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Among the most revolutionary concepts introduced by Dalton's atomic theory was the idea of atomic mass. Dalton proposed that each element possessed a characteristic atomic mass, distinct from other elements. This groundbreaking concept was particularly powerful because it provided a robust explanation for the Law of Constant Proportions, inspiring scientists to devise methods for measuring these elusive atomic masses.

However, directly measuring the mass of an individual atom was, and still is, an incredibly difficult task due to their minuscule size. To overcome this, scientists initially focused on determining relative atomic masses. They achieved this by using the laws of chemical combination and observing the compounds formed between elements.

Early Attempts at an Atomic Mass Unit:

Initially, scientists took 1/16th of the mass of an atom of naturally occurring oxygen as their standard unit for atomic mass. This choice was favored for two main reasons:

1. Oxygen readily reacted with a large number of elements, forming compounds, which made it a convenient reference point for comparisons.
2. Using oxygen as a standard often resulted in the atomic masses of most other elements being close to whole numbers, which was convenient for calculations.

The Carbon-12 Standard:

Despite the convenience of oxygen, in 1961, a universally accepted atomic mass unit was finally adopted. This standard chooses carbon-12 isotope as the reference.

The modern definition states that one atomic mass unit (amu), now officially written as 'u' (for unified mass) according to IUPAC recommendations, is a mass unit exactly equal to one-twelfth (1/12th) the mass of one atom of carbon-12.

Think of it like a fruit seller (Fig. 3.4) who doesn't have standard weights. They might take a watermelon, declare it to have a mass of "12 units," then cut it into 12 equal pieces. They can then weigh other fruits relative to these watermelon pieces. Similarly, all relative atomic masses of elements are now determined with respect to an atom of carbon-12. The relative atomic mass of the atom of an element is therefore defined as the average mass of the atom, as compared to 1/12th the mass of one carbon-12 atom.

This standardization provided a consistent and precise way to compare the masses of different atoms, further solidifying the quantitative nature of chemistry. It allows us to assign specific atomic masses to elements, which are crucial for understanding chemical reactions and the composition of compounds.

Example Atomic Masses (u):

• Hydrogen: 1 u
• Carbon: 12 u
• Nitrogen: 14 u
• Oxygen: 16 u
• Sodium: 23 u
• Magnesium: 24 u
• Sulphur: 32 u
• Chlorine: 35.5 u
• Calcium: 40 u

Ultimately, while atoms are too small to be seen individually by the naked eye, the concept of atomic mass allows chemists to work with them indirectly, making accurate predictions and calculations about how matter behaves.

Question for You: Why do you think scientists chose a specific isotope (carbon-12) as the standard for atomic mass rather than a naturally occurring average of carbon atoms?