The size of atoms is a fundamental property that dictates how elements interact with one another. Atomic radius is defined as one-half the distance between the nuclei of two identical atoms that are chemically bonded together. A predictable pattern emerges when moving across the periodic table from left to right: the atomic radius generally decreases. This means that an atom like Lithium (Li) is larger than an atom like Neon (Ne). The contraction of the atom’s size across a row, or period, is a consequence of competing forces that result in a greater inward pull on the electrons.
The Role of Electron Shells
As you move from one element to the next across a single horizontal row on the periodic table, electrons are progressively added to the atom. Crucially, these electrons are being placed into the same principal energy level, also known as the valence shell. For example, elements in the second period, from Lithium to Neon, add their outermost electrons to the second energy level. This constancy is significant because adding a new principal energy level, which occurs when moving down a column, causes a large increase in atomic size. The fact that the outermost electron shell remains the same sets the stage for the forces that cause the decrease in atomic size.
Understanding Effective Nuclear Charge
The primary reason for the reduction in atomic size is the effective nuclear charge (\(Z_{eff}\)). As you move across a period, the number of protons in the nucleus increases by one for each element. This sequential addition of positive charge creates a stronger attractive force originating from the nucleus. \(Z_{eff}\) is the net positive charge experienced by an atom’s outermost valence electrons, which is less than the total number of protons due to electron shielding. Because the number of protons steadily increases, the effective nuclear charge experienced by the valence electrons also increases, pulling the electron cloud closer to the nucleus.
The Screening Effect
The counteracting force to the nuclear attraction is the screening effect, also known as electron shielding. This effect occurs because the inner, non-valence electrons repel the outer valence electrons, partially blocking the full positive charge of the nucleus from reaching them. Across a period, the number of inner-shell electrons remains constant because new electrons are added to the same outermost shell. The existing inner electrons provide the same amount of shielding for each successive element. Since the newly added valence electrons do not effectively screen each other, the overall shielding effect remains relatively unchanged across the period.
The Combined Effect: Why Atoms Shrink
The trend of decreasing atomic radius results from the interplay between the increasing effective nuclear charge and the constant screening effect. As you move across a period, the nucleus gains an additional proton, increasing the strength of the inward pull. Simultaneously, the shielding provided by the core electrons remains the same. Because the increase in nuclear attraction far outweighs the unchanging repulsive effects, the net force on the outermost electron cloud is a stronger attraction toward the center.
This stronger net inward pull draws the valence electrons closer to the nucleus, physically reducing the size of the atom. For instance, Lithium (Li) has an atomic radius of about 152 picometers, while Neon (Ne) has a radius of approximately 38 picometers, demonstrating significant contraction. The electrons remain in the same principal energy level, but that energy level is compressed closer to the nucleus due to the stronger electrostatic attraction.