Resources

Periodic Trends

Learning Objectives

Studying periodic trends in AP Chemistry involves understanding and predicting how atomic radius, ionization energy, electron affinity, and electronegativity vary across periods and down groups in the periodic table. This includes explaining the underlying reasons for these trends based on atomic structure and nuclear charge. Additionally, it involves the ability to compare these properties among different elements, make predictions about their chemical behavior, and apply this knowledge to solve problems related to element reactivity, bonding, and periodicity.

Periodic trends are patterns observed in the periodic table that help explain the chemical and physical behavior of elements. These trends include variations in atomic radius, ionization energy, electron affinity, and electronegativity as you move across periods and down groups. Understanding these trends is essential for predicting how elements will react, bond, and interact with each other. In AP Chemistry, mastering these concepts provides a foundation for understanding more complex chemical principles and enhances problem-solving skills related to element properties and behaviors.

Periodic trends are the predictable patterns that occur within the periodic table as a result of the arrangement of elements in order of increasing atomic number. These trends include changes in atomic radius, ionization energy, electron affinity, and electronegativity, and they help explain the chemical behavior and properties of the elements. Understanding these trends allows for the prediction of element reactivity and bonding characteristics.

Atomic Radius

Definition: The atomic radius is the distance from the center of an atom’s nucleus to the outermost electron shell. It is typically measured in picometers (pm) or angstroms (Å).

Factors Affecting Atomic Radius

  1. Nuclear Charge: The number of protons in the nucleus. A higher nuclear charge pulls electrons closer, reducing the atomic radius.
  2. Electron Shielding: Inner electrons shield outer electrons from the full effect of the nuclear charge, causing a larger atomic radius.
  3. Electron Shells: More electron shells mean a larger atomic radius, as outer electrons are farther from the nucleus.

Trends in Atomic Radius

Across a Period (Left to Right):

  • Decreases.
  • As you move across a period, the number of protons in the nucleus increases. This increased nuclear charge attracts the electron cloud closer to the nucleus, decreasing the atomic radius.

Example:

  • Lithium (Li) vs. Fluorine (F):
    • Lithium: Atomic radius ≈ 152 pm
    • Fluorine: Atomic radius ≈ 60 pm
    • Fluorine has a smaller atomic radius than lithium due to its higher nuclear charge pulling electrons closer.

Down a Group (Top to Bottom):

  • Increases.
  • As you move down a group, each successive element has an additional electron shell. This increases the distance between the nucleus and the outermost electrons, resulting in a larger atomic radius.

Example:

  • Fluorine (F) vs. Chlorine (Cl):
    • Fluorine: Atomic radius ≈ 60 pm
    • Chlorine: Atomic radius ≈ 99 pm
    • Chlorine has a larger atomic radius than fluorine because it has an additional electron shell.

Key Points

  • Smallest atomic radius: Top-right corner of the periodic table (e.g., Helium).
  • Largest atomic radius: Bottom-left corner of the periodic table (e.g., Francium).

Key Examples

  1. Hydrogen (H):
    • Atomic radius: ≈ 53 pm
    • Small due to having only one electron and no inner shells.
  2. Sodium (Na):
    • Atomic radius: ≈ 186 pm
    • Larger than lithium due to the addition of an electron shell.

Ionization Energy

Definition: Ionization energy is the amount of energy required to remove an electron from a gaseous atom or ion. It is typically measured in kilojoules per mole (kJ/mol).

Factors Affecting Ionization Energy

  1. Nuclear Charge: A higher nuclear charge increases the attraction between the nucleus and the electrons, increasing the ionization energy.
  2. Atomic Radius: A larger atomic radius means the outer electrons are farther from the nucleus and less tightly held, decreasing the ionization energy.
  3. Electron Shielding: Inner electrons shield the outer electrons from the full effect of the nuclear charge, reducing the ionization energy.

Trends in Ionization Energy

Across a Period (Left to Right):

  • Increases.
  • As you move across a period, the nuclear charge increases, making it more difficult to remove an electron because the electrons are held more tightly by the nucleus.

Example:

  • Boron (B) vs. Nitrogen (N):
    • Boron: First ionization energy ≈ 800 kJ/mol
    • Nitrogen: First ionization energy ≈ 1402 kJ/mol
    • Nitrogen has a higher ionization energy than boron due to its higher nuclear charge and smaller atomic radius.

Down a Group (Top to Bottom):

  • Decreases.
  • As you move down a group, the atomic radius increases due to the addition of electron shells. The outer electrons are farther from the nucleus and more shielded by inner electrons, making them easier to remove.

Example:

  • Nitrogen (N) vs. Phosphorus (P):
    • Nitrogen: First ionization energy ≈ 1402 kJ/mol
    • Phosphorus: First ionization energy ≈ 1012 kJ/mol
    • Phosphorus has a lower ionization energy than nitrogen because its outer electrons are farther from the nucleus and more shielded by inner electron shells.

