Structure of the Atom - Comprehensive Notes

Charged Particles in Matter

Atoms and molecules are fundamental building blocks of matter.
Different kinds of matter are due to different atoms.
Key Questions:
- What differentiates atoms of different elements?
- Are atoms indivisible as Dalton proposed, or do they have smaller constituents?

Static Electricity

Evidence that atoms are not indivisible comes from studying static electricity.
Activity 4.1:
- Combing dry hair and observing attraction to paper.
- Rubbing a glass rod with silk and observing its interaction with a balloon.
- Conclusion: Rubbing objects can cause them to become electrically charged.

Discovery of Sub-Atomic Particles

By 1900, it was known that atoms contained at least one sub-atomic particle.
Electron: Discovered by J.J. Thomson.
Canal Rays:
- Discovered by E. Goldstein in 1886 during gas discharge experiments.
- Positively charged radiations.
- Led to the discovery of the proton.
Proton:
- Charge equal in magnitude but opposite in sign to the electron.
- Mass approximately 2000 times that of the electron.
Electron Representation: e-
Proton Representation: p+
Proton Mass: 1 unit
Proton Charge: +1
Electron Mass: Considered negligible
Electron Charge: -1
Initial Model: Atoms composed of protons and electrons, balancing charges.
Protons appeared to be in the interior of the atom, harder to remove than electrons.

Structure of the Atom

Dalton's Atomic Theory

Dalton’s theory suggested atoms were indivisible and indestructible.
Discovery of electrons and protons disproved this.

Thomson’s Model

Proposed by J.J. Thomson.
Model:
- Atom is a sphere of positive charge with electrons embedded in it, like currants in a Christmas pudding, or seeds in a watermelon.
Postulates:
- Atom consists of a positively charged sphere.
- Electrons are embedded in it.
- Negative and positive charges are equal in magnitude, making the atom electrically neutral.
Limitations:
- Could not explain experimental results obtained by other scientists.

Rutherford’s Model

Ernest Rutherford designed an experiment to understand electron arrangement.
Experiment: Alpha (α)-particles were directed at a thin gold foil.
- Gold foil was about 1000 atoms thick (thin as possible).
- α-particles are doubly-charged helium ions with a mass of 4 u.
- High energy due to fast movement.
- Expected α-particles to be deflected by sub-atomic particles, but not largely due to their mass.
Observations:
- Most α-particles passed straight through the gold foil.
- Some α-particles were deflected by small angles.
- About 1 in 12000 α-particles rebounded.
Conclusions:
- Most of the space inside the atom is empty.
- Positive charge occupies very little space.
- All positive charge and mass are concentrated in a small volume within the atom (nucleus).
- Radius of the nucleus is about 10^5 times less than the radius of the atom.
  *Model Features (Nuclear Model):
- Positively charged center called the nucleus.
- Nearly all the mass resides in the nucleus.
- Electrons revolve around the nucleus in circular paths.
- Nucleus size is very small compared to the atom size.
Drawbacks:
- Revolution of electrons in a circular orbit is unstable because any particle in a circular orbit would undergo acceleration.
- During acceleration, charged particles radiate energy.
- Revolving electrons would lose energy and fall into the nucleus.
- This would make atoms highly unstable, contradicting observed stability.

Bohr’s Model

Neils Bohr proposed postulates to address the issues with Rutherford's model.
Postulates:
- Electrons can only occupy certain special orbits (discrete orbits) within the atom.
- Electrons do not radiate energy while revolving in these discrete orbits.
These orbits or shells are called energy levels.
Energy levels are represented by letters K, L, M, N,… or numbers n = 1, 2, 3, 4,…. (See Fig. 4.3)

Neutrons

Discovered by J. Chadwick in 1932.
Sub-atomic particle with no charge and a mass nearly equal to that of a proton.
Present in the nucleus of all atoms except hydrogen.
Neutron Representation: ‘n’
Mass of an atom: Sum of the masses of protons and neutrons in the nucleus.

Distribution of Electrons in Orbits (Shells)

Bohr and Bury proposed rules for the distribution of electrons into different orbits.
Rules:
- The maximum number of electrons present in a shell is given by the formula 2n^2, where ‘n’ is the orbit number.
  - K-shell (n=1): 2 \times 1^2 = 2 electrons
  - L-shell (n=2): 2 \\times 2^2 = 8 electrons
  - M-shell (n=3): 2 \times 3^2 = 18 electrons
  - N-shell (n=4): 2 \times 4^2 = 32 electrons
- The maximum number of electrons that can be accommodated in the outermost orbit is 8.
- Electrons are not accommodated in a given shell unless the inner shells are filled in a stepwise manner.

Valency

Electrons in the outermost shell are called 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).
Inert elements, such as helium (2 electrons) or other elements (8 electrons), have a valency of zero.
Combining capacity (valency): Tendency of atoms 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. Finding Valency:
- The number of electrons gained, lost, or shared to make an octet gives the valency.
- Example: Hydrogen, lithium, and sodium have 1 electron in their outermost shell and can lose 1 electron; their valency is 1.
- If the number of electrons in the outermost shell is close to its full capacity
  - Fluorine: 7 electrons in the outermost shell. It is easier to gain 1 electron than lose 7, so its valency is 1 (8-7=1).

Atomic Number and Mass Number

Atomic Number

The number of protons in the nucleus determines the atomic number.
Denoted by ‘Z’.
All atoms of an element have the same atomic number.
Elements are defined by their number of protons.
- Hydrogen (H): Z = 1 (1 proton)
- Carbon (C): Z = 6 (6 protons)
Definition: The total number of protons present in the nucleus of an atom.

Mass Number

Mass of an atom is primarily due to protons and neutrons (nucleons).
Located in the nucleus.
Mass Number (A): Sum of the total number of protons and neutrons in the nucleus.
- Carbon: 12 u (6 protons + 6 neutrons)
- Aluminum: 27 u (13 protons + 14 neutrons)
Notation:
- ^A_Z \text{X}, where A is the mass number, Z is the atomic number, and X is the symbol of the element.
- Example: Nitrogen is written as ^{14}_7 \text{N}.

Isotopes

Atoms of some elements have the same atomic number but different mass numbers.
Example: Hydrogen has three isotopes:
- Protium (_1^1 \text{H}
- Deuterium (_1^2 \text{H} or D)
- Tritium (_1^3 \text{H} or T)
Atomic number is 1 for all, but mass numbers are 1, 2, and 3, respectively.
Other examples:
- Carbon: 6^{12} \text{C} and 6^{14} \text{C}
- Chlorine: {17}^{35} \text{Cl} and {17}^{37} \text{Cl}
Isotopes: Atoms of the same element with the same atomic number but different mass numbers.
Many elements consist of a mixture of isotopes, each being a pure substance.
Chemical properties of isotopes are similar, but physical properties are different.
Chlorine occurs in nature as two isotopes with masses 35 u and 37 u in a 3:1 ratio.
- The average atomic mass of chlorine is calculated as: \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 an atom of any natural element is taken as the average mass of all its naturally occurring isotopes.
If an element has no isotopes, its atomic mass is the sum of protons and neutrons.

Applications of Isotopes

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

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