Periodic Table Trends: Atomic Radius

• Periodic table trends are specific patterns that are present in the periodic table that illustrate different aspects of a certain element, including its size and electronic properties.
• These trends include atomic radius, ionization energy, electronegativity, ionic radii, and electron configuration.

Atomic Radius

• Atomic radius generally decreases from left to right across a period.
• Within a period, electrons are added to the same energy level as protons which are added to the nucleus.
• The increased positive charge in the nucleus attracts the negatively charged electrons more strongly, pulling them closer to the nucleus, decreasing the atomic radius.
• Atomic radius generally increases moving down a group.
• As you move down a group, electrons are added to higher energy levels, farther from the nucleus.
• Core electrons shield the outer electrons from the full positive charge of the nucleus, reducing the effective nuclear charge experienced by the outer electrons.
• Increased distance from the nucleus and the shielding effect cause the atomic radius to increase.

Ionization Energy

• Ionization energy is the energy required to remove an electron from a neutral atom in its gaseous phase.
• Ionization energy generally increases when moving from left to right across a period.
• As you move across a period, the effective nuclear charge increases, making it more difficult to remove an electron.
• Smaller atoms hold their electrons more tightly, requiring more energy to remove them.
• Ionization energy generally decreases moving down a group.
• As you move down a group, the outermost electrons are farther from the nucleus, shielded by core electrons.
• This reduces the effective nuclear charge experienced by the outer electrons, making them easier to remove, decreasing the ionization energy.

Electronegativity

• Electronegativity measures the ability of an atom to attract electrons in a chemical bond.
• Electronegativity generally increases when moving from left to right across a period.
• As you move across a period, the effective nuclear charge increases, making the atom more attractive to electrons.
• Electronegativity generally decreases when moving down a group.
• As you move down a group, outermost electrons are farther from the nucleus and shielded by core electrons.
• This reduces the effective nuclear charge experienced by the outer electrons, making the atom less attractive to electrons, decreasing the electronegativity.
• Fluorine is the most electronegative element.

Ionic Radii

• Ionic radii refers to the radius of an atom's ion.
• Cations (positive ions) form when an atom loses electrons.
• Cations are smaller than their neutral atoms because they have fewer electrons but the same number of protons.
• The remaining electrons are more strongly attracted to the nucleus, causing the ion to shrink.
• Anions (negative ions) form when an atom gains electrons.
• Anions are larger than their neutral atoms because they have more electrons but the same number of protons.
• Increased electron-electron repulsion causes the electron cloud to expand, increasing the size of the ion.
• For isoelectronic species (ions with the same number of electrons), the ion with the greater nuclear charge will have a smaller radius.

Electron Configuration

• Electron configuration describes the arrangement of electrons within an atom.
• Filling of electron orbitals follows specific rules and principles:
  • Aufbau Principle: Electrons fill the lowest energy levels first.
  • Hund's Rule: Within a subshell, electrons individually occupy each orbital before any orbital is doubly occupied; all electrons in singly occupied orbitals have the same spin.
  • Pauli Exclusion Principle: No two electrons in an atom can have the same set of four quantum numbers; therefore, each orbital can hold a maximum of two electrons with opposite spins.
• Electron configuration is closely related to position on the periodic table; predictable based on the group and period of the element.
• Elements in the same group typically have similar outer electron configurations, leading to similar chemical properties.
• Knowing the electron configuration helps predict trends in atomic size, ionization energy, and chemical reactivity.
• Valence electrons (outermost electrons) are responsible for the chemical properties of an element.
• Atoms gain, lose, or share electrons to achieve a stable electron configuration, typically a full outer shell (octet rule).