AP Chemistry 1.7 Periodic Trends Study Notes - New Syllabus 2024-2025
AP Chemistry 1.7 Periodic Trends Study Notes- New syllabus
AP Chemistry 1.7 Periodic Trends Study Notes – AP Chemistry – per latest AP Chemistry Syllabus.
LEARNING OBJECTIVE
- Explain the relationship between trends in atomic properties of elements and electronic structure and periodicity.
Key Concepts:
- Periodic Trends
- Trends in Atomic Properties (Periodicity)
Periodic Trends
The periodic table is organized to reflect patterns of recurring properties among the elements. These patterns are determined primarily by the arrangement of electrons in atoms.
Ground-State Electron Configurations
- The ground-state electron configuration of an element describes the lowest-energy arrangement of electrons in shells and subshells.
- Elements with similar outer electron configurations appear in the same group, which explains similarities in chemical behavior.
Filled and Partially Filled Shells
- Completely filled shells (noble gas configuration) make atoms particularly stable and less chemically reactive.
- Partially filled shells or subshells influence reactivity, bonding patterns, and periodic properties such as atomic radius and ionization energy.
Periodic Properties Arising from Electron Configurations
- Elements in the same group have similar valence electron configurations → similar chemical properties.
- Atomic size, ionization energy, electronegativity, and electron affinity follow predictable trends across periods and down groups.
- These trends are a direct result of the number of electron shells, effective nuclear charge, and electron shielding.
Example
Explain why fluorine is more reactive than chlorine, even though both are in Group 17.
▶️ Answer/Explanation
• Both fluorine and chlorine are halogens (Group 17) and have 7 valence electrons.
• Fluorine is smaller in size than chlorine, so its nucleus has a stronger attraction for additional electrons (higher effective nuclear charge).
• This stronger attraction makes fluorine more reactive than chlorine, even though they have the same number of valence electrons.
• This illustrates how periodic trends, such as atomic size and electron shielding, influence chemical reactivity within a group.
Trends in Atomic Properties (Periodicity) with Exceptions
The periodic table allows us to predict periodic trends in atomic properties. These trends can be qualitatively explained using:
Coulomb’s law:
The force of attraction between electrons and the nucleus depends on nuclear charge and distance: \( F = k \dfrac{q_1 q_2}{r^2} \).
Shell model:
Electrons occupy shells; outer electrons (valence) are farther from the nucleus and less tightly held.
Shielding:
Inner electrons partially block the nuclear charge, reducing attraction on outer electrons.
Effective nuclear charge (\( Z_{eff} \)):
The net positive charge felt by a valence electron, accounting for shielding.
Main Periodic Properties
1. Ionization Energy (IE)
Energy required to remove 1 mole of electrons from 1 mole of gaseous atoms.
- Trend across a period: Increases → higher nuclear charge, same shielding, electrons held more tightly.
- Trend down a group: Decreases → more electron shells, outer electrons farther from nucleus, more shielding.
- Exceptions:
- Be → B: 2p electron in B is higher in energy and easier to remove → IE(B) < IE(Be).
- N → O: Half-filled 2p subshell in N is stable → O’s extra electron is easier to remove → IE(O) < IE(N).
2. Atomic and Ionic Radii
Atomic radius: Distance from nucleus to outermost electron.
- Trend across a period: Decreases → increasing nuclear charge pulls electrons closer.
- Trend down a group: Increases → more electron shells added, more distance from nucleus.
- Ionic radius: Cations are smaller (lose electrons, less electron-electron repulsion), anions are larger (gain electrons, more repulsion).
- Exceptions:
- Transition metals: Atomic radii change less regularly due to d-orbital electron shielding.
3. Electron Affinity (EA)
Energy change when 1 mole of electrons is added to 1 mole of gaseous atoms.
- Trend across a period: Generally becomes more negative → atoms more eager to gain electrons as nuclear charge increases.
- Trend down a group: Generally less negative → outer electrons farther from nucleus, weaker attraction.
- Exceptions:
- Be and N: Have filled or half-filled subshells → electron affinity is less negative.
- Group 2 elements: Low EA because adding an electron requires starting a new subshell.
4. Electronegativity
Ability of an atom in a molecule to attract shared electrons.
- Trend across a period: Increases → higher nuclear charge, smaller radius, stronger attraction for electrons.
- Trend down a group: Decreases → larger radius, more shielding, weaker pull on bonding electrons.
- Exceptions:
- Transition metals: Slight irregularities due to d-electron shielding.
Use of Periodicity
By recognizing trends, chemists can predict atomic and ionic properties even if experimental data are not available.
For example, knowing the trend of ionization energy in Group 1 allows prediction for elements like rubidium or francium.
Example
Predict and explain which element has a higher ionization energy: sodium (Na) or chlorine (Cl).
▶️ Answer/Explanation
• Sodium is in Group 1, Cl is in Group 17 (same period).
• Sodium has 1 valence electron far from the nucleus with more shielding → low ionization energy.
• Chlorine has 7 valence electrons, smaller radius, higher nuclear charge → electrons held more tightly.
• Therefore, Cl has a higher ionization energy than Na.
Example
Explain why fluorine is more electronegative than iodine.
▶️ Answer/Explanation
• Both are halogens (Group 17) → same number of valence electrons.
• Fluorine has fewer electron shells → smaller radius.
• Effective nuclear charge felt by valence electrons is stronger in fluorine.
• Therefore, fluorine attracts shared electrons in a bond more strongly → higher electronegativity than iodine.