Within an atom, electrons are influenced by two primary forces: the attractive force from the positively charged nucleus and the repulsive force from other negatively charged electrons. This duality is crucial for understanding atomic structure and behavior. For instance, an electron experiences attraction towards the nucleus due to the positive charge of protons, while simultaneously facing repulsion from other electrons in the outer shell, which are also negatively charged.

The concept of effective nuclear charge, denoted as \( Z_{\text{eff}} \), quantifies the net attractive force experienced by an electron from the nucleus after accounting for the repulsion from other electrons. A higher effective nuclear charge indicates a stronger attraction between the nucleus and the electron, leading to the electron being drawn closer to the nucleus. This relationship can be summarized as follows: as \( Z_{\text{eff}} \) increases, the attractive force between the nucleus and the electron also increases.

In addition to effective nuclear charge, the shielding constant plays a significant role in electron interactions. The shielding constant measures the extent to which inner core electrons repel outer valence electrons. As the shielding constant increases, the repulsive force among valence electrons also increases, pushing them further away from the nucleus. This dynamic illustrates how the interplay of attraction and repulsion shapes the arrangement and behavior of electrons within an atom.

In summary, the balance between the attractive force from the nucleus and the repulsive forces among electrons is fundamental to understanding atomic structure, influencing properties such as atomic size and ionization energy.