S1.3.7 Successive ionization energies (Higher Level Only)

Using Ionization Energy Data to Determine Group Numbers and Analyze Trends

Successive Ionization Energy and Group Number

What Is Ionization Energy?

Definition

Ionization energy

Ionization energy ( $I E$ ) is the minimum energy required to remove an electron from a gaseous atom in its ground state.

Example

The first ionization energy ( $I E_{1}$ ) refers to the energy needed to remove the first electron.
The second ionization energy ( $I E_{2}$ ) is the energy required to remove the second electron, and so on.

These processes can be represented as follows:

For the first ionization:
$X(g) + energy \to X^{+} (g) + e^{-}$

For the second ionization:
$X^{+} (g) + energy \to X^{2 +} (g) + e^{-}$

Hint

Successive ionization energies always increase because, as electrons are removed, the remaining electrons experience a stronger electrostatic attraction to the positively charged nucleus.

Note

Ionization energy is measured for gaseous atoms to ensure that interactions between atoms don’t affect the energy values.

Graph of successive ionization energies of sodium: a visual representation of energy required to remove electrons from 3s, 2p, 2s, and 1s subshells.

Using Ionization Energy Data to Deduce Group Number

The key to identifying an element’s group lies in identifying large jumps in successive ionization energy values.
These jumps occur when an electron is removed from a stable, filled energy level (or noble gas configuration).
Let’s break this process down:

Identify the Large Jump: Examine the ionization energy data for significant increases. For example:
- $I E_{1} = 500 kJ/mol$
- $I E_{2} = 4560 kJ/mol$
- $I E_{3} = 6910 kJ/mol$
A massive jump between $I E_{1}$ and $I E_{2}$ suggests that the first electron is being removed from a new energy level, while the second electron is removed from a stable noble gas configuration.
- This indicates the element has one valence electron, placing it in Group 1.
Relate to the Periodic Table:
- For main group elements (Groups 1, 2, and 13–18), the number of valence electrons corresponds to the group number.
- For transition metals, interpreting ionization energy trends requires understanding $d$ -sublevel electron configurations.

Example

Successive ionization energies for an unknown element:

$I E_{1} = 800 kJ/mol$
$I E_{2} = 2420 kJ/mol$
$I E_{3} = 3660 kJ/mol$
$I E_{4} = 25000 kJ/mol$ The large jump occurs between $I E_{3}$ and $I E_{4}$ , indicating the element has three valence electrons. This places the element in Group 13.

Common Mistake

Students often mistake small increases in ionization energy as significant jumps. Always look for a substantial increase, typically several times larger than the previous step.

Graphical Trends in Ionization Energy

Trends Across a Period

As you move across a period in the periodic table (from left to right), the first ionization energy generally increases. This happens because:

The nuclear charge (number of protons) increases, pulling the electrons closer to the nucleus.
Shielding remains relatively constant as electrons are added to the same energy level.

Example

Example 2: First ionization energy values for Period 2 elements:

Lithium $L i$ : $520 kJ/mol$
Beryllium $B e$ : $900 kJ/mol$
Boron $B$ : $800 kJ/mol$
Carbon $C$ : $1100 kJ/mol$
Nitrogen $N$ : $1400 kJ/mol$
Oxygen $O$ : $1300 kJ/mol$
Fluorine $F$ : $1680 kJ/mol$
Neon $N e$ : $2080 kJ/mol$ Notