Structure of the Atom

Property	Electron	Proton	Neutron
Charge	-1 (negative)	+1 (positive)	0 (neutral)
Mass	\(9.1 \times 10^{-31}\) kg (≈ 1/1837 u)	\(1.67 \times 10^{-27}\) kg (≈ 1 u)	\(1.67 \times 10^{-27}\) kg (≈ 1 u)
Location	Outside nucleus (in orbits)	Inside nucleus	Inside nucleus
Discovery	J.J. Thomson (1897)	E. Goldstein (1886)	James Chadwick (1932)
Symbol	\(e^-\)	\(p^+\)	\(n\)

Limitations of Thomson's "Plum Pudding" Model:

1. Could not explain Rutherford's α-particle scattering experiment – Most α-particles passed straight through, suggesting atoms are mostly empty space, not a solid positive sphere.

2. Failed to explain the existence of a dense, positively charged nucleus – The model distributed positive charge uniformly, but Rutherford's experiment showed concentration of positive charge in a tiny nucleus.

3. Did not account for electron orbits – No explanation for how electrons are arranged or why they don’t fall into the positive sphere.

4. Could not explain atomic stability – According to classical physics, electrons in motion should radiate energy and spiral into the nucleus, but atoms are stable.

Limitations of Rutherford's Nuclear Model:

1. Could not explain atomic stability – According to Maxwell's electromagnetic theory, an electron revolving around the nucleus should continuously lose energy and spiral into the nucleus, but atoms are stable.

2. Did not explain the arrangement of electrons – The model didn't specify how electrons are distributed around the nucleus.

3. Could not account for discrete atomic spectra – If electrons could orbit at any radius, they would emit a continuous spectrum, but atoms emit only specific line spectra.

4. No explanation for energy levels – The model lacked the concept of quantized energy levels for electrons.

Bohr's Model of the Atom (1913):

1. Electrons revolve in specific orbits called "stationary orbits" or "energy levels" without radiating energy.

2. Each orbit has a fixed energy – Energy levels are quantized (discrete).

3. Orbits are labeled as K, L, M, N,… or n = 1, 2, 3, 4,…

4. Electrons can jump between orbits by absorbing or emitting energy in the form of photons.

5. Maximum electrons in a shell is given by \(2n^2\), where n is the orbit number.

6. Outermost shell can hold maximum 8 electrons (octet rule for stability).

Significance: Explained atomic stability, discrete spectra, and introduced quantized energy levels.

Model	Key Features	Limitations
Thomson's Model (Plum Pudding)	• Atom = positive sphere with embedded electrons • Electrically neutral	• No nucleus • Could not explain α-scattering
Rutherford's Model (Nuclear Model)	• Dense, positive nucleus • Electrons orbit nucleus • Mostly empty space	• Electrons should spiral into nucleus • No energy levels
Bohr's Model	• Fixed energy levels (orbits) • Electrons don’t radiate in orbits • Explains atomic spectra	• Only works for hydrogen-like atoms • Doesn't explain fine spectra

Rules for Electron Distribution (Bohr-Bury Scheme):

1. Maximum electrons in a shell = \(2n^2\)
  K shell (n=1): 2 electrons
  L shell (n=2): 8 electrons
  M shell (n=3): 18 electrons (but only 8 filled first in first 18 elements)

2. Outermost shell cannot have more than 8 electrons (octet rule).

3. Shells are filled stepwise – Inner shells must be filled before moving to the next shell.

4. First 18 elements follow:
  • K shell fills first (max 2 e⁻)
  • L shell fills next (max 8 e⁻)
  • M shell starts filling after L is full (first 8 e⁻ in M)

Example: Sodium (Na, Z=11): K=2, L=8, M=1

Valency is the combining capacity of an element, determined by the number of electrons it can lose, gain, or share to achieve a stable octet (8 electrons in outermost shell).

Example 1: Silicon (Si, atomic number 14)
Electronic configuration: K=2, L=8, M=4
Valency = 4 (can share 4 electrons to complete octet)

Example 2: Oxygen (O, atomic number 8)
Electronic configuration: K=2, L=6
Valency = 2 (can gain 2 electrons to complete octet, or share 2 electrons)

(i) Atomic Number (Z): Number of protons in the nucleus of an atom.
Example: Carbon has Z = 6 (6 protons).

