Chapter 4: Structure Of Atom

Structure of the Atom

Key Questions Addressed in the Chapter:

1. What makes atoms of one element different from another?

2. Are atoms truly indivisible, or do they have smaller constituents?

Learning Objectives:

• Understand sub-atomic particles.

• Explore models of atomic structure.

Historical Context:

• By the 19th century, scientists aimed to:

  • Reveal the structure of the atom.

  • Explain its key properties.

Insights into the Atom's Divisibility:

• Initial indication of atoms being divisible emerged from studies on:

  • Static electricity.

  • Conditions under which substances conduct electricity.

Significance:

• Series of experiments led to understanding:

  • Sub-atomic particles.

  • Their arrangement within the atom.

Charged Particles in Matter

Key Concepts:

1. Rubbing and Electric Charge:

   • Rubbing two objects (e.g., comb and hair, glass rod and silk) can make them electrically charged.

   • This charge arises from the divisibility of atoms into smaller charged particles.

Discovery of Sub-atomic Particles:

• Electron:

  • Identified by J.J. Thomson (1900).

  • Negatively charged, represented as e⁻.

  • Charge: −1, Mass: negligible.

• Proton:

  • Discovered by E. Goldstein (1886) through canal rays (positively charged radiations).

  • Charge: +1+1, Mass: approximately 2000 times the mass of an electron.

  • Represented as p⁺.

Canal Rays:

• Positively charged radiations observed in gas discharge experiments.

• Led to the discovery of protons.

Neutral Atoms:

• An atom with one proton (p+) and one electron (e⁻) is electrically neutral.

• Charges balance out: +1+(−1)=0

Structure of the Atom:

• Electrons are easier to remove compared to protons, suggesting protons are located in the interior of the atom.

• The arrangement of protons and electrons in the atom remained an open question for further exploration.

Structure of an Atom

Background:

• Dalton’s atomic theory proposed atoms were indivisible and indestructible.

• Discovery of electrons and protons disproved this aspect of Dalton's theory.

Need for Atomic Models:

• Scientists sought to explain the arrangement of electrons and protons within the atom.

• J.J. Thomson proposed the first model of atomic structure.

Thomson’s Model of an Atom

Key Features:

1. Positive Sphere with Embedded Electrons:

   • Atom resembles a Christmas pudding or a watermelon:

     • Positive charge: Spread throughout the sphere (like the red part of a watermelon).

     • Electrons: Embedded within the positive sphere (like watermelon seeds).

2. Electrically Neutral Atom:

   • Positive and negative charges are equal in magnitude, making the atom neutral as a whole.

Strengths:

• Explained why atoms are electrically neutral.

Limitations:

• Thomson's model failed to explain the results of experiments conducted by other scientists.

• Required further refinement and alternative models to understand the atom’s structure.

J.J. Thomson

Personal Background:

• Name: Joseph John Thomson.

• Born: December 18, 1856, in Cheetham Hill, Manchester, England.

• Profession: British physicist.

Major Achievements:

1. Discovery of the Electron:

   • Pioneered work on sub-atomic particles.

   • Awarded the Nobel Prize in Physics (1906) for the discovery of the electron.

2. Role as a Mentor:

   • Directed the Cavendish Laboratory at Cambridge for 35 years.

   • Mentored seven research assistants who later won Nobel Prizes.

Legacy:

• Significant contributions to understanding atomic structure.

• Proposed the plum pudding model of the atom.

• His work laid the foundation for modern atomic and particle physics.

Rutherford’s Model of an Atom

Background:

• Ernest Rutherford designed an experiment to study the arrangement of electrons in an atom.

• The experiment involved bombarding a thin gold foil with fast-moving alpha (α\alpha) particles.

The Alpha-Particle Scattering Experiment

Key Features of the Experiment:

1. Gold Foil Selection:

   • A very thin layer of gold foil (1000 atoms thick) was used.

   • Gold was chosen because it can be hammered into extremely thin sheets.

2. Alpha (α\alpha) Particles:

   • Doubly charged helium ions (He²⁺).

   • Have a mass of 4u and considerable energy.

3. Expectations:

   • α\alpha-particles were expected to pass through the foil with minimal deflection, as they are much heavier than protons.

Observations:

1. Most particles passed straight through the gold foil:

   • Indicates that most of the atom is empty space.

2. Some particles were deflected by small angles:

   • Suggests the presence of a concentrated positive charge in the atom.

3. A very small fraction (1 in 12,000) rebounded at large angles (180°):

   • Indicates that nearly all the mass and positive charge of the atom are concentrated in a very small region.