Key Points

  • Highest ionization energy: Top-right corner of the periodic table (e.g., Helium).
  • Lowest ionization energy: Bottom-left corner of the periodic table (e.g., Francium).

Key Examples

  1. Helium (He):
    • First ionization energy: ≈ 2372 kJ/mol
    • Highest in the periodic table due to its small size and high nuclear charge relative to its electron configuration.
  2. Magnesium (Mg):
    • First ionization energy: ≈ 737 kJ/mol
    • Higher than sodium due to a higher nuclear charge and smaller atomic radius.

Electron Affinity

Definition: Electron affinity is the amount of energy released when an electron is added to a neutral atom in the gaseous state to form a negative ion. It is typically measured in kilojoules per mole (kJ/mol).

Factors Affecting Electron Affinity

  1. Nuclear Charge: A higher nuclear charge attracts additional electrons more strongly, increasing electron affinity.
  2. Atomic Radius: A smaller atomic radius means the added electron is closer to the nucleus, increasing electron affinity.
  3. Electron Shielding: Less shielding by inner electrons means the nucleus can attract additional electrons more strongly, increasing electron affinity.

Trends in Electron Affinity

Across a Period (Left to Right):

  • Generally increases (becomes more negative).
  • As you move across a period, the nuclear charge increases, making the atom more likely to gain an electron and release energy in the process.

Example:

  • Carbon (C) vs. Oxygen (O):
    • Carbon: Electron affinity ≈ -122 kJ/mol
    • Oxygen: Electron affinity ≈ -141 kJ/mol
    • Oxygen has a more negative electron affinity than carbon due to its higher nuclear charge and smaller atomic radius.

Down a Group (Top to Bottom):

  • Generally decreases (becomes less negative).
  • As you move down a group, the atomic radius increases due to the addition of electron shells. The added electron is farther from the nucleus and experiences more shielding, decreasing the electron affinity.

Example:

  • Chlorine (Cl) vs. Iodine (I):
    • Chlorine: Electron affinity ≈ -349 kJ/mol
    • Iodine: Electron affinity ≈ -295 kJ/mol
    • Chlorine has a more negative electron affinity than iodine because its smaller atomic radius and less shielding effect result in a stronger attraction for the added electron.

Key Points

  • Highest electron affinity: Halogens (e.g., Chlorine).
  • Lowest electron affinity: Noble gases and some alkali metals.

Key Examples

  1. Fluorine (F):
    • Electron affinity: ≈ -328 kJ/mol
    • High electron affinity due to high nuclear charge and small atomic radius, but not the highest due to repulsion in its small electron cloud.
  2. Chlorine (Cl):
    • Electron affinity: ≈ -349 kJ/mol
    • Highest electron affinity in the periodic table, as it readily gains an electron to achieve a stable noble gas configuration.

Electronegativity

Definition: Electronegativity is a measure of the tendency of an atom to attract a bonding pair of electrons. It is a dimensionless quantity and is often measured on the Pauling scale, where fluorine has the highest value of 4.0.

Factors Affecting Electronegativity

  1. Nuclear Charge: A higher nuclear charge increases electronegativity because the nucleus more strongly attracts bonding electrons.
  2. Atomic Radius: A smaller atomic radius means the bonding electrons are closer to the nucleus, increasing electronegativity.
  3. Electron Shielding: Less shielding by inner electrons allows the nucleus to more effectively attract bonding electrons, increasing electronegativity.

Trends in Electronegativity

Across a Period (Left to Right):

  • Increases.
  • As you move across a period, the nuclear charge increases while the atomic radius decreases, resulting in a stronger attraction for bonding electrons.

Example:

  • Lithium (Li) vs. Fluorine (F):
    • Lithium: Electronegativity ≈ 1.0
    • Fluorine: Electronegativity ≈ 4.0
    • Fluorine has a much higher electronegativity than lithium due to its higher nuclear charge and smaller atomic radius.

Down a Group (Top to Bottom):

  • Decreases.
  • As you move down a group, the atomic radius increases due to the addition of electron shells, and the increased shielding effect reduces the nucleus’s ability to attract bonding electrons.

Example:

  • Fluorine (F) vs. Chlorine (Cl):
    • Fluorine: Electronegativity ≈ 4.0
    • Chlorine: Electronegativity ≈ 3.0
    • Chlorine has a lower electronegativity than fluorine because its larger atomic radius and greater electron shielding reduce the attraction for bonding electrons.

Key Points

  • Highest electronegativity: Fluorine (top-right corner of the periodic table).
  • Lowest electronegativity: Cesium and Francium (bottom-left corner of the periodic table).

Key Examples

  1. Fluorine (F):
    • Electronegativity: 4.0
    • Highest electronegativity in the periodic table, due to its small atomic radius and high nuclear charge.
  2. Sulfur (S):
    • Electronegativity: 2.5
    • Lower than oxygen but still relatively high within its period, making it more likely to attract bonding electrons than many other elements.