Structure of the Atom Summary Class 9 Science Chapter 4

Introduction

• John Dalton considered atom to be an indivisible entity, but his concept had to be discarded at the end of nineteenth century, when scientists through experiments were able to find existence of charged (electrons and protons) and neutral particles (neutrons) in the atom. These particles were called the ‘Sub-atomic Particles’.

Discovery of Electrons – Cathode Rays (By J. J. Thomson)

• Thomson explained presence of electrons by cathode rays experiment.

Facts about Electrons

• Charge on electron = −1.6 × 10^-19 C (C = Coloumb)
  (As calculated by Robert E. Millikan)
• Mass of electron = 9.1 × 10-31 kg

Discovery of Protons – Anode Rays/Canal Rays (By E. Goldstein)

• E. Goldstein by his famous anode rays/canal rays experiment was able to detect presence of positively charged particles called protons in the atom.

Facts about Protons

• Charge on proton = + 1.6 × 10-19 C
• Mass of proton = 1.673 × 10-24 gm
  i.e.,  Mass of proton ≅ 1840 × Mass of electron

Discovery of Neutrons (By J. Chadwick)

• J. Chadwick bombarded lighter elements (like lithium, boron etc.) with α-particles and observed emission of new particles having zero charge but having mass equal to that of proton.
• These particles were called ‘Neutron’ i.e., neutral particle of the atom.
• Neutron are absent in Protium isotope of hydrogen atom.(1H1)
• Since, mass of electrons are negligible as compared to that of proton and neutrons hence, sum of masses of protons and neutrons in an atom will compose its atomic mass.

Atomic Models

• From the knowledge of existence of subatomic particles like electron, proton and neutron in an atom, various atomic models were proposed by different scientists.

Some of the atomic models:

• Thomson’s Model of Atom
• Rutherford’s Model of Atom
• Bohr’s Model of Atom

The most trusted and scientifically established model of atom which is adopted these days is ‘Quantum Mechanical Model of Atom’. It will be dealt in higher classes.

Thomson’s Atomic Model

• This model is often called the ‘Water Melon Model’.
• In this model, Thomson predicted the presence of electrons inside positive sphere (made up of protons), just same as seeds of watermelon are embedded in red edible part of watermelon.
  
• Although this model explained neutrality of atom but couldn’t able to explain other scientific experiments conducted on atom. Hence it was discarded.

Rutherford’s Atomic Model

In his famous ‘α-ray Scattering Experiment’, Rutherford bombarded α-ray (Helium nucleus ₂He4) upon thin gold foil.

• Observations made by Rutherford in his experiment:
  • Most of α-particles passed through gold foil undeflected.
  • Some of the α-particles deflected by foil by small angles.
  • One out of every 12000 particles appeared to rebound.
    
• Conclusions made by Rutherford:
  • Atom consists of predominantly empty space as most of α-particles passed through gold foil undeflected.
  • Atom contains centrally placed positively charged nucleus (carrying positively charged particles), because few α-particles suffered deflected and very few i.e., one in 12000 bounced back.
  • Since a minute fraction of α-particles suffered deflections and very few bounced back, this lead to conclusion that most of the space an atom is empty and the space occupied by nucleus is negligible compared to this empty space.
    (i) Size of nucleus was about 10-5 times that of size of atom.
  • Whole of the atomic mass concentrated in the nucleus.
• Features of Rutherford proposed model of atom:
  • There is positively placed nucleus in an atom. Nearly all the mass resides in nucleus (Proton + Neutron).
  • Electrons revolves round the nucleus in well defined orbits.
  • Size of nucleus is very small compared to the size of atom.
• Drawbacks of Rutherford’s Model (Unstability of Atom)
  •  According to Rutherford, electrons revolve round the nucleus in well-defined orbits, but electrons being charged particles will lose their energy and finally will fall into the nucleus. 
  • This will make atom highly unstable.
  • This was the major drawback of Rutherford which was unexplained by him.
  • To overcome drawbacks of Rutherford’s Model, Neil Bohr in 1912 proposed modified model of structure of atom.

