Structure of the Atom

Structure of the Atom

In Chapter 3, we learned that atoms and molecules are the basic building blocks of matter. Different kinds of matter exist because they are made up of different atoms. Now, let’s explore:

1. What makes one atom different from another?
2. Are atoms really indivisible, or do they have smaller parts inside them?

We’ll discover answers to these questions in this chapter. We’ll learn about sub-atomic particles and different models explaining their arrangement in an atom.

Charged Particles in Matter

These activities show that rubbing objects together makes them electrically charged. This charge indicates that atoms are divisible and have charged particles.

Discovery of Sub-Atomic Particles

• Electron: J.J. Thomson discovered the electron, a negatively charged particle.
• Proton: E. Goldstein discovered canal rays (positively charged radiations) leading to the discovery of the proton, a positively charged particle.

Characteristics of Electrons and Protons

• Electron (e–):
  • Charge: -1
  • Mass: Negligible
• Proton (p+):
  • Charge: +1
  • Mass: 2000 times that of an electron

The Structure of an Atom

Dalton’s atomic theory suggested atoms were indivisible. However, the discovery of electrons and protons showed this was not true. Scientists proposed different models to explain how these particles are arranged in an atom.

a. Thomson’s Model of an Atom

Thomson’s Christmas Pudding Model

• Analogy: Imagine an atom like a Christmas pudding or a watermelon:
  • Positive charge = pudding or watermelon flesh
  • Electrons = currants or watermelon seeds

Key Points of Thomson’s Model

• An atom has a positively charged sphere with electrons embedded in it.
• The positive and negative charges balance each other, making the atom neutral.

Why It Didn’t Work

• Although it explained electrical neutrality, it couldn’t explain results from later experiments by other scientists.

About J.J. Thomson

• Born: December 18, 1856, in Manchester.
• Discovered electrons.
• Won the Nobel Prize in Physics in 1906.
• Directed Cavendish Laboratory at Cambridge for 35 years.

b. Rutherford’s Model of an Atom

Rutherford’s Experiment

• Purpose: To understand how electrons are arranged.
• Method:
  • Used fast-moving alpha (α)-particles and a very thin gold foil.
  • α-particles are heavy, doubly-charged helium ions.

• Gold Foil Experiment: Fast-moving alpha (α) particles were directed at a thin gold foil.
  • Alpha Particles: Doubly-charged helium ions with considerable energy.
  • Observations:
    • Most alpha particles passed straight through the foil.
    • Some were deflected at small angles.
    • A few bounced back, showing large deflections.

Conclusions

• Most of the atom is empty space.
• Positive charge and most of the atom’s mass are concentrated in a tiny region called the nucleus.
• Electrons revolve around the nucleus in circular paths.
• The nucleus is very small compared to the entire atom.

Drawbacks

• Rutherford’s model suggested that electrons in circular orbits would lose energy and spiral into the nucleus, making atoms unstable. But we know that atoms are stable.

About Ernest Rutherford

• Born: August 30, 1871, at Spring Grove.
• Known as the ‘Father’ of nuclear physics.
• Discovered the atomic nucleus.
• Won the Nobel Prize in Chemistry in 1908.

c. Bohr’s Model of an Atom

Bohr’s Postulates

• Discrete Orbits: Electrons move in specific orbits or shells around the nucleus without radiating energy.
• Energy Levels: These orbits or shells are called energy levels, labeled as K, L, M, N, or numbered as n=1, 2, 3, 4,…

Key Points

• Electrons in these orbits do not lose energy.
• This model explained the stability of atoms and a solution to Rutherford’s model Problems.

About Neils Bohr

• Born: October 7, 1885, in Copenhagen.
• Became a professor of physics at Copenhagen University in 1916.
• Won the Nobel Prize in 1922 for his work on the atomic structure.
• Wrote important books like “The Theory of Spectra and Atomic Constitution,” “Atomic Theory,” and “The Description of Nature.”

Summary

• Thomson’s Model: Positively charged sphere with embedded electrons.
• Rutherford’s Model: Atom mostly empty space, with a dense nucleus and electrons orbiting it.
• Bohr’s Model: Electrons in specific orbits, preventing energy loss and ensuring stability of atoms.

Neutrons

Discovery of Neutrons

• Discovered by J. Chadwick in 1932.
• Neutrons have no charge.
• Mass is nearly equal to that of a proton.
• Represented as ‘n’.

Role in the Atom

• Present in the nucleus of all atoms except hydrogen.
• The mass of an atom is the sum of the masses of protons and neutrons in the nucleus.

