4.3.3 Ionization Energy

This topic explains ionization energy — the energy required to remove the outermost electron from an atom.

Definition

Ionization energy is defined as the minimum amount of energy required to remove the outermost electron from an isolated gaseous atom.

Simple explanation: Ionization energy is the energy needed to pull an electron away from an atom.

Main Idea

Ionization energy is an important property of atoms that explains cation formation.

Equation:

M(g) + ionization energy → M⁺(g) + e^-

Simple explanation: When an atom loses an electron, it becomes a positive ion (cation). The energy needed for this process is the ionization energy.

Core Concept

What Ionization Energy Tells Us

Ionization energy is a measure of the extent to which the nucleus attracts the outermost electron.

• High ionization energy → stronger attraction between nucleus and outermost electron.
• Low ionization energy → weaker force of attraction between nucleus and outermost electron.

Simple explanation: If the nucleus holds the electron tightly, it takes more energy to remove it (high ionization energy). If the hold is weak, less energy is needed (low ionization energy).

Trends in Ionization Energy

Detailed Explanation

Trend 1: Ionization Energy Decreases Down a Group

The value of the ionization energy decreases from top to bottom in the group.

• The shielding effect of the atoms increases down the group.
• Greater shielding effects result in a weaker attraction of the valence electrons to the nucleus.
• So they are easier to remove.
• This leads to a decrease in ionization energy from top to bottom in the group.

• Na has a greater shielding effect than Li.
• Therefore, Na has lower ionization energy than Li.

Simple explanation: Down a group, the outer electron is farther from the nucleus and shielded by more inner electrons, so it's easier to remove.

Trend 2: Ionization Energy Increases Across a Period

• The shielding effect remains unchanged.
• But little by little the nuclear charge increases.
• The stronger attraction between the nucleus and the valence electron increases.
• As a result, the ionization energy increases from left to right in a period.

• Be has higher ionization energy than Li.
• Because Be has one more proton than Li, so the nucleus holds the electrons more tightly.

Simple explanation: Across a period, the nucleus gets stronger (more protons), but the shielding stays the same, so the electrons are held more tightly.

Book Example: Identifying the Element with Smaller Ionization Energy

Question: Choose the element whose atom you expect to have smaller ionization energy in each of the following pairs. (a) B, C (b) N, P

Problem Solving Strategy: Remember that ionization energy: (a) Increases across a period. The element that has smaller ionization energy will be further to the left in the periodic table. (b) Decreases down a group. The element that has smaller ionization energy will be further down in the periodic table.

Solution: (a) B has smaller ionization energy (B is to the left of C). (b) P has smaller ionization energy (P is below N).

Values from Figure 4.5 (page 58)

Some ionization energy values from the book (in kJ/mol):

• Across Period 2: Li (520) < Be (899) < B (801) < C (1086) < N (1402) > O (1314) < F (1681) — general increase, with small irregularities.
• Down Group 1: Li (520) > Na (496) > K (419) — decreases down the group.

Intermediate Level

Why Ionization Energy Decreases Down a Group

• New shells are added → valence electrons are farther from the nucleus.
• Shielding effect increases → the nucleus's pull on valence electrons is reduced.
• Effective nuclear charge increases only slightly.
• So the valence electron is easier to remove → lower ionization energy.

Why Ionization Energy Increases Across a Period

• Protons increase → nuclear charge increases.
• Shielding stays constant (inner shells don't change).
• Effective nuclear charge increases → the nucleus pulls valence electrons more tightly.
• So it becomes harder to remove an electron → higher ionization energy.

Factors Affecting Ionization Energy

Successive Ionization Energies

• First ionization energy (IE₁)
• Second ionization energy (IE₂)
• Third ionization energy (IE₃)

• The remaining electrons are held more tightly by the increasingly positive ion.
• Removing an electron from a positive ion requires more energy.

Relation to Cation Formation

• Group 1 metals (Li, Na, K) have the lowest ionization energies → form +1 ions easily.
• Group 2 metals (Mg, Ca) have slightly higher ionization energies → form +2 ions.
• Noble gases have the highest ionization energies → do not form ions.

Important Points

• - Ionization energy: minimum energy required to remove the outermost electron from an isolated gaseous atom.
• - Equation: M(g) + energy → M⁺(g) + e^-
• - Down a group: ionization energy decreases (shielding increases, valence electrons farther from nucleus).
• - Across a period: ionization energy increases (nuclear charge increases, shielding constant).
• - High ionization energy → strong attraction between nucleus and outermost electron.
• - Low ionization energy → weak attraction.
• - Elements with low ionization energy form cations easily.
• - Successive ionization energies are always larger than previous ones.

Remember

• - Ionization energy = energy to remove outermost electron.
• - Down a group → decreases.
• - Across a period → increases.
• - High IE = strong attraction; low IE = weak attraction.
• - Group 1 metals → lowest IE → form +1 ions easily.
• - Noble gases → highest IE → no ions.
• - Successive IEs are always larger.

Common Mistakes

• 1. Thinking ionization energy increases down a group. It decreases.
• 2. Thinking ionization energy decreases across a period. It increases.
• 3. Confusing ionization energy with electronegativity. Ionization energy is about removing an electron; electronegativity is about attracting an electron in a bond.
• 4. Forgetting that ionization energy is for a gaseous atom. The definition specifies "isolated gaseous atom."
• 5. Thinking the second ionization energy is smaller than the first. It is always larger.
• 6. Confusing shielding effect with ionization energy. Greater shielding → lower ionization energy.

Quick Revision

• Ionization energy: minimum amount of energy required to remove the outermost electron from an isolated gaseous atom.
• Equation: M(g) + ionization energy → M⁺(g) + e^-
• Down a group: ionization energy decreases (shielding increases, valence electrons farther from nucleus).
• Across a period: ionization energy increases (nuclear charge increases, shielding constant).
• High ionization energy = strong attraction between nucleus and outermost electron.
• Low ionization energy = weak attraction.
• Elements with low ionization energy form cations easily.
• Successive ionization energies are always larger than the previous ones.

60-Second Revision

• Ionization energy = energy to remove outermost electron.
• Down a group → decreases.
• Across a period → increases.
• Group 1 metals → low IE → form +1 ions.
• Noble gases → high IE → no ions.
• Warning: ionization energy increases across a period, not decreases.