PERIODIC TRENDS

Question 1

Ionization Energy (IE):

Answer 1

Energy required to remove an electron.
Increases across a period (left to right) because of stronger nuclear attraction.
Decreases down a group because of increased shielding and atomic radius.

Question 2

Electronegativity

Answer 2

Ability of an atom to attract electrons in a bond.
Increases across a period due to increasing nuclear charge.
Decreases down a group because of greater atomic size and shielding.

Question 3

Atomic Radius:

Answer 3

Distance from the nucleus to the outermost electron.
Decreases across a period due to stronger nuclear pull.
Increases down a group due to the addition of electron shells.

Question 4

Shielding Effect:

Answer 4

Inner electron shells block outer electrons from nuclear attraction.
Increases down a group as more shells are added.
Remains constant across a period, as electrons are added to the same energy level

Question 5

Metallic Character:

Answer 5

Increases down a group (easier electron loss).

Decreases across a period (stronger nuclear attraction).

Question 6

Reactivity Trends

Metals

Answer 6

Metals: More reactive down a group (easier electron loss) and increases as you move to the left of the periodic table

Question 7

Reactivity Trends

Non-Metals

Answer 7

Non-Metals: More reactive up a group (stronger electron attraction).

Question 8

Effective Nuclear Charge

Answer 8

= protons – shielding (inner) electrons
Increases across a period (more protons, but similar shielding).
stays roughly the same down a group

Question 9

why do atoms get smaller across the periodic table

Answer 9

Atomic radius decreases across a period because increasing protons (+ charge) pull electrons closer. Since shielding remains nearly constant, effective nuclear charge (Z_eff) increases, drawing valence electrons inward and shrinking the atom.

Question 10

Why does effective nuclear charge increase from left to right?

Answer 10

Effective nuclear charge increases from left to right across the periodic table because of the increasing atomic charge and constant shielding effect. Due of a similar number of shielding electrons across a period, the valence electrons are pulled more tightly to the nucleus. This creates a smaller atomic radius, higher ionization energy, and higher net positive charge on the atom.