What is the trend for ionization energy across a period in the periodic table?

Step 1: Analyze the question You're asked how first ionization energy (the energy needed to remove the outermost electron from a gaseous atom) changes as you move across a period, from left to right.

Step 2: Explain the concept Across a period:

• The nuclear charge (number of protons) increases by one for each successive element.
• Electrons are added to the same shell, so shielding by inner electrons stays about the same.
• The effective nuclear charge therefore increases, and the atomic radius decreases.
• The outer electrons are held more tightly, so more energy is needed to remove one.

Step 3: Worked example (Period 2) Li → Be → B → C → N → O → F → Ne

First ionization energies (approx., kJ/mol): Li 520, Be 900, B 800, C 1086, N 1402, O 1314, F 1681, Ne 2081.

The overall trend is a clear increase from Li to Ne. There are small dips (Be → B and N → O). These happen because of subshell stability (a filled 2s subshell in Be, and a half-filled 2p subshell in N, with electron-pair repulsion in O). They do not change the general trend.

Step 4: Evaluate the options

• A. Increases from left to right. Correct.
• B. Remains constant. Wrong, because nuclear charge keeps changing.
• C. Varies randomly. Wrong, because there is a clear overall pattern.
• D. Decreases from left to right. Wrong, because this is the trend for atomic radius, not ionization energy.

Step 5: Exam tip Remember: across a period, ionization energy, electronegativity and electron affinity generally increase, while atomic radius decreases. Down a group, ionization energy decreases because the atoms get larger and shielding increases. The noble gases have the highest ionization energy in each period.