Write and compare the ground-state electron configurations for each pair of elements. (a) Carbon and silicon (b) Oxygen and sulfur (c) Nitrogen and phosphorus

When we talk about the ground-state electron configuration of an element, we're referring to the arrangement of electrons in an atom's orbitals when the atom is in its lowest energy state. Each element has a unique electron configuration that can be determined by its atomic number, which is the number of protons (and electrons in a neutral atom) it has.

For instance, when comparing elements like carbon and silicon, the ground-state electron configuration reveals not only their placement in the periodic table but also their chemical properties. Both elements have a similar configuration in the first two shells, but silicon has additional electrons that 'fill up' the next shell. This pattern is consistent across the periodic table and is pivotal for understanding chemical reactions and bonds. Here are a few key points to remember:

• The electrons fill orbitals starting from the lowest energy level to the highest, following specific rules and principles.
• Electron configurations explain the reactivity and bonding of elements.
• The canonical form includes a series of numbers and letters that depict energy levels, subshells, and the number of electrons in those subshells, e.g., the configuration of carbon is denoted as 1s² 2s² 2p².

The Aufbau principle, which translates to 'building up' in German, is integral when determining the ground-state electron configurations of elements. It's a rule of thumb that prescribes the order in which subshells are filled as the atomic number increases.

An easy way to visualize the Aufbau principle is by using an energy-level diagram where subshells are represented as 'steps' that electrons 'climb'. The rule specifies that electrons occupy the lowest energy orbital available, and this is why we always fill the 1s orbital before moving to 2s, then 2p, and so on.

However, some complexities require attention, such as subshell crossover energies, where, for instance, the 4s subshell is filled before the 3d. To make this concept clearer for students, always start with the simple cases to build their understanding before introducing exceptions.

The atomic number of an element is perhaps the most fundamental attribute in chemistry. It is more than just a number - it represents the number of protons in the nucleus of an atom and by extension, assuming no electric charge, also equals the number of electrons.

Knowledge of the atomic number is the starting point for writing out the electron configuration. For example, nitrogen, with an atomic number of 7, will have seven electrons to place in the orbitals, following the rules we've outlined above. This information directly informs the chemical characteristics of the element and its position on the periodic table. Remember, the atomic number is unique to each element and is the primary identifier used by chemists.

In quantum chemistry, orbitals are like 'rooms' for electrons, and there are several types of orbitals - s, p, d, and f - each with a different shape and energy level. An orbital comparison is useful when examining how elements within the same group or period of the periodic table relate to each other.

Comparing orbitals between elements allows us to predict chemical behavior and reactivity. Elements with similar valence electron configurations, for example, tend to exhibit similar properties. Understanding orbital filling across elements (like nitrogen and phosphorus or oxygen and sulfur) is key in grasping periodic trends, such as ionization energy and electronegativity. Here, the idea of 'like fills like' will guide students as they learn to compare and contrast the electron configurations of different elements.