Implications:

1. Empty Space in Atoms:

   • Most of the atom is empty space, allowing particles to pass through.

2. Concentration of Positive Charge and Mass:

   • Positive charge and mass are concentrated in a small central region called the nucleus.

3. Relative Size of the Nucleus:

   • The radius of the nucleus is approximately 10510^5 times smaller than the radius of the atom.

Rutherford’s Nuclear Model of the Atom

Key Features:

1. Nucleus:

   • A small, dense, positively charged center.

   • Contains nearly all the mass of the atom.

2. Electrons:

   • Electrons revolve around the nucleus in circular orbits.

3. Size of the Nucleus:

   • The nucleus is very small compared to the size of the atom.

Drawbacks of Rutherford’s Model:

1. Instability of Revolving Electrons:

   • According to classical physics, an accelerating charged particle (electron) would radiate energy.

   • Continuous energy loss would cause the electron to spiral into the nucleus, making the atom unstable.

2. Matter Stability:

   • Rutherford’s model could not explain why atoms are stable despite the predicted energy loss of revolving electrons.

   • This contradiction indicated the need for a new atomic model to explain atomic stability.

Significance:

• Rutherford’s model introduced the concept of the nucleus, a fundamental breakthrough in understanding atomic structure.

• However, it failed to explain the stability of atoms, paving the way for Bohr’s atomic model.

Ernest Rutherford

Personal Background:

• Name: Ernest Rutherford.

• Born: August 30, 1871, in Spring Grove, New Zealand.

• Known As: "Father of Nuclear Physics."

Major Contributions:

1. Discovery of the Nucleus:

   • Conducted the gold foil experiment, leading to the discovery of the atomic nucleus.

2. Work on Radioactivity:

   • Made significant contributions to understanding radioactive decay.

   • Coined terms like alpha and beta radiation.

3. Rutherford’s Atomic Model:

   • Proposed the nuclear model of the atom, introducing the concept of a dense, positively charged nucleus.

Achievements:

• Awarded the Nobel Prize in Chemistry (1908) for his investigations into the chemistry of radioactive substances.

Legacy:

• Pioneered nuclear physics and inspired future research in atomic structure.

• His discoveries laid the foundation for modern physics, including the development of quantum mechanics and nuclear energy.

 Bohr’s Model of the Atom

Background:

• Bohr’s model was proposed to address the limitations of Rutherford’s atomic model, specifically the instability of revolving electrons.

Postulates of Bohr’s Atomic Model:

1. Discrete Orbits of Electrons:

   • Electrons revolve around the nucleus in specific, well-defined orbits called discrete orbits.

   • These orbits are also known as energy levels.

2. Energy Conservation in Orbits:

   • While in these discrete orbits, electrons do not radiate energy.

   • This ensures the stability of the atom.

Energy Levels in an Atom:

• Represented by letters: K, L, M, N, …

  • K-shell: n=1

  • L-shell: n=2

  • M-shell: n=3

  • N-shell: n=4, and so on.

• Alternatively represented by numbers: n=1,2,3,4,…n = 1, 2, 3, 4, \dots.

Significance of Bohr’s Model:

1. Explained atomic stability by preventing energy loss during electron motion.

2. Paved the way for understanding atomic spectra and quantum mechanics.

3. Successfully explained the spectral lines of hydrogen.

Limitations:

• Could not explain spectral lines of atoms with more than one electron.

• Did not account for the finer details of atomic spectra or the effects of magnetic fields (Zeeman effect).

Diagram Reference:

• Energy levels are represented as concentric circles around the nucleus.

• Each orbit corresponds to a specific energy level.

Neils Bohr: 

Personal Background:

• Name: Neils Bohr.

• Born: October 7, 1885, in Copenhagen, Denmark.

• Profession: Physicist.

Major Contributions:

1. Atomic Structure:

   • Proposed the Bohr Model of the Atom to address the limitations of Rutherford’s model.

   • Explained the concept of discrete energy levels and stability of electrons in orbits.

2. Work on Atomic Spectra:

   • His model successfully explained the spectral lines of hydrogen, a significant achievement in atomic physics.

Achievements:

• Appointed Professor of Physics at Copenhagen University in 1916.

• Awarded the Nobel Prize in Physics (1922) for his contributions to the understanding of atomic structure and quantum mechanics.

Notable Writings:

1. The Theory of Spectra and Atomic Constitution.

2. Atomic Theory.

3. The Description of Nature.

Legacy:

• Bohr’s work laid the foundation for modern quantum mechanics.

• His insights into atomic behavior continue to influence physics and chemistry.

Neutrons: 

Discovery:

• Year: 1932.

• Scientist: J. Chadwick.