(ii) Mass Number (A): Sum of protons and neutrons in the nucleus.
Example: Carbon-12 has A = 12 (6 protons + 6 neutrons).

(iii) Isotopes: Atoms of the same element with same Z but different A (different neutrons).
Example: Hydrogen isotopes: Protium (\(_1^1H\)), Deuterium (\(_1^2H\)), Tritium (\(_1^3H\)).

(iv) Isobars: Atoms of different elements with same A but different Z.
Example: \(_{20}^{40}Ca\) and \(_{18}^{40}Ar\) both have mass number 40.

Two uses of isotopes:
1. Medical: Cobalt-60 for cancer treatment.
2. Archaeological: Carbon-14 for radiocarbon dating.

Sodium atom (Na): Atomic number = 11
Electronic configuration: K=2, L=8, M=1

Sodium ion (Na⁺): Formed by losing 1 electron from M shell.
New configuration: K=2, L=8 (M shell becomes empty)

Result:
• K shell is filled with 2 electrons (max capacity = 2)
• L shell is filled with 8 electrons (max capacity = 8)
• Thus, Na⁺ has completely filled K and L shells, making it stable like noble gas Neon (Ne).

Given:
Isotope 1: \(_{35}^{79}Br\), abundance = 49.7% = 0.497
Isotope 2: \(_{35}^{81}Br\), abundance = 50.3% = 0.503

Formula:
Average atomic mass = (Mass₁ × Abundance₁) + (Mass₂ × Abundance₂)

Calculation:
= (79 × 0.497) + (81 × 0.503)
= (39.263) + (40.743)
= 80.006 ≈ 80 u

Result: Average atomic mass of bromine ≈ 80 u

Let:
Abundance of \(_8^{16}X\) = x%
Abundance of \(_8^{18}X\) = (100 - x)%

Equation:
Average mass = \(\frac{(16 \times x) + (18 \times (100-x))}{100} = 16.2\)

Solve:
16x + 1800 - 18x = 1620
-2x = 1620 - 1800
-2x = -180
x = 90%

Result:
\(_8^{16}X\) = 90%
\(_8^{18}X\) = 10%

Given: Atomic number Z = 3

Element: Lithium (Li)

Electronic configuration: K=2, L=1

Valency: 1 (can lose 1 electron to achieve stable configuration)

For X:
Mass number = Protons + Neutrons = 6 + 6 = 12
So X is \(_6^{12}C\) (Carbon-12)

For Y:
Mass number = 6 + 8 = 14
So Y is \(_6^{14}C\) (Carbon-14)

Relation: X and Y are isotopes of carbon (same Z, different A).

(a) F – Thomson did not propose nucleus; Rutherford discovered nucleus.

(b) F – Neutron is a fundamental particle, not formed by electron+proton combination.

(c) T – Electron mass ≈ 1/1837 of proton mass (~1/2000 approx).

(d) T – Iodine-131 is used in medicine for thyroid treatment.

Correct option: (a) Atomic Nucleus

Rutherford's α-scattering experiment (1911) led to the discovery of the dense, positively charged nucleus at the center of the atom.

Correct option: (c) different number of neutrons

Isotopes have same atomic number (same protons) but different mass numbers due to different neutrons.

Correct option: (b) 8

Chlorine atom (Cl, Z=17): K=2, L=8, M=7 (valence electrons = 7)
Chloride ion (Cl⁻): Gains 1 electron → M shell becomes 8
Valence electrons in Cl⁻ = 8

Correct option: (d) 2,8,1

Sodium (Na, Z=11):
K shell (n=1): 2 electrons
L shell (n=2): 8 electrons
M shell (n=3): 1 electron
Configuration: 2,8,1

Atomic Number	Mass Number	Neutrons	Protons	Electrons	Name
9	19	10	9	9	Fluorine
16	32	16	16	16	Sulphur
12	24	12	12	12	Magnesium
1	2	1	1	1	Deuterium
1	1	0	1	0	Proton/H⁺ ion