Nucle

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Atomic mass unit (u) is defined as 1/12th of the mass of a carbon-12 atom. 1u = 1.660539 × 10⁻²⁷ kg. This unit is used because atomic masses are extremely small compared to kilogram, making calculatio

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Isotopes are atomic species of the same element that have the same chemical properties but differ in mass. Example: Chlorine has two isotopes with masses 34.98 u (75.4% abundance) and 36.98 u (24.6% a

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Z = atomic number (number of protons) N = neutron number (number of neutrons) A = mass number (total nucleons) Relationship: A = Z + N Nuclear notation: ᴬZ X where X is the chemical symbol

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Mass of proton (mp) = 1.00727 u = 1.67262 × 10⁻²⁷ kg Mass of neutron (mn) = 1.00866 u = 1.6749 × 10⁻²⁷ kg Note: Neutron is slightly heavier than proton. Free neutron is unstable with mean life ~1000s.

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Nuclear radius formula: R = R₀ A^(1/3) where R₀ = 1.2 × 10⁻¹⁵ m = 1.2 fm This shows that nuclear volume is proportional to mass number A, making nuclear density constant for all nuclei (~2.3 × 10¹⁷ kg

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Einstein's mass-energy relation: E = mc² Significance: Mass can be converted to energy and vice versa. In nuclear reactions, mass defect is converted to binding energy. This principle explains the eno

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Mass defect (ΔM) = [Z·mp + (A-Z)·mn] - M where M is the actual nuclear mass. Binding energy: Eb = ΔM·c² This is the energy required to separate a nucleus into its constituent nucleons, or energy relea

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For ¹⁶O: Z = 8, N = 8, A = 16 Mass of constituents = 8(1.00727) + 8(1.00866) = 16.12744 u Nuclear mass = 15.99493 - 8(0.00055) = 15.99053 u Mass defect = 16.12744 - 15.99053 = 0.13691 u Binding energy