Characteristics of Neutrons:

1. Charge:

   • Neutrons are electrically neutral (no charge).

2. Mass:

   • Mass is nearly equal to that of a proton.

   • Represented as ‘n’.

3. Location:

   • Neutrons are present in the nucleus of an atom.

   • Exception: Hydrogen atoms, which typically lack neutrons.

Role in the Atom:

• Neutrons, along with protons, contribute to the atomic mass.

• Atomic Mass = Mass of Protons + Mass of Neutrons.

Significance:

• Neutrons help stabilize the nucleus by reducing electrostatic repulsion between positively charged protons.

• Essential for nuclear reactions and radioactive decay.

Practice Question 👈click it 

Electron Distribution in Different Orbits (Shells)

Rules for Electron Distribution:

1. Maximum Number of Electrons in a Shell:

   • The maximum number of electrons that can occupy a shell is given by the formula: Maximum electrons=2n²where n is the orbit number (energy level index).

     • For each shell, the maximum number of electrons is calculated as follows:

       • K-shell (n=1): 2×1² = 2 electrons.

       • L-shell (n=2): 2×2²= 8 electrons.

       • M-shell (n=3): 2×3² = 18 electrons.

       • N-shell (n=4): 2×4² = 32 electrons, and so on.

2. Outer Shell Limitation:

   • The outermost shell of an atom can hold a maximum of 8 electrons, irrespective of the formula. This is a key factor in determining the chemical behavior of elements.

3. Filling of Shells:

   • Electrons are placed in shells in a step-wise manner.

   • Inner shells must be completely filled before electrons are placed in outer shells.

   • This rule ensures that the lower energy levels are filled first, following the principle of energy minimization.

Summary of Electron Distribution in Shells:

• K-shell (n=1): 2 electrons

• L-shell (n=2): 8 electrons

• M-shell (n=3): 18 electrons

• N-shell (n=4): 32 electrons

• The maximum for the outermost shell is 8 electrons, regardless of the shell number.

These rules help determine the electron configuration of atoms, which plays a crucial role in their chemical properties.

Valency:

Definition:

• Valency is the combining capacity of an atom, defined by its ability to form bonds with other atoms. This capacity is determined by the number of electrons an atom can gain, lose, or share to achieve a stable outer electron configuration, usually with 8 electrons in the outermost shell (an octet), or 2 electrons in the case of hydrogen and helium.

Outer Electrons (Valence Electrons):

• The electrons present in the outermost shell of an atom are called valence electrons. The number of valence electrons influences the chemical properties and reactivity of the element.

Bohr-Bury Scheme:

• The outermost shell can hold a maximum of 8 electrons (except hydrogen and helium, which can hold 2 electrons). Atoms tend to react to fill their outermost shell to achieve a stable electron configuration.

Inert Elements:

• Elements that have their outermost shell completely filled with 8 electrons are inert (such as neon and helium). These elements show little chemical reactivity, meaning their valency is zero.

Formation of Molecules and Reactivity:

• Elements tend to react and combine with other atoms in order to attain a stable outer electron shell (octet). This can happen by sharing, gaining, or losing electrons.

• Example: Hydrogen, lithium, and sodium have 1 electron in their outermost shell. Each can lose 1 electron to achieve a stable configuration, resulting in a valency of 1.

Valency of Elements:

• Magnesium (Mg): Has 2 electrons in the outer shell. It can lose 2 electrons to achieve a stable configuration, so its valency is 2.

• Aluminium (Al): Has 3 electrons in the outer shell. It can lose 3 electrons to achieve a stable configuration, so its valency is 3.

Determining Valency When Electrons Are Close to Full:

• Elements with a nearly full outer shell will typically gain electrons rather than lose them to achieve a stable octet.

• Example: Fluorine (F) has 7 electrons in the outer shell. It is easier for fluorine to gain 1 electron to complete its octet, so the valency of fluorine is 1.

• Similarly, oxygen (O) has 6 electrons in its outer shell. Oxygen can gain 2 electrons to complete the octet, so the valency of oxygen is 2.

