13.2. ATOMIC MASSES AND COMPOSITION OF NUCLEUS

The mass of an atom is very small, compared to a kilogram; for example, the mass of a carbon atom, 12C, is 1.992647 × 10–26 kg. Kilogram is not a very convenient unit to measure such small quantities. Therefore, a different mass unit is used for expressing atomic masses. This unit is the atomic mass unit (u), defined as 1/12th of the mass of the carbon (12C) atom. According to this definition

mass of one 12C atom

1u =

12

1.992647 · 10^-26 kg

=

12

=1.660539· 10^-27 kg (13.1)

Atoms are incredibly tiny, and measuring their mass in kilograms does not make much sense because those numbers would be extremely small. Instead, scientists use a unit called the atomic mass unit (u). This unit acts as a more practical way to express atomic masses. By defining 1 u as 1/12th the mass of a carbon-12 (12C) atom, we can approximate the sizes of other atoms in a more manageable scale.

Think of measuring very small things in grams rather than kilograms. For example, just like it's easier to say a chocolate bar weighs 100 grams instead of 0.1 kilograms, it's easier to use atomic mass units when discussing the weights of atoms.

The atomic masses of various elements expressed in atomic mass unit (u) are close to being integral multiples of the mass of a hydrogen atom. There are, however, many striking exceptions to this rule. For example, the atomic mass of chlorine atom is 35.46 u. Accurate measurement of atomic masses is carried out with a mass spectrometer. The measurement of atomic masses reveals the existence of different types of atoms of the same element, which exhibit the same chemical properties but differ in mass. Such atomic species of the same element differing in mass are called isotopes.

Most atomic masses reflect a simple relationship to hydrogen's mass, often presenting as whole numbers. However, some elements, like chlorine, have atomic masses that aren't whole numbers, suggesting that atomic masses are averages. These averages account for isotopes, which are variants of the same element that have different numbers of neutrons, affecting their mass but not their chemical behavior.

Consider a family with siblings: they have the same parents (like isotopes having the same number of protons), but can vary in height or weight (like variations in neutron count). They share many characteristics but still have their individual differences.

It was found that practically every element consists of a mixture of several isotopes. The relative abundance of different isotopes differs from element to element. Chlorine, for example, has two isotopes having masses 34.98 u and 36.98 u, which are nearly integral multiples of the mass of a hydrogen atom. The relative abundances of these isotopes are 75.4 and 24.6 per cent, respectively. Thus, the average mass of a chlorine atom is obtained by the weighted average of the masses of the two isotopes.

Most elements exist as mixtures of several isotopes, with each having a set percentage of abundance. For chlorine, it shows two primary isotopes with known masses. To find its average atomic mass, we perform calculations based on their relative abundances—figuring out how much each isotope contributes to the total mass.

Just as a smoothie might be made from different fruits that vary in sweetness, the average sweetness is a blend of each fruit's contribution, weighted by how much of each fruit is used in the recipe.

Since the nuclei of deuterium and tritium are isotopes of hydrogen, they must contain only one proton each. But the masses of the nuclei of hydrogen, deuterium, and tritium are in the ratio of 1:2:3. Therefore, the nuclei of deuterium and tritium must contain, in addition to a proton, some neutral matter. This fact indicates that the nuclei of atoms contain, in addition to protons, neutral matter in multiples of a basic unit. This hypothesis was verified in 1932 by James Chadwick who observed emission of neutral radiation when beryllium nuclei were bombarded with alpha-particles.

The study of isotopes led to the conclusion that for certain isotopes like deuterium and tritium, their mass ratios indicated additional, neutral particles must exist alongside protons in the nucleus. Chadwick’s experiments in the 1930s confirmed the presence of neutrons, which explained why atomic mass wasn't always a simple sum of protons.

Imagine a team where some members are visible (protons) and some are invisible but still contribute to the team's performance (neutrons). Just as the invisible members have a real effect on the team, so do neutrons in the atomic nucleus.