Atomic radius

Definition

Atomic radius is a measure of the size of an atom, typically defined as the distance from the nucleus to the outermost electron shell. This concept is essential for understanding how atoms interact with each other and how their size affects chemical properties, such as reactivity and bonding. The atomic radius varies across the periodic table, influenced by factors like atomic number and electron shielding, which in turn affect an element's physical and chemical behavior.

5 Must Know Facts For Your Next Test

1. Atomic radius generally increases down a group in the periodic table due to the addition of electron shells, which outweighs the increase in nuclear charge.
2. Atomic radius tends to decrease across a period from left to right because of increased effective nuclear charge, pulling electrons closer to the nucleus.
3. Metals typically have larger atomic radii compared to nonmetals within the same period, as metals tend to lose electrons more easily.
4. Cations have a smaller atomic radius than their neutral atoms because losing electrons reduces electron-electron repulsion and allows the remaining electrons to be pulled closer to the nucleus.
5. Anions have a larger atomic radius than their neutral atoms due to the addition of extra electrons, increasing electron-electron repulsion and pushing them further apart.

Review Questions

• How does atomic radius change across a period and down a group in the periodic table, and what are the reasons for these trends?
  • Atomic radius decreases across a period due to an increase in effective nuclear charge, which pulls electrons closer to the nucleus. In contrast, atomic radius increases down a group because additional electron shells are added, making the atoms larger despite an increasing nuclear charge. The interplay of these two trends reflects how atomic structure influences size.
• Discuss how atomic radius impacts ionization energy and electronegativity for elements in the periodic table.
  • Atomic radius significantly affects both ionization energy and electronegativity. As atomic radius decreases across a period, ionization energy increases because it requires more energy to remove an electron that is closer to the nucleus. Similarly, smaller atoms with shorter distances between their nucleus and electrons exhibit higher electronegativity since they can attract bonding electrons more effectively than larger atoms.
• Evaluate how changes in atomic radius can influence chemical reactivity among different elements in groups and periods.
  • Changes in atomic radius play a crucial role in determining chemical reactivity. For example, larger atoms with greater atomic radii tend to be more reactive as they have outer electrons that are farther from the nucleus and experience less attraction, making it easier for them to lose or share electrons during reactions. Conversely, smaller atoms have stronger attractions to their valence electrons, which can make them less reactive. Thus, understanding these size differences provides insights into an element's likelihood of participating in chemical reactions.

Related terms

ionization energy: The energy required to remove an electron from an atom or ion, which is influenced by the atomic radius and the effective nuclear charge.

electronegativity: The tendency of an atom to attract electrons in a chemical bond, which is affected by atomic radius among other factors.

electron shielding: The phenomenon where inner-shell electrons partially block the attractive force of the nucleus on outer-shell electrons, impacting the atomic radius.