chem review

1. Front: What is effective nuclear charge?

Back: The net positive charge felt by an electron, after accounting for shielding by inner electrons.

2. Front: What is the shielding effect?

Back: Inner electrons block outer electrons from the full pull of the nucleus, reducing the effective nuclear charge.

3. Front: Define ionization energy (IE).

Back: Energy needed to remove an electron from a gaseous atom. Higher IE = harder to remove an electron.

4. Front: Define atomic radius.

Back: The size of an atom, measured as half the distance between two bonded nuclei.

5. Front: What is ionic radius?

Back: The size of an ion. Cations (+) are smaller than their parent atom; anions (−) are larger.

6. Front: Define electronegativity (EN).

Back: An atom’s ability to attract electrons in a bond. Highest in fluorine (EN = 4.0).

Periodic Trends Explained (No Z_eff)

7. Front: How does ionization energy (IE)change across a period (left to right)?

Back: Increases because:

• More protons = stronger nuclear pull.

• Electrons are added to the same shell (no extra shielding).

8. Front: How does IE change down a group?

Back: Decreases because:

• More electron shells = more shielding.

• Outer electrons are farther from the nucleus.

9. Front: How does atomic radius change across a period?

Back: Decreases because:

• More protons pull electrons closer.

• No new shells are added.

10. Front: How does atomic radius change down a group?

Back: Increases because:

• More electron shells = larger atom.

11. Front: Why are cations smaller than their parent atoms?

Back: Lost electrons = less repulsion + same nuclear pull "shrinks" the ion.

12. Front: Why are anions larger than their parent atoms?

Back: Gained electrons = more repulsion + same nuclear pull "expands" the ion.

13. Front: How does electronegativity (EN)change across a period?

Back: Increases because:

• Stronger nuclear pull attracts electrons better.

14. Front: How does EN change down a group?

Back: Decreases because:

• More shielding + farther electrons = weaker attraction.

Bonus Application Questions

15. Front: Which has a larger atomic radius: Na or K? Why?

Back: K (potassium)—it has an extra electron shell compared to Na.

16. Front: Which has higher IE: O or S? Why?

Back: O (oxygen)—fewer electron shells = less shielding = harder to remove an electron.

17. Front: Why is fluorine the most electronegative element?

Back: Small size + strong nuclear pull (many protons, few shells).