Atoms of the same chemical element all have the same number of protons, but they can contain different numbers of neutrons. Such variants are isotopes.

An isotope is commonly written

$$^{A}_{Z}X,$$

where $X$ is the element symbol, $Z$ is the proton number, and

$$A=Z+N$$

is the mass number, the total number of protons and neutrons. For example, carbon-12 and carbon-13 both have $Z=6$, but contain six and seven neutrons respectively.

Atomic masses are conveniently measured in unified atomic mass units, symbol $\mathrm u$, defined so that a neutral carbon-12 atom has mass exactly $12,\mathrm u$.

A naturally occurring element may contain several isotopes. The atomic mass listed for such an element is therefore generally a weighted average rather than an integer. If isotope $i$ has fractional abundance $f_i$ and atomic mass $m_i$, then

$$\bar m=\sum_i f_i m_i,$$

with $\sum_i f_i=1$.

For chlorine, using approximate natural abundances and isotopic masses,

$$\bar m\approx(0.7578)(34.969,\mathrm u)+(0.2422)(36.966,\mathrm u) \approx35.45,\mathrm u.$$

This explains why periodic-table atomic masses such as chlorine's are usually not whole numbers. The average describes a macroscopic sample with the stated isotopic composition; an individual chlorine atom is one isotope, not an atom of mass $35.45,\mathrm u$.