MCQs — Atomic Structure

Because atoms are extremely light, their masses are expressed using the atomic mass unit rather than grams.

One atomic mass unit is defined as exactly 1/12 the mass of a carbon-12 atom.

Relative atomic mass compares an element's average atomic mass to the carbon-12 standard.

Since relative atomic mass compares one mass to another, the units cancel out, leaving a pure ratio.

Averaging different isotopic masses, weighted by their abundance, usually produces a decimal value.

Carbon-12 is the reference standard and is defined as having a mass of exactly 12 amu.

Carbon-12 is a stable, abundant isotope that serves as a reliable reference for atomic mass measurements.

Relative atomic mass directly reflects how many times heavier the atom is than 1/12 the mass of carbon-12.

The relative atomic mass equals the ratio of the atom's mass to 1/12 the mass of carbon-12, which is 16 here.

Relative atomic mass is a comparative average mass value, not a whole number in every case.

Using a relative scale based on carbon-12 avoids working with impractically tiny numbers in grams.

If X is twice as massive as Y, its relative atomic mass is roughly double that of Y: 2 x 7 = 14.

Periodic table values represent a weighted average across all naturally occurring isotopes of an element.

Atomic mass unit is a unit of measurement, while mass number is simply a count of nucleons, with no unit.

Both protons and neutrons have a mass close to 1 atomic mass unit each.

Isotopes share the same number of protons but differ in the number of neutrons, giving different mass numbers.

Since atomic number (protons) stays the same, isotopes differ only in neutron number.

Chemical behaviour depends mainly on electron arrangement, which is the same across isotopes of an element.

Since mass depends on neutron count, isotopes with more neutrons are typically denser or heavier.

Isotopic notation displays the mass number and atomic number to specify a particular isotope.