Structure of the Atom

Sub-Atomic Particles and Atomic Models

Introduction

  • Atoms and molecules are fundamental building blocks of matter.
  • Different kinds of matter are due to different atoms.
  • Questions to be addressed:
    • What differentiates atoms of different elements?
    • Are atoms indivisible as Dalton proposed?
  • This chapter explores sub-atomic particles and their arrangement within the atom.
  • Late 19th-century challenge: to reveal the structure of the atom and explain its properties.
  • Elucidation of atomic structure is based on experiments.
  • Early indication that atoms are not indivisible comes from studying static electricity and electrical conductivity.

Charged Particles in Matter

  • Activities demonstrating static electricity:
    • Combing dry hair and attracting paper pieces.
    • Rubbing a glass rod with silk and bringing it near a balloon.
  • Rubbing objects together results in electrical charges.
  • Atoms are divisible and consist of charged particles.
  • Scientists who contributed to revealing charged particles.
  • By 1900, it was known that atoms contain at least one sub-atomic particle - the electron (J.J. Thomson).
  • E. Goldstein (1886) discovered canal rays in a gas discharge.
    • These rays are positively charged radiations.
    • Led to the discovery of the proton, another sub-atomic particle.
    • Proton charge: equal in magnitude but opposite in sign to the electron.
    • Proton mass: approximately 2000 times that of the electron.
  • Representation:
    • Electron: e-
    • Proton: p+
  • Mass and charge units:
    • Proton mass: 1 unit, charge: +1
    • Electron mass: negligible, charge: -1
  • Atoms consist of protons and electrons, balancing charges.
  • Protons are located in the interior of the atom, making them difficult to remove compared to electrons.

The Structure of an Atom

  • Dalton’s atomic theory (Chapter 3) suggested indivisibility and indestructibility of atoms.
  • Discovery of electrons and protons led to the failure of Dalton’s theory.
  • It became necessary to understand how electrons and protons are arranged within an atom.
  • Many scientists proposed atomic models.
Thomson’s Model of an Atom
  • J.J. Thomson proposed the first atomic model.
  • Model: similar to a Christmas pudding or watermelon.
    • Electrons are like currants in a sphere of positive charge.
    • Positive charge spread like the red edible part of a watermelon.
    • Electrons studded in the positively charged sphere, like seeds.
  • Thomson’s postulates:
    • Atom consists of a positively charged sphere with electrons embedded in it.
    • Negative and positive charges are equal in magnitude, making the atom electrically neutral.
  • Thomson's model explained the electrical neutrality of atoms.
  • However, experiments by other scientists could not be explained by this model.
Rutherford’s Model of an Atom
  • Ernest Rutherford sought to understand electron arrangement within the atom.
  • Designed an experiment involving bombarding a thin gold foil with fast-moving alpha ($\alpha$)-particles.
    • Gold foil was chosen for its thinness (approximately 1000 atoms thick).
    • $\alpha$-particles are doubly-charged helium ions with a mass of 4 u and considerable energy.
    • Expected outcome: $\alpha$-particles would be deflected by sub-atomic particles in gold atoms, but not largely due to their higher mass.
Alpha-Particle Scattering Experiment
  • Unexpected results from the α-particle scattering experiment:
    • Most α-particles passed straight through the gold foil.
    • Some α-particles were deflected by the foil at small angles.
    • One out of every 12000 particles appeared to rebound.
  • Rutherford’s interpretation: "This result was almost as incredible as if you fire a 15-inch shell at a piece of tissue paper and it comes back and hits you".
  • Analogy to understand the experiment:
    • A child throwing stones at a wall versus a barbed-wire fence.
    • Stones at a wall: sound heard each time.
    • Stones at a fence: most stones pass through gaps without hitting, thus no sound.
  • Conclusions from the α-particle scattering experiment:
    • Most of the space inside the atom is empty because most α-particles passed through without deflection.
    • Very few particles were deflected, indicating that the positive charge of the atom occupies very little space.
    • A very small fraction of α-particles were deflected by 1800180^0, indicating that all the positive charge and mass of the gold atom were concentrated in a very small volume within the atom.
    • Radius of the nucleus is about 10510^5 times less than the radius of the atom.
Rutherford’s Nuclear Model of an Atom
  • Features:
    • Positively charged center in an atom called the nucleus.
    • Nearly all the mass of an atom resides in the nucleus.
    • Electrons revolve around the nucleus in circular paths.
    • The size of the nucleus is very small compared to the size of the atom.
Drawbacks of Rutherford’s Model
  • Revolution of electrons in a circular orbit is not expected to be stable.
  • Particles in circular orbits undergo acceleration.
  • During acceleration, charged particles radiate energy.
  • Revolving electrons would lose energy and eventually fall into the nucleus.
  • This would make the atom highly unstable, contradicting the stability of matter.
  • Atoms are known to be quite stable.

Bohr’s Model of Atom

  • Neils Bohr proposed postulates to address objections to Rutherford’s model:
    • Only certain special orbits, known as discrete orbits of electrons, are allowed inside the atom.
    • While revolving in discrete orbits, electrons do not radiate energy.
Energy Levels
  • These orbits or shells are called energy levels.
  • Energy levels are designated as K, L, M, N,… or n=1, 2, 3, 4,…

Neutrons

  • In 1932, J. Chadwick discovered the neutron.
    • A sub-atomic particle with no charge and a mass nearly equal to that of a proton.
  • Neutrons are present in the nucleus of all atoms, except hydrogen.
  • Neutron representation: ‘n’
  • The mass of an atom is given by the sum of the masses of protons and neutrons in the nucleus.

