Which statement best describes the relationship between atomic radius, ionization energy, and electronegativity across a period in the periodic table?

• A. As atomic radius decreases, ionization energy decreases and electronegativity increases.
• B. As atomic radius decreases, both ionization energy and electronegativity increase.
• C. As atomic radius increases, both ionization energy and electronegativity increase.
• D. As atomic radius increases, ionization energy increases and electronegativity decreases.

Answer: (B)

Verified step by step guidance

1. Step 1: Understand the periodic trends across a period (left to right) in the periodic table. As you move from left to right across a period, the number of protons in the nucleus increases, which affects atomic properties.
2. Step 2: Analyze how atomic radius changes across a period. With increasing nuclear charge and electrons added to the same energy level, the effective nuclear attraction pulls electrons closer, causing the atomic radius to decrease.
3. Step 3: Examine ionization energy across a period. Ionization energy is the energy required to remove an electron. As atomic radius decreases and nuclear charge increases, electrons are held more tightly, so ionization energy increases.
4. Step 4: Consider electronegativity across a period. Electronegativity is the tendency of an atom to attract electrons in a chemical bond. With a smaller atomic radius and higher nuclear charge, electronegativity increases across a period.
5. Step 5: Summarize the relationship: As atomic radius decreases across a period, both ionization energy and electronegativity increase. This matches the correct statement given.