13.2- ch 8 periodic trends

Periodic Trends

• Periodic trends refer to observable patterns in elemental properties as you move across or down the periodic table.

• There are four periodic trends, with three essential ones to study for the exam: atomic radius, ionization energy, and electron affinity.

Atomic Radius

• Definition: The atomic radius is the distance from the nucleus of an atom to the outer edge of its electron cloud or the last orbital.

• Measurement:

  • Measured using experimental methods like x-ray diffraction in crystalline materials.

  • In diatomic molecules, distances between nuclei can vary slightly due to asymmetry, but approximations can be made.

Effective Nuclear Charge

• Concept: Inner shell electrons shield valence electrons from the nucleus, impacting atomic size.

  • Valence electrons do not shield each other.

  • As atomic number increases (going left to right), protons and core electrons increase, enhancing effective nuclear charge experienced by valence electrons.

• Trend Explanation:

  • Adding protons without additional shielding leads to a stronger attraction for valence electrons.

  • The atomic size decreases as you move from left to right on the periodic table due to increasing effective nuclear charge.

  • The relationship between atomic radius and effective nuclear charge explains why atoms become smaller across a period.

Trends in Atomic Radius

• Left to Right: Atomic radius generally decreases across a period due to increasing effective nuclear charge.

• Top to Bottom: Atomic radius increases down a group due to increased shell numbers (higher valence orbitals) and inner electron shielding effects.

  • Electrons in higher shells are farther from the nucleus, contributing to a larger radius.

Ionic Radii

• Definition: Ionic radius refers to the size of an ion, which can differ from its neutral atom size due to loss or gain of electrons.

• Cations:

  • Formed by the loss of electrons; for instance, Group 1A elements lose one electron to form +1 cations.

  • Ionic radii decrease as you go up the group (e.g., Li+ smaller than Na+).

• Anions:

  • Formed by gaining electrons; for example, halogens which gain electrons tend to increase in size.

  • As anions (e.g., F-, Cl-, Br-, I-), they are larger than their respective neutral atoms due to increased electron-electron repulsion and decreased effective nuclear charge.

Ionization Energy

• Definition: Ionization energy is the energy required to remove an electron from an atom. The first ionization energy measures energy to remove the first electron.

• Trend:

  • Metals (left side of the periodic table) have lower ionization energies compared to nonmetals (right side).

  • Noble gases exhibit the highest ionization energies due to their stable electron configurations.

• Trends in Ionization Energy:

  • Ionization energy increases as you move right across a period and decreases as you move down a group due to increased electron shielding and distance from the nucleus.

Electron Affinity

• Definition: The change in energy when an atom gains an electron.

• Trend: Generally, nonmetals have higher electron affinities as they tend to gain electrons, while metals have lower values.

  • Electron affinity values are typically expressed as positive numbers, reflecting the energy change associated with gaining an electron, despite representing a favorable energy decrease.

Sample Problems

1. Largest Ionic Radius: Compare Sn2+ vs. Sn4+. Sn2+ is larger because it retains more electrons, experiencing less nuclear attraction.

2. Highest and Lowest Ionization Energy: In elements like arsenic, bromine, and cobalt, bromine has the highest ionization energy as it’s a nonmetal; cobalt (metal) has the lowest.

3. Atomic Radius Comparison: Among iron, strontium, and carbon, carbon has the smallest atomic radius due to proximity to the top right of the periodic table, while strontium has the largest atomic radius.