Nuclei | Class 12 Physics Notes

13.2 ATOMIC MASSES AND COMPOSITION OF NUCLEUS

This section explains the measurement of atomic masses and the composition of the nucleus. Since atomic masses are extremely small compared to everyday units like kilograms, the atomic mass unit (u) is introduced. One atomic mass unit is defined as 1/12th the mass of a carbon-12 atom, which equals approximately 1.660539 × 10^-27 kg. Atomic masses of elements expressed in u are close to integral multiples of the hydrogen atom mass, but there are exceptions such as chlorine with an atomic mass of 35.46 u. Mass spectrometry is used to measure atomic masses accurately, revealing the existence of isotopes—atoms of the same element with identical chemical properties but different masses. For example, chlorine has two isotopes with masses 34.98 u and 36.98 u and relative abundances of 75.4% and 24.6%, respectively, resulting in an average atomic mass of 35.47 u. Hydrogen has three isotopes: protium (mass 1.0078 u), deuterium (2.0141 u), and tritium (3.0160 u). The proton is the nucleus of the lightest hydrogen isotope and carries a positive charge equal to the fundamental charge. The number of protons in a nucleus equals the atomic number Z, which also equals the number of electrons in a neutral atom. The section also introduces the discovery of the neutron, a neutral particle in the nucleus, which was confirmed by James Chadwick in 1932. Neutrons have a mass nearly equal to protons and are stable inside the nucleus but unstable when free. The composition of a nucleus is described by the atomic number Z (number of protons), neutron number N, and mass number A = Z + N. Nuclides are represented as A/Z X, where X is the chemical symbol. Isotopes have the same Z but different N; isobars have the same A but different Z; isotones have the same N but different Z.

🔗 Connection: Prepares for the next section on the size of the nucleus, building on the understanding of nuclear composition.