Assumption made by Neil Bohr

• Only certain special orbits known as discrete orbits of electrons are allowed inside the atom.
• While revolving in discrete orbits, the electrons do not radiate energy.
• Energy is emitted or absorbed by an atom only when an electron moves from one orbit to another.
  

Atomic Number

• The total number of proton lying in the nucleus of any atom is called the atomic number.
• An atomic number is the identity of an atom, changing atomic number means changing the atom.
• Atomic number is denoted by ‘Z’.
• Atomic number = no. of protons or a neutral atom, no. of protons and electrons are equal.

Mass Number

• It is the sum of total number of protons and no. of neutrons lying in the nucleus of an atom.
• It is denoted by ‘A’.
• Mass number = no. of protons + no. neutrons
• Representation of an atom: _ZE^A or ^AZ^E (E= symbol of an element)

Example: Calculate number of protons, electrons and neutrons for ₁₇Cl³⁵ or ³⁵₁₇Cl
Since Cl is neutral,
No. of electrons = no. of protons = 17
Mass no. of Cl = 35
No. of neutrons = 35 - 17 =18

Distribution of Electrons In Various Shells

The distribution of electrons in various shells is done in accordance to ‘Bohr-Bury Scheme’.

Bohr-Bury Scheme

• The filling of electrons in an atom is done in accordance to ‘2n2’, where ‘n’ is the number of shell and ‘2n2’ represents the total number of electrons that can be accommodated in that particular shell.
  • Maximum number of electrons that can be filled in particular shell.
    If n = 1, i.e., K = shell, 2n² = 2 × 1² = 2 electrons
    If n = 2, i.e., L = shell, 2n² = 2 × 2² = 8 electrons
    If n = 3, i.e., M = shell, 2n²2 = 2 × 3² = 18 electrons
    If n = 4, i.e., N = shell, 2n² = 2 × 4² = 32 electrons
• The outermost shell can’t hold more than 8 electrons, while second last shell can’t have more than 18 electrons, even though they may have capacity to hold more electrons.
  Example: ‘Ca₂₀’, the electron distribution will be :
  Ca_{20 =}  2(K), 8(L), 8(M), 2(N)
  • But Ca₂₀ = 2, 8, 10 is wrong although ‘M’ shell can contain upto 18 electrons.
• The outermost shell can’t hold more than 2 electrons and the penultimate shell can’t hold more than 8 electrons unless the preceding inner shell (antepenultimate shell) is filled completely obeying ‘2n²’ rule.
  Some examples:
  Ka₁₉ = 2, 8, 8, 1
  Al₁₃ = 2, 8, 3
  F₉ = 2, 7
  Ne₁₀ = 2, 8
   Na₁₁ = 2, 8, 1

Valence Shell and Valence Electrons

• From Bohr-Bury sequence, we know that maximum number of electrons which can be accommodated in outermost shell is 8.
• Every element has an urge to have 8 electrons in its outermost shell, in achieving 8 electrons an atom can either gain electrons or loose electrons.
• The number of electrons lost or gained by an element in achieving 8 electrons in its outermost shell will be called its Valence.
• For elements like H, He, Li, Be and B, these elements lose their outermost electron to achieve 2 electrons in their outermost shell. These elements will have valence in accordance to this act.
  

Isotopes

• Isotopes are atoms of same elements having same atomic number and different mass numbers.
  Example: Chlorine has two isotopes of mass numbers 35 and 37 respectively.
  ₁₇Cl³⁵, ₁₇Cl³⁷

Uses of isotopes

• Uranium isotope is used as fuel in nuclear rector.
• Isotope of cobalt is useful in treatment of cancer.
• An isotope of iodine is used in the treatment of goiter.

Isobars

• Isobars are the atoms of those elements which have the same mass number but different atomic numbers are called isobars.
• ₂₀Ca⁴⁰ and ₁₈Ar⁴⁰ have same mass number and different atomic number. ₁₁Na²⁴ and ₁₂Mg²⁴ are another examples.