Electron Distribution in Different Orbits (Shells)

Bohr and Bury’s Rules

• Rule 1: Maximum number of electrons in a shell = 2n² (where ‘n’ is the shell number).
  • K-shell (1st shell): 2 electrons
  • L-shell (2nd shell): 8 electrons
  • M-shell (3rd shell): 18 electrons
  • N-shell (4th shell): 32 electrons
• Rule 2: Maximum 8 electrons in the outermost shell.
• Rule 3: Electrons fill inner shells first before moving to outer shells.

Valency

Understanding Valency

• Valence Electrons: Electrons in the outermost shell.
• Atoms with 8 valence electrons (octet) are chemically inactive (inert).

Valency Rules

• Octet Rule: Atoms react to achieve 8 electrons in their outermost shell.
• Types of Valency:
  • Losing Electrons: Atoms with 1-3 valence electrons (e.g., Sodium, valency = 1).
  • Gaining Electrons: Atoms with 5-7 valence electrons (e.g., Fluorine, valency = 1).

Examples

• Hydrogen, Lithium, Sodium: 1 valence electron → Valency = 1.
• Magnesium: 2 valence electrons → Valency = 2.
• Aluminium: 3 valence electrons → Valency = 3.
• Fluorine: 7 valence electrons, easier to gain 1 electron → Valency = 1.
• Oxygen: 6 valence electrons, easier to gain 2 electrons → Valency = 2.

Definite Combining Capacity

• Each element has a specific combining capacity known as valency.
• Valency of the first 18 elements can be found in below 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	3,5
Sulphur	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

Atomic Number and Mass Number

Atomic Number (Z)

• Protons in the nucleus determine the atomic number.
• Atomic number (Z) = Number of protons.
• Elements are identified by their atomic number.
  • Example: Hydrogen (Z = 1), Carbon (Z = 6).

Mass Number (A)

• Mass number (A) = Number of protons + Number of neutrons.
• Protons and neutrons are in the nucleus and called nucleons.
  • Example: Carbon has 6 protons and 6 neutrons, so A = 12.
  • Example: Aluminium has 13 protons and 14 neutrons, so A = 27.

Isotopes

Definition

• Atoms of the same element with the same atomic number but different mass numbers.
  • Example: Hydrogen has three isotopes:
    • Protium (¹₁H)
    • Deuterium (²₁H or D)
    • Tritium (³₁H)

Examples

• Carbon: ¹²C and ¹⁴C
• Chlorine: ³⁵Cl and ³⁷Cl

Properties

• Chemical properties are the same.
• Physical properties are different.
• Average atomic mass is calculated based on the isotopes’ masses and their abundance.
  • Example: Chlorine’s average atomic mass is 35.5 u.

Applications of Isotopes

• Uranium isotope in nuclear reactors.
• Cobalt isotope in cancer treatment.
• Iodine isotope in goitre treatment.

Isobars

Definition

• Atoms of different elements with the same mass number but different atomic numbers.
  • Example: Calcium (Z = 20) and Argon (Z = 18) both have a mass number of 40.

Quick Recap

• Atomic Number (Z): Number of protons.
• Mass Number (A): Number of protons + neutrons.
• Isotopes: Same element, different mass numbers.
• Isobars: Different elements, same mass number.

Chapter Summary:

• Credit for the discovery of electron and proton goes to J.J. Thomson and E. Goldstein, respectively.
• J.J. Thomson proposed that electrons are embedded in a positive sphere.
• Rutherford’s alpha-particle scattering experiment led to the discovery of the atomic nucleus.
• Rutherford’s model of the atom proposed:
  • A very tiny nucleus is present inside the atom.
  • Electrons revolve around this nucleus.
  • The stability of the atom could not be explained by this model.
• Neils Bohr’s model of the atom was more successful:
  • Electrons are distributed in different shells with discrete energy around the nucleus.
  • If the atomic shells are complete, the atom will be stable and less reactive.
• J. Chadwick discovered the presence of neutrons in the nucleus of an atom:
  • The three sub-atomic particles of an atom are:
    1. Electrons (negatively charged)
    2. Protons (positively charged)
    3. Neutrons (no charge)
  • The mass of an electron is about 1/2000 times the mass of a hydrogen atom.
  • The mass of a proton and a neutron is taken as one unit each.
• Shells of an atom are designated as K, L, M, N, etc.
• Valency is the combining capacity of an atom.
• The atomic number of an element is the same as the number of protons in the nucleus of its atom.
• The mass number of an atom is equal to the number of nucleons in its nucleus.
• Isotopes are atoms of the same element, which have different mass numbers.
• Isobars are atoms having the same mass number but different atomic numbers.
• Elements are defined by the number of protons they possess.