Atomic structure

An atom has a central nucleus of protons (relative charge $+1$, mass $1$) and neutrons (charge $0$, mass $1$), surrounded by electrons (charge $-1$, mass $\approx \frac{1}{1840}$) in shells. Almost all the mass is in the nucleus.

The atomic number $Z$ = number of protons; the mass number $A$ = protons + neutrons. In notation ${}_Z^A\text{X}$, a neutral atom has $Z$ electrons and $A-Z$ neutrons; for an ion, adjust the electrons by the charge.

Isotopes are atoms of the same element (same $Z$) with different numbers of neutrons (different $A$). Mass-spectrometer data gives their abundances.

Electrons fill shells from the inside out, holding at most $2,8,8$ for the first 20 elements. Write the configuration as comma-separated numbers, e.g. potassium (19) is 2,8,8,1.

The relative atomic mass $A_r$ is the average mass of an element's atoms, weighting each isotope by its abundance.

A mass spectrometer measures the masses and relative abundances of isotopes, giving a mass spectrum:

• x-axis: mass-to-charge ratio ($m/z$); since ions carry a $1+$ charge, $m/z$ equals the mass number.
• y-axis: relative abundance (height of the peak).

The tallest peak is the most abundant isotope. To calculate $A_r$, multiply each $m/z$ by its abundance, sum them, and divide by the total abundance:

$$A_r=\frac{\sum (m/z\times \text{relative abundance})}{\sum \text{relative abundance}}$$

• Forgetting to adjust electrons for an ion's charge.
• Confusing mass number (protons + neutrons) with $A_r$ (a weighted average).
• Overfilling a shell (the first three hold 2, 8, 8).