) of the electron. This was a significant step toward understanding the properties of the electron.

Charge on the Electron

The charge of an electron was determined by Robert Millikan through the oil drop experiment. The elementary charge was found to be approximately 1.602×10−191.602 \times 10^{-19}1.602×10−19 coulombs.

Discovery of Protons and Neutrons

• Protons: Positively charged particles found in the nucleus of the atom. Their discovery by Ernest Rutherford followed the identification of the electron.
• Neutrons: Discovered by James Chadwick, neutrons are neutral particles also located in the nucleus and contribute to the atom’s mass but not its charge.

Atomic Models

Various atomic models have been proposed to explain the structure of atoms:

Thomson model, also known as the “plum pudding model” of the atom, can be described with the following key points:

1. Uniform Positive Charge: J.J. Thomson proposed that an atom consists of a uniformly distributed positive charge (like a “pudding”) in which electrons (the “plums”) are embedded. This positive charge was thought to balance out the negative charge of the electrons, making the atom electrically neutral.
2. Electron Distribution: According to this model, electrons were scattered throughout the positively charged “pudding” like raisins in a plum pudding. This suggested that electrons were stationary within the atom and did not have any specific arrangement or orbit.

Rutherford’s Nuclear Model of Atom

Rutherford’s gold foil experiment led to the nuclear model of the atom, proposing that an atom consists of a dense, positively charged nucleus surrounded by electrons.

Explanation of Rutherford’s Nuclear Model of Atom

In Rutherford’s model:

• The nucleus contains protons and neutrons, making it the atom’s center of mass.
• Electrons orbit the nucleus, similar to planets orbiting the sun.
• Most of the atom’s volume is in space.

Main Points of Rutherford’s Nuclear Model of Atom

1. Nucleus: A dense core containing protons and neutrons.
2. Electrons: Negatively charged particles orbiting the nucleus.
3. Empty Space: Most of the atom’s volume is empty.

Atomic Number and Mass Number

• Atomic Number (Z): The number of protons in the nucleus of an atom, defining the element.
• Mass Number (A): The total number of protons and neutrons in an atom’s nucleus.

Isobars and Isotopes

• Isobars: Atoms with the same mass number but different atomic numbers (e.g., 1840Ar^{40}_{18}\text{Ar}1840​Ar and 2040Ca^{40}_{20}\text{Ca}2040​Ca).
• Isotopes: Atoms with the same atomic number but different mass numbers (e.g., 612C^{12}_{6}\text{C}612​C, 613C^{13}_{6}\text{C}613​C, and 614C^{14}_{6}\text{C}614​C).

Drawbacks of Rutherford’s Model

While revolutionary, Rutherford’s model had drawbacks:

• It couldn’t explain the stability of atoms, as electrons spiraling into the nucleus should theoretically collapse.
• It didn’t account for the discrete line spectra of elements.

Developments Leading to The Bohr’s Model of Atom

The limitations of Rutherford’s model paved the way for new theories, incorporating quantum concepts and leading to Niels Bohr’s model.

Particle Nature of Electromagnetic Radiation: Planck’s Quantum Theory

Let’s now discuss Plank’s quantum theory covered in the chapter Structure of Atom. Max Planck proposed that electromagnetic energy is quantized and can be emitted or absorbed in discrete quantities called quanta or photons.

Planck’s Quantum Theory II

Photoelectric Effect

Einstein’s explanation of the photoelectric effect showed that light behaves as particles (photons). When light of sufficient energy strikes a metal surface, it emits electrons. This phenomenon confirmed the particle nature of electromagnetic radiation.

Dual Behaviour of Electromagnetic Radiation

Electromagnetic radiation, as stated in the chapter Structure of Atom, exhibits both wave-like and particle-like properties, a concept known as wave-particle duality.

Evidence for The Quantized Electronic Energy Levels: Atomic Spectra

As covered in the chapter structure of atoms, Atoms emit or absorb light at specific wavelengths, forming an atomic spectrum. This spectrum provides evidence for quantized energy levels within an atom.

Emission and Absorption Spectra

• Emission Spectrum: Produced when an electron drops from a higher energy level to a lower one, emitting energy as light.
• Absorption Spectrum: Formed when electrons absorb energy and move to higher energy levels, leaving dark lines in the spectrum.

Line Spectrum of Hydrogen

The hydrogen atom’s line spectrum, with distinct lines in the visible region, led to the understanding that electrons occupy specific energy levels.

Bohr’s Model for Hydrogen Atom

Within the chapter structure of atom, Bohr’s model addressed the shortcomings of Rutherford’s model by introducing quantized orbits for electrons.