Distribution of Electrons in Different Orbits (Shells)

  • Bohr and Bury suggested the distribution of electrons into different orbits.
Rules for Electron Distribution
  • The maximum number of electrons present in a shell is given by the formula 2n22n^2, where ‘n’ is the orbit number or energy level index (1, 2, 3,…).
    • First orbit or K-shell: 2×12=22 \times 1^2 = 2
    • Second orbit or L-shell: 2×22=82 \times 2^2 = 8
    • Third orbit or M-shell: 2×32=182 \times 3^2 = 18
    • Fourth orbit or N-shell: 2×42=322 \times 4^2 = 32
  • The maximum number of electrons in the outermost orbit is 8.
  • Electrons are not accommodated in a given shell unless the inner shells are filled stepwise.

Valency

  • Electrons present in the outermost shell of an atom are valence electrons.
  • The outermost shell can accommodate a maximum of 8 electrons.
  • Atoms with 8 electrons in the outermost shell show little chemical activity (inert elements).
  • Inert elements: helium (2 electrons) and other elements (8 electrons in the outermost shell).
  • Combining capacity (valency): tendency to react and form molecules to attain a fully-filled outermost shell (octet).
  • Atoms react by sharing, gaining, or losing electrons to achieve an octet.
  • Valency is the number of electrons gained, lost, or shared to make an octet in the outermost shell.
  • Examples:
    • Hydrogen, lithium, sodium: 1 electron in the outermost shell, valency of 1 (lose 1 electron).
    • Magnesium: 2 electrons in the outermost shell, valency of 2.
    • Aluminum: 3 electrons in the outermost shell, valency of 3.
  • If the number of electrons in the outermost shell is close to its full capacity:
    • Valency is determined differently.
    • Fluorine: 7 electrons in the outermost shell, easier to gain 1 electron than lose 7; valency = 1 (8 - 7).
    • Oxygen: valency calculated similarly.
  • Each element has a definite combining capacity called valency.

Atomic Number and Mass Number

Atomic Number
  • Protons in the nucleus determine the atomic number.
  • It is denoted by ‘Z’.
  • All atoms of an element have the same atomic number.
  • Elements are defined by the number of protons.
    • Hydrogen: Z = 1 (1 proton).
    • Carbon: Z = 6.
  • The atomic number is the total number of protons in the nucleus of an atom.
Mass Number
  • The mass of an atom is due to protons and neutrons (nucleons) in the nucleus.
  • Mass resides in the nucleus.
  • Carbon mass: 12 u (6 protons + 6 neutrons).
  • Aluminum mass: 27 u (13 protons + 14 neutrons).
  • The mass number is the sum of protons and neutrons in the nucleus, denoted by ‘A’.
  • Notation for an atom:

AZX\frac{A}{Z}X

  • Example: Nitrogen is written as 147N\frac{14}{7}N

Isotopes

  • Atoms of some elements have the same atomic number but different mass numbers.
  • Hydrogen isotopes:
    • Protium: 11H\frac{1}{1}H
    • Deuterium: 21H\frac{2}{1}H or D
    • Tritium: 31H\frac{3}{1}H or T
  • Each has an atomic number of 1, but mass numbers of 1, 2, and 3, respectively.
  • Other examples:
    • Carbon: 126C\frac{12}{6}C and 146C\frac{14}{6}C
    • Chlorine: 3517Cl\frac{35}{17}Cl and 3717Cl\frac{37}{17}Cl
  • Isotopes are atoms of the same element with the same atomic number but different mass numbers.
  • Three isotopes of hydrogen: protium, deuterium, tritium.

Average Atomic Mass

  • Many elements consist of a mixture of isotopes.
  • Each isotope of an element is a pure substance.
  • Isotopes have similar chemical properties but different physical properties.
  • Chlorine occurs in two isotopic forms with masses 35 u and 37 u in a 3:1 ratio.
  • The average atomic mass of chlorine is calculated as follows:

(35×75100)+(37×25100)=[(1054)+(374)]=1424=35.5 u(\frac{35 \times 75}{100}) + (\frac{37 \times 25}{100}) = [(\frac{105}{4}) + (\frac{37}{4})] = \frac{142}{4} = 35.5 \text{ u}

  • The mass of a natural element is the average mass of all its naturally occurring atoms.
  • If an element has no isotopes, its mass is the sum of protons and neutrons.
  • If an element occurs in isotopic forms, the percentage of each isotopic form must be known to calculate the average mass.
  • This does not mean that one atom of chlorine has a fractional mass of 35.5 u.
  • It means that a certain amount of chlorine will contain both isotopes, and the average mass is 35.5 u.

Applications of Isotopes

  • The chemical properties of all isotopes of an element are the same.
  • Some isotopes have special properties useful in various fields:
    • Uranium isotope: fuel in nuclear reactors.
    • Cobalt isotope: treatment of cancer.
    • Iodine isotope: treatment of goiter.

Isobars

  • Calcium (atomic number 20) and argon (atomic number 18) have different numbers of protons but the same mass number of 40.
  • Isobars are atoms of different elements with different atomic numbers but the same mass number.