1. Atomic radius increases as you go down a group/family.
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/row.
4. This is caused by increasing nuclear charge (more protons) pulling electrons closer to the nucleus, while electrons are added to the same shell.
5.
a) Al
b) S
c) Br
d) Na
e) O
f) Ca
6. C < O < Sn < Sr. Carbon and oxygen are in period 2; oxygen is smaller due to higher effective nuclear charge. Tin and strontium are in period 5; strontium is larger because it is further left in the period. Sn is larger than C or O because it has more electron shells. Sr is largest because it is in group 2 and has the same number of shells as Sn but less effective nuclear charge pulling electrons inward.
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 nonmetals gain electrons to form anions, increasing electron-electron repulsion and expanding the electron cloud.
9. Electronegativity decreases as you go down a group/family.
10. This is caused by increasing atomic size and greater shielding of the nucleus by inner electrons, reducing the atom’s ability to attract bonding electrons.
11. Electronegativity increases as you go across a period/row.
12. This is caused by increasing nuclear charge and decreasing atomic radius, allowing the nucleus to exert a stronger pull on bonding electrons.
13.
a) Ga
b) O
c) Cl
d) Br
e) Ba
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.