Bohr’s Model for Hydrogen Atom

Key postulates of Bohr’s model:

1. Electrons revolve around the nucleus in certain fixed orbits without emitting radiation.
2. Energy is emitted or absorbed only when an electron transitions between these orbits.
3. The angular momentum of an electron in these orbits is quantized and given by mvr=nh2πmvr = n\frac{h}{2\pi}mvr=n2πh​, where nnn is the principal quantum number.

Explanation of Line Spectrum of Hydrogen

Bohr’s model explained the hydrogen atom’s line spectrum by associating spectral lines with electron transitions between quantized orbits.

Limitations of Bohr’s Model

Despite its success, Bohr’s model had limitations:

• It couldn’t explain the spectra of atoms with more than one electron.
• It didn’t consider electron-electron interactions and relativistic effects.

Reasons for The Failure of The Bohr Model

As mentioned in the chapter Structure of Atom, the Bohr model failed for multi-electron systems because:

• It did not incorporate the wave nature of electrons.
• It could not explain the fine structure and splitting of spectral lines.

Towards Quantum Mechanical Model of The Atom

The quantum mechanical model developed from the need to address the limitations of Bohr’s model and incorporate the wave-particle duality of electrons.

Heisenberg’s Uncertainty Principle

Heisenberg’s uncertainty principle states that it is impossible to simultaneously determine the exact position and momentum of an electron.

Significance of the Uncertainty Principle

This principle implies that the concept of definite paths or orbits for electrons, as in the Bohr model, is incorrect.

Example

If the uncertainty in position (Δx\Delta xΔx) is reduced, the uncertainty in momentum (Δp\Delta pΔp) increases, and vice versa.

Quantum Mechanical Model of Atom

The quantum mechanical model describes electrons in terms of probability distributions, rather than fixed orbits, using wave functions.

Significance of Ψ

Ψ (Psi) represents the wave function of an electron, and ∣Ψ∣2|Ψ|^2∣Ψ∣2 gives the probability density of finding an electron in a particular region around the nucleus.

Orbitals and Quantum Numbers

An orbital is a region in space where the probability of finding an electron is highest. Quantum numbers describe the properties of these orbitals.

The Principal Quantum Number (n)

Indicates the main energy level or shell, with values of n=1,2,3,…n = 1, 2, 3, \ldotsn=1,2,3,….

Azimuthal Quantum Number (l)

Defines the shape of the orbital, with values ranging from 000 to n−1n-1n−1.

Magnetic Orbital Quantum Number (ml)

Describes the orientation of the orbital in space, with values ranging from −l-l−l to +l+l+l.

Spin Quantum Number (ms)

Specifies the spin orientation of the electron, with values of ±12\pm \frac{1}{2}±21​.

Shapes of Atomic Orbitals

Atomic orbitals have different shapes based on the azimuthal quantum number:

Shapes of Atomic Orbitals I

• s-orbital: Spherical in shape.
• P-orbital: Dumbbell-shaped.

Shapes of Atomic Orbitals II

• d-orbital: Four-lobed or cloverleaf shape.
• F-orbital: Complex shapes with multiple lobes.

Charge Cloud Picture of Orbitals

The charge cloud model represents the probability distribution of an electron’s position, emphasizing areas of higher probability as denser clouds.

Energies of Orbitals

The energy of an orbital depends on its principal quantum number and azimuthal quantum number.

Energies of Orbitals I

In a single-electron atom, the energy depends only on the principal quantum number.

Energies of Orbitals II

In multi-electron atoms, electron-electron interactions cause energy levels to split further.

Filling of Orbitals in Atom

Electrons fill orbitals in a way that minimizes the atom’s energy, following the Aufbau principle, Pauli exclusion principle, and Hund’s rule.

Hund’s Rule of Maximum Multiplicity

Electrons occupy degenerate orbitals singly before pairing up to maximize the number of unpaired electrons.

Electronic Configuration of Atoms

The chapter structure of atoms also covers the Electronic Configuration of Atoms. The electronic configuration of an atom describes the distribution of electrons among orbitals.

Electronic Configuration of Atoms I

Configuration follows the order of increasing energy levels: 1s,2s,2p,3s,3p,4s,1s, 2s, 2p, 3s, 3p, 4s,1s,2s,2p,3s,3p,4s, etc.

Electronic Configuration of Atoms II

• Example: The electronic configuration of carbon (Z=6Z = 6Z=6) is 1s22s22p21s^2 2s^2 2p^21s22s22p2.

