An atom’s electrons are arranged in distinct energy levels, also known as electron shells, designated by the principal quantum number, \(n\). The structure of the Periodic Table reflects this organization. The row, or period, an element occupies corresponds precisely to the highest energy level containing at least one electron. Therefore, elements in the fourth row (Period 4) have electrons populating the fourth energy level (\(n=4\)).
Identifying the Elements with Four Energy Levels
The elements with four energy levels belong to Period 4 of the Periodic Table. This row is composed of 18 elements, beginning with Potassium (K, atomic number 19) and concluding with the noble gas Krypton (Kr, atomic number 36). The elements in this period are highly diverse, encompassing three different blocks of the table. The period starts with Potassium and Calcium (the two \(s\)-block elements). Following these are ten transition metals, from Scandium (Sc) through Zinc (Zn), which constitute the \(d\)-block, also known as the first transition series. The row concludes with six \(p\)-block elements, starting with Gallium (Ga) and ending with Krypton.
How Electrons Fill the Fourth Shell
Electron filling in the fourth energy level follows the Aufbau principle, which dictates that electrons occupy the lowest available energy orbitals first. Although the fourth principal energy level (\(n=4\)) includes \(4s\), \(4p\), \(4d\), and \(4f\) subshells, the \(4s\) subshell has lower energy than the \(3d\) subshell from the third level. Consequently, the first two electrons in the fourth period enter the \(4s\) orbital, as seen in Potassium (K, \([Ar]4s^1\)) and Calcium (Ca, \([Ar]4s^2\)).
Once the \(4s\) orbital is full, the next electrons begin to fill the \(3d\) orbitals. The \(3d\) subshell can hold ten electrons, which accounts for the ten transition metals from Scandium to Zinc. For example, Iron (Fe) has the configuration \([Ar]3d^64s^2\), showing that its valence electrons are distributed between the \(3d\) and \(4s\) orbitals.
This filling pattern is not perfectly linear due to the added stability of half-filled and completely filled subshells. Chromium (Cr) and Copper (Cu) are two exceptions in this period. Chromium promotes one electron from the \(4s\) to the \(3d\) orbital, resulting in a stable \(3d^54s^1\) configuration. Similarly, Copper achieves a fully filled \(3d\) shell with a \(3d^{10}4s^1\) configuration. After the \(3d\) orbitals are filled at Zinc, the remaining six electrons fill the \(4p\) subshell, completing the period with Krypton’s inert gas configuration.
Common Uses and Properties of These Elements
The elements in the fourth period exhibit a wide array of properties, ranging from highly reactive metals to an inert gas. The \(d\)-block transition metals are notable for their distinctive characteristics, which arise from their partially filled \(d\) orbitals. These properties include the ability to form compounds with multiple oxidation states and the tendency to produce colored salts and solutions.
Many of these metals are crucial for industrial applications and biological processes. Iron (Fe) is the most prominent, serving as a structural material in steel and forming the active site in the oxygen-carrying protein hemoglobin. Copper (Cu) is valued for its exceptional electrical conductivity, making it the standard for wiring. Zinc (Zn) is used to galvanize steel, protecting it from corrosion, and is an important trace element for human health.
Other elements in this period are also highly utilized. Titanium (Ti) is prized in aerospace and medical implants for its high strength-to-weight ratio and corrosion resistance. Chromium (Cr) is a component of stainless steel and is used in electroplating for a durable, shiny finish. Even the non-metals and metalloids in the \(p\)-block, such as Germanium (Ge), are indispensable as semiconductors in the electronics industry.