1 Atomic Structure and Electron Configurations

Atomic Identity and Particle Counts

An atom has a tiny, dense nucleus surrounded by electrons. The nucleus contains positively charged protons and, in most atoms, neutral neutrons. Electrons occupy regions of space called orbitals. An orbital describes where an electron is likely to be found; it is not a fixed path around the nucleus.

The number of protons is the atomic number, represented by $Z$, and determines the element’s identity. In a neutral atom, the number of electrons equals the number of protons. The mass number, represented by $A$, counts the protons and neutrons together:

For example, carbon-14 has atomic number $6$ and mass number $14$. It therefore has $6$ protons and $14-6=8$ neutrons. Because the atom is neutral, it also has $6$ electrons. Its isotope notation is $^{14}_{6}\mathrm{C}$.

Takeaway: Proton count identifies the element; the total of protons and neutrons gives the mass number.

Isotopes and Ions

Isotopes are atoms of the same element with different numbers of neutrons. Carbon-12 and carbon-14 are both carbon because each has $6$ protons. Carbon-12 has $6$ neutrons, while carbon-14 has $8$. The mass number applies to a particular isotope, whereas the decimal atomic mass shown on many periodic tables is a weighted average of naturally occurring isotopes.

An ion forms when an atom or group of atoms gains or loses electrons, producing a net electric charge. Its charge is calculated as:

Losing electrons produces a positively charged cation; gaining electrons produces a negatively charged anion. For example, sodium has $11$ protons. The ion $\mathrm{Na}^{+}$ has lost one electron and has $10$ electrons. The oxide ion $\mathrm{O}^{2-}$ has $8$ protons and $10$ electrons. Ordinary ion formation changes the electron count, not the element’s identity: changing the proton count would change the element.

Takeaway: Isotopes differ in neutron count; ions differ in electron count and carry a charge.

Electron Configurations and Filling Rules

An electron configuration records how an atom’s electrons occupy energy levels and subshells. The subshell labels are $s$, $p$, $d$, and $f$, with maximum capacities of $2$, $6$, $10$, and $14$ electrons, respectively. Each orbital can hold up to two electrons.

For a ground-state atom, electrons generally fill lower-energy orbitals first, following the Aufbau principle. Within a set of equal-energy orbitals, Hund’s rule says that electrons occupy separate orbitals before pairing. The Pauli exclusion principle requires two electrons in the same orbital to have opposite spins.

A common filling sequence is $1s, 2s, 2p, 3s, 3p, 4s, 3d, 4p$. Oxygen has $8$ electrons, giving the configuration $1s^2\,2s^2\,2p^4$. Sodium has $11$ electrons, giving $1s^2\,2s^2\,2p^6\,3s^1$. Shorthand notation uses a preceding noble gas to represent the configuration of the inner electrons, so sodium can also be written $[\mathrm{Ne}]\,3s^1$.

Some elements do not follow the simplest filling prediction because subshell energies are close. Chromium’s ground-state configuration is $[\mathrm{Ar}]\,3d^5\,4s^1$, rather than the simplest predicted $[\mathrm{Ar}]\,3d^4\,4s^2$. For transition-metal cations, electrons are generally removed from the outermost principal energy level first. Neutral iron is $[\mathrm{Ar}]\,3d^6\,4s^2$, while $\mathrm{Fe}^{2+}$ is $[\mathrm{Ar}]\,3d^6$.

Takeaway: Filling rules provide a useful method for writing configurations, but some elements are exceptions; when forming transition-metal cations, account for which electrons are removed.

Periodic Table Patterns

The periodic table arranges elements in order of increasing atomic number. Its horizontal rows are periods, and its vertical columns are groups or families. Elements in the same group often behave similarly because they have similar arrangements of outer electrons, called valence electrons.

The table’s blocks identify the subshell being filled:

• The $s$-block occupies the left two columns, with helium’s configuration as an exception to its position at the far right.

• The $p$-block occupies the right six columns.

• The $d$-block is the central transition-metal region.

• The $f$-block consists of the lanthanides and actinides, commonly displayed in two rows beneath the main table.

For main-group elements, group position is a useful guide to valence-electron count. Group one elements have one valence electron, group two elements have two, and groups thirteen through eighteen generally have three through eight. Helium is an exception, with two. As electron configurations repeat across periods, similar outer-electron arrangements recur in groups, helping explain recurring chemical similarities.

Takeaway: Periods, groups, and blocks connect an element’s place in the table to its electron configuration and, in turn, to patterns in chemical behavior.

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