1. Atomic radius increases as you go down a group.
2. This is caused by the addition of electron shells (principal energy levels), which increases the distance between the nucleus and the outermost electrons.
3. Atomic radius decreases as you go across a period.
4. This is caused by increasing effective nuclear charge; protons are added to the nucleus while electrons are added to the same shell, pulling electrons closer.
5. a) Al b) S c) Br d) Na e) O f) Ca
6. Order: O < C < Sn < Sr. Oxygen and carbon are in period 2; O is farther right so smaller. Sn and Sr are in period 5; Sr is farther left so larger. Sn is below C, so larger than C. Sr is below Mg (in same group as Ca), and since Sr is farther down and left than Sn, it is largest.
7. Electronegativity is the ability of an atom to attract shared electrons in a chemical bond.
8. The ionic radius of a nonmetal is larger than its atomic radius because it gains electrons, increasing electron-electron repulsion and expanding the electron cloud.
9. Electronegativity decreases as you go down a group.
10. This is caused by increasing atomic size and greater shielding effect from inner electrons, reducing the nucleus’s pull on bonding electrons.
11. Electronegativity increases as you go across a period.
12. This is caused by increasing nuclear charge with no additional shielding (electrons added to same shell), resulting in stronger attraction for bonding electrons.
13. a) Ga b) O c) Cl d) Br e) Sr f) O
14. Group 1 (alkali metals)
15. Group 2 (alkaline earth metals)
16. Group 17 (halogens)
17. Group 18 (noble gases)
Parent Tip: Review the logic above to help your child master the concept of periodic table trends worksheet answers.
