Unit 6: Periodic Trends
Shielding (also called the "shielding effect"): Inner-shell electrons block (shield) the outer electrons from feeling the full attractive force of the nucleus.
Effective Nuclear Charge (Zeff): The NET positive charge that valence electrons actually experience. It accounts for shielding.
Zeff = Actual nuclear charge (Z) − Shielding from inner electrons
Example:
• Sodium (Na): Z = 11 protons, but has 10 inner electrons (1s²2s²2p⁶) shielding the one 3s valence electron.
Approximate Zeff on the valence electron ≈ 11 − 10 = +1
• Chlorine (Cl): Z = 17 protons, with 10 inner shielding electrons.
Approximate Zeff ≈ 17 − 10 = +7
Conclusion: As you move ACROSS a period (left to right), Zeff INCREASES because protons are added but inner shielding electrons stay the same — the valence electrons feel a stronger pull.
Atomic radius: the size of an atom, measured approximately as half the distance between two bonded nuclei.
Trend across a PERIOD (left → right): Atomic radius DECREASES
Why: Increasing Zeff pulls electrons closer to the nucleus. More protons, same number of shells, so the atom contracts.
Trend down a GROUP (top → bottom): Atomic radius INCREASES
Why: Each new period adds a new electron shell that is farther from the nucleus. More shielding and more distance result in a larger atom.
Memory aid:
• Largest atoms: bottom-left corner of periodic table (Cs, Ba)
• Smallest atoms: top-right corner (except noble gases) — like F and Ne
Ionization energy (IE): the energy required to REMOVE one valence electron from a neutral gas-phase atom. Higher IE = harder to remove an electron = electron is more tightly held.
Trend across a PERIOD (left → right): Ionization energy INCREASES
Why: Greater Zeff means valence electrons are held more tightly — it takes more energy to remove them.
Trend down a GROUP (top → bottom): Ionization energy DECREASES
Why: Valence electrons are farther from the nucleus and more shielded. Less energy needed to remove them.
Memory aid: IE trends are the OPPOSITE of atomic radius trends.
• Large atom (bottom-left) = low IE (easy to remove electrons)
• Small atom (top-right) = high IE (hard to remove electrons)
Electronegativity (EN): the tendency of an atom to attract electrons TOWARD itself in a chemical bond. Higher EN = greater ability to attract bonding electrons.
Trend across a PERIOD (left → right): EN INCREASES
Why: Greater Zeff means the nucleus pulls shared electrons more strongly.
Trend down a GROUP (top → bottom): EN DECREASES
Why: Valence electrons are farther from the nucleus and more shielded — less attraction for bonding electrons.
Note: Fluorine (F) has the HIGHEST electronegativity of all elements. Francium (Fr) has the LOWEST.
You will NOT be given the electronegativity chart or asked to calculate EN differences. You need to know the general trends to make comparisons between elements.
The potential energy of valence electrons relates to their distance from the nucleus and the strength of the nuclear attraction (Zeff).
Higher potential energy = electrons farther from nucleus OR weaker nuclear attraction
Lower potential energy = electrons closer to nucleus OR stronger nuclear attraction
Trend across a period: Zeff increases → valence electrons drawn closer → LOWER potential energy
Trend down a group: Electrons are in higher shells, farther from nucleus → HIGHER potential energy
This connects to ionization energy: higher potential energy electrons are easier to remove (lower IE), while lower potential energy electrons require more energy to remove (higher IE).