Electron Distribution in First 18 Elements (Table 4.1)

Element Information Table

Element	Symbol	Atomic Number	Protons	Neutrons	Electrons	Electron Distribution	Valency
Hydrogen	H	1	1	0	1	K = 1	1
Helium	He	2	2	2	2	K = 2	0
Lithium	Li	3	3	4	3	K = 2, L = 1	1
Beryllium	Be	4	4	5	4	K = 2, L = 2	2
Boron	B	5	5	6	5	K = 2, L = 3	3
Carbon	C	6	6	6	6	K = 2, L = 4	4
Nitrogen	N	7	7	7	7	K = 2, L = 5	3
Oxygen	O	8	8	8	8	K = 2, L = 6	2
Fluorine	F	9	9	10	9	K = 2, L = 7	1
Neon	Ne	10	10	10	10	K = 2, L = 8	0
Sodium	Na	11	11	12	11	K = 2, L = 8, M = 1	1
Magnesium	Mg	12	12	12	12	K = 2, L = 8, M = 2	2
Aluminium	Al	13	13	14	13	K = 2, L = 8, M = 3	3
Silicon	Si	14	14	14	14	K = 2, L = 8, M = 4	4
Phosphorus	P	15	15	16	15	K = 2, L = 8, M = 5	3 or 5
Sulfur	S	16	16	16	16	K = 2, L = 8, M = 6	2
Chlorine	Cl	17	17	18	17	K = 2, L = 8, M = 7	1
Argon	Ar	18	18	22	18	K = 2, L = 8, M = 8	0

Summary:

• Valency is an atom’s ability to combine with other atoms to form molecules by either gaining, losing, or sharing electrons.

• Atoms react to achieve a stable outer shell, typically with 8 electrons (octet rule).

• Valency can be easily determined by considering the number of electrons in the outermost shell and how an atom can either lose, gain, or share electrons to fill it.

• Inert gases (like neon and helium) have a valency of 0 because they already have full outer shells.

• Elements with fewer electrons in their outer shell tend to lose electrons, while those with more electrons tend to gain electrons to form a stable octet.

These concepts form the foundation for understanding chemical bonding and reactions.

Atomic Number:

Definition:

• The atomic number of an element is the number of protons in the nucleus of an atom. It is denoted by the symbol Z.

Key Points:

1. Determines Identity of Elements:

   • The atomic number is unique to each element. It is this number of protons that defines the element.

   • For example:

     • Hydrogen (H) has 1 proton, so its atomic number is Z = 1.

     • Carbon (C) has 6 protons, so its atomic number is Z = 6.

2. Atomic Number and Electron Count:

   • In a neutral atom (no charge), the number of protons is equal to the number of electrons.

   • For instance, a hydrogen atom (Z = 1) has 1 proton and 1 electron, and a carbon atom (Z = 6) has 6 protons and 6 electrons.

3. Significance:

   • The atomic number determines the chemical properties of an element, as it influences how the electrons are arranged in the atom’s electron shells.

   • The periodic table of elements is arranged in order of increasing atomic number. This arrangement reflects the increasing number of protons in the nucleus, which leads to a systematic progression in the chemical properties of elements.

Summary:

• The atomic number (Z) of an element is the total number of protons in the nucleus of its atom.

• It uniquely identifies an element and determines its position in the periodic table.

• For neutral atoms, the atomic number also tells us the number of electrons in the atom.

Mass Number: Detailed Notes

Definition:

• The mass number (denoted by A) is the total number of protons and neutrons in the nucleus of an atom. It is a measure of the atomic mass of the atom.

Key Points:

1. Mass of an Atom:

   • The mass of an atom is essentially due to the mass of the protons and neutrons, as the electrons have negligible mass.

   • Protons and neutrons are collectively called nucleons, as they are found in the nucleus of an atom.

2. Formula for Mass Number:

   • The mass number (A) is calculated as: A=Number of Protons+Number of Neutrons

   • Example:

     • For Carbon (C), which has 6 protons and 6 neutrons: A=6+6=12

     • Therefore, the mass number of carbon is 12 u (atomic mass unit).

     • For Aluminium (Al), which has 13 protons and 14 neutrons: A=13+14=27

     • Therefore, the mass number of aluminium is 27 u.

3. Atomic Number and Mass Number:

   • The atomic number (Z) tells us the number of protons in the nucleus.

   • The mass number (A) gives the total number of protons and neutrons in the nucleus.

4. Notation for an Atom:

   • An atom is often represented using the following notation: ^A _ZSymbol of element

     • Where:

       • A is the mass number,

       • Z is the atomic number,

       • The element's symbol represents the element.

     • Example: For nitrogen (which has 7 protons and 7 neutrons), the notation is written as:This shows that nitrogen has an atomic number of 7 and a mass number of 14.

Summary:

• The mass number (A) of an atom is the sum of the number of protons and neutrons in the nucleus.

• It is a key property that helps determine the mass of an atom and is written in the atomic notation alongside the atomic number and the element symbol.

• Mass Number is used to distinguish between isotopes of the same element, as isotopes have the same atomic number but different mass numbers.