Stability of Completely Filled and Half-Filled Subshells

According to the structure of atoms, Atoms with filled or half-filled subshells exhibit extra stability due to symmetrical distribution and exchange of energy.

Conclusion

This comprehensive guide on “Structure of Atom” provides an in-depth exploration of the fundamental aspects of CBSE Chemistry as outlined in Class 11. It covers the core concepts of atomic theory, including the historical development of atomic models and the current quantum mechanical model. The guide delves into the arrangement of electrons in different orbitals and shells, and how this arrangement affects chemical properties and bonding. It also explores the periodic trends in the structure of atoms and their relevance to the behavior of elements in biological systems. By integrating these concepts with biological applications, the guide helps students understand the role of the structure of atoms in the chemistry of life.

Welcome to iPrep, your Learning Super App. Our learning resources for Chapter 1, “Some Basic Concepts of Chemistry,” in Class 11 Biology provide a solid foundation in essential chemical principles. They include detailed notes on the nature of matter, molecular mass, and the mole concept, helping students understand the fundamental units of chemical measurement. The resources also cover the laws of chemical combination, stoichiometry, and the periodic table, ensuring a thorough grasp of chemical properties and behaviors. Additionally, practice problems and sample questions are included to reinforce learning and prepare students for examinations.

The concept of “Some Basic Concepts of Chemistry” in Class 11 Biology delves into the foundational principles of life by exploring the chemical basis of biological processes. This chapter introduces students to essential topics such as matter and molecular structure, which are crucial for understanding how biological molecules are formed and interact. It also covers the mole concept and stoichiometry, which are key for quantifying substances and reactions in biological systems. Additionally, the chapter discusses the periodic table, providing insights into the properties and behaviors of different elements that are fundamental to life processes. By grasping these basic concepts of chemistry, students gain a deeper understanding of the chemical underpinnings of biological functions.

Development of Chemistry and Its Importance

Chemistry has evolved over centuries from alchemy to a structured scientific discipline. It is essential for understanding biological processes, environmental changes, and material innovations. Chemistry forms the basis of numerous industries such as pharmaceuticals, petrochemicals, and food processing.

Nature of Matter

Matter is anything that has mass and occupies space. It is composed of particles, which can be atoms, molecules, or ions. Understanding the nature of matter is fundamental to the study of chemistry, as it forms the basis for explaining various physical and chemical phenomena.

Physical States of Matter

Matter exists in three primary physical states:

1. Solid: Has a definite shape and volume due to closely packed particles.
2. Liquid: Has a definite volume but takes the shape of its container due to loosely packed particles.
3. Gas: Has neither a definite shape nor volume, as particles are far apart and move freely.

Classification of Matter

Matter can be classified based on its composition and properties into mixtures and pure substances.

Mixtures

A mixture consists of two or more substances physically combined. Mixtures can be classified into:

1. Homogeneous Mixtures: The composition is uniform throughout, and the different components are not visibly distinguishable. Examples include saltwater and air.
2. Heterogeneous Mixtures: The composition is not uniform, and the different components are visibly distinguishable. Examples include a salad and sand in water.

Pure Substances

A pure substance has a uniform composition and distinct properties. Pure substances are further classified into:

1. Elements: Substances that cannot be broken down into simpler substances by chemical means. Examples include oxygen, hydrogen, and iron.
2. Compounds: Substances composed of two or more elements chemically combined in a fixed ratio. Examples include water (H₂O) and carbon dioxide (CO₂).

Properties of Matter

Properties of matter are classified into physical and chemical properties:

• Physical Properties: Characteristics that can be observed or measured without changing the substance’s composition. Examples include color, odor, melting point, boiling point, and density.
• Chemical Properties: Characteristics that describe a substance’s ability to undergo a specific chemical change. Examples include flammability, reactivity with acids, and oxidation states.

Measurement in Chemistry

Accurate measurement is essential in chemistry for quantitative analysis and experiments. The International System of Units (SI) is the standard for measurements.

The International System of Units (SI)

The SI system includes seven base units for fundamental quantities:

Mass and Weight

• Mass: The amount of matter in an object, measured in kilograms (kg).
• Weight: The force exerted by gravity on an object, which varies depending on the gravitational field strength.

Volume and Density

• Volume: The amount of space occupied by a substance, measured in cubic meters (m³) or liters (L).
• Density: The mass of a substance per unit volume, calculated as density = mass/volume. It is expressed in kilograms per cubic meter (kg/m³) or grams per cubic centimeter (g/cm³).

Temperature

Temperature is a measure of the average kinetic energy of particles in a substance. It is measured in Kelvin (K) in the SI system.

Uncertainty in Measurement and Significant Figures

All measurements have some degree of uncertainty due to the limitations of measuring instruments. Significant figures in a measurement include all certain digits plus one uncertain digit.

Rules for Significant Figures:

1. All non-zero digits are significant.
2. Zeros between non-zero digits are significant.
3. Leading zeros are not significant.
4. Trailing zeros in a decimal number are significant.

Calculations Involving Significant Figures

• Addition and Subtraction: The result should have the same number of decimal places as the measurement with the least number of decimal places.
• Multiplication and Division: The result should have the same number of significant figures as the measurement with the least number of significant figures.

Dimensional Analysis

Dimensional analysis, also known as the factor-label method, is used to convert units from one system to another using conversion factors. It ensures that the final answer has the correct units.

Laws of Chemical Combination

Chemistry is governed by several fundamental laws that describe how substances combine and react.

Law of Conservation of Mass

Proposed by Antoine Lavoisier, this law states that mass is neither created nor destroyed in a chemical reaction. The total mass of the reactants equals the total mass of the products.

Law of Definite Proportions

Also known as the Law of Constant Composition, this law states that a chemical compound always contains the same elements in the same proportion by mass, regardless of the sample size or source.

Law of Multiple Proportions

This law states that if two elements combine to form more than one compound, the masses of one element that combine with a fixed mass of the other are in simple whole-number ratios.

Gay-Lussac’s Law of Gaseous Volumes

This law states that when gases react together at constant temperature and pressure, the volumes of the reacting gases and the products (if gaseous) are in whole number ratios.

Avogadro’s Law

Avogadro’s Law states that equal volumes of gases, at the same temperature and pressure, contain an equal number of molecules.

Dalton’s Atomic Theory

Proposed by John Dalton, this theory states that all matter is composed of atoms, which are indivisible and indestructible particles. Atoms of the same element are identical in mass and properties, while atoms of different elements have different masses and properties. Atoms combine in simple whole-number ratios to form compounds.

Atomic Mass and Average Atomic Mass

• Atomic Mass: The mass of a single atom, typically expressed in atomic mass units (amu).
• Average Atomic Mass: The weighted average of the atomic masses of an element’s isotopes, based on their natural abundance.

Molecular Mass and Formula Mass

• Molecular Mass: The sum of the atomic masses of all atoms in a molecule.
• Formula Mass: The sum of the atomic masses of all atoms in a formula unit of a compound.

Percentage Composition

Percentage composition refers to the percentage by mass of each element in a compound. It is calculated using the formula:

Percentage of an element=Mass of the element in 1 mole of the compoundMolar mass of the compound×100\text{Percentage of an element} = \frac{\text{Mass of the element in 1 mole of the compound}}{\text{Molar mass of the compound}} \times 100 Percentage of an element = Molar mass of the compound mass of the element in 1 mole of the compound​×100

Empirical Formula and Molecular Formula

• Empirical Formula: The simplest whole-number ratio of atoms of each element in a compound.
• Molecular Formula: The actual number of atoms of each element in a molecule of the compound.

Stoichiometry and Stoichiometric Calculations

Stoichiometry involves calculating the quantities of reactants and products in chemical reactions based on the balanced chemical equation.

Mole Concept and Molar Masses

• Mole: A mole is the amount of substance that contains as many entities (atoms, molecules, ions) as there are in 12 grams of carbon-12.
• Molar Mass: The mass of one mole of a substance, expressed in grams per mole (g/mol).

Reactions in Solutions

Chemical reactions often occur in solutions, where reactants are dissolved in a solvent. Understanding the concentration of solutions is crucial for stoichiometric calculations.

Terms Expressed Concentrations

• Molarity (M): The number of moles of solute per liter of solution.
• Molality (m): The number of moles of solute per kilogram of solvent.
• Normality (N): The number of equivalents of solute per liter of solution.
• Mass Percent: The mass of solute per mass of solution, expressed as a percentage.

Conclusion

This comprehensive guide on “Some Basic Concepts of Chemistry” provides an in-depth exploration of the fundamental aspects of biology as outlined in Class 11. It covers the core concepts of matter and its nature.. The guide also delves into molecular structure and the nature of chemical bonds, essential for understanding biological macromolecules like proteins, nucleic acids, and carbohydrates. Key topics such as the mole concept, stoichiometry, and chemical reactions are discussed to illustrate how chemical principles apply to biological systems. By connecting these chemical concepts to biological processes, the guide helps students appreciate the integral role of chemistry in life sciences.