* Mass Difference: Protons and neutrons have roughly the same mass (approximately 1 atomic mass unit, or amu). Electrons, on the other hand, are much lighter, with a mass of only about 1/1836 amu.
* Relative Abundance: Atoms have a similar number of protons and neutrons. The number of electrons is equal to the number of protons. This means the total mass of electrons is significantly less than the total mass of protons and neutrons.
Example:
Consider a carbon-12 atom:
* Protons: 6 (6 x 1 amu = 6 amu)
* Neutrons: 6 (6 x 1 amu = 6 amu)
* Electrons: 6 (6 x 1/1836 amu = 0.0033 amu)
The total mass of the protons and neutrons is 12 amu, while the electrons contribute only a tiny fraction of the total mass.
Practical Implications:
* Atomic Mass Units: The atomic mass unit (amu) is defined based on the mass of carbon-12, where the mass of the electrons is considered negligible.
* Calculations: In most chemical calculations, the mass of the electrons is insignificant and can be ignored for simplicity.
* Nuclear Reactions: In nuclear reactions where mass-energy conversions are significant, the electron mass is considered.
In summary: The mass of electrons is not entirely disregarded, but its contribution is so small compared to the nucleus that it is often ignored for practical purposes.