Isotopes: Detailed Notes

Definition:

• Isotopes are atoms of the same element that have the same atomic number but different mass numbers. This means they have the same number of protons but a different number of neutrons in their nuclei.

Key Points:

1. Same Atomic Number, Different Mass Numbers:

   • Isotopes of an element have the same number of protons but a different number of neutrons, leading to different mass numbers.

   • Example 1: Hydrogen has three isotopes:

     • Protium (¹H): 1 proton, 0 neutrons (Mass number = 1)

     • Deuterium (²H or D): 1 proton, 1 neutron (Mass number = 2)

     • Tritium (³H or T): 1 proton, 2 neutrons (Mass number = 3)

   • Example 2: Carbon has two common isotopes:

     • Carbon-12 (¹²C): 6 protons, 6 neutrons (Mass number = 12)

     • Carbon-14 (¹⁴C): 6 protons, 8 neutrons (Mass number = 14)

   • Example 3: Chlorine has two isotopes:

     • Chlorine-35 (³⁵Cl): 17 protons, 18 neutrons

     • Chlorine-37 (³⁷Cl): 17 protons, 20 neutrons

2. Chemical and Physical Properties:

   • Chemical Properties: The chemical properties of isotopes are generally the same because they have the same number of electrons and the same electron configuration.

   • Physical Properties: The physical properties (such as mass, density, and rate of diffusion) of isotopes are different due to the difference in mass.

3. Natural Occurrence of Isotopes:

   • Many elements in nature consist of a mixture of isotopes. The ratio of different isotopes can vary, affecting the average atomic mass of the element.

   • Example:

     • Chlorine: In nature, chlorine exists as two isotopes, Cl-35 and Cl-37, in the ratio of 3:1. The average atomic mass of chlorine is calculated by considering the abundance of each isotope:

       

     • Note: This does not mean that each chlorine atom has a mass of 35.5 u, but rather that the average mass of chlorine atoms in nature is 35.5 u due to the presence of both isotopes.

4. Isotopic Mass and Atomic Mass:

   • The atomic mass of an element is typically the weighted average of the masses of its naturally occurring isotopes, taking into account their relative abundances.

   • If an element has no isotopes, its atomic mass is simply the sum of the protons and neutrons in its nucleus.

5. Applications of Isotopes:

   • Some isotopes have unique properties that make them useful in various fields:

     • Uranium Isotope (U-235): Used as fuel in nuclear reactors.

     • Cobalt Isotope (Co-60): Used in cancer treatment, particularly in radiation therapy.

     • Iodine Isotope (I-131): Used in the treatment of goitre, a thyroid condition.

Summary:

• Isotopes are atoms of the same element with the same number of protons but different numbers of neutrons, leading to different mass numbers.

• Although isotopes of an element have similar chemical properties, they may differ in physical properties.

• The average atomic mass of an element is determined by the weighted average of the masses of its isotopes, taking into account their natural abundances.

• Isotopes have various practical applications in medicine, industry, and research due to their unique properties.

Isobars:

Definition:

• Isobars are atoms of different elements that have the same mass number but different atomic numbers. This means that while the total number of protons and neutrons (nucleons) is the same, the number of protons (atomic number) is different.

Key Points:

1. Same Mass Number, Different Atomic Numbers:

   • Isobars have the same total number of nucleons (protons + neutrons), but a different number of protons.

   • This leads to different atomic numbers (Z), meaning the elements involved are different.

2. Example of Isobars:

   • Consider the elements Calcium (Ca) and Argon (Ar):

     • Calcium (Ca):

       • Atomic Number (Z) = 20 (20 protons)

       • Mass Number (A) = 40 (20 protons + 20 neutrons)

     • Argon (Ar):

       • Atomic Number (Z) = 18 (18 protons)

       • Mass Number (A) = 40 (18 protons + 22 neutrons)

     • Both Calcium and Argon have the same mass number (40), but different atomic numbers (20 and 18, respectively), making them isobars.

3. Difference in Atomic Number:

   • The different atomic numbers of isobars indicate that they belong to different elements, and thus, they have different chemical properties.

4. Isobar Examples:

   • Other examples of isobars include:

     • Argon (Ar) with atomic number 18 and Calcium (Ca) with atomic number 20, both having a mass number of 40.

     • Carbon-14 (¹⁴C) and Nitrogen-14 (¹⁴N) are also isobars because they both have a mass number of 14, but carbon has 6 protons and nitrogen has 7 protons.

Summary:

• Isobars are atoms of different elements that have the same mass number but different atomic numbers.

• The number of nucleons (protons + neutrons) is the same in isobars, but their atomic numbers (protons) differ, which means they are different elements with distinct chemical properties.

Practice Question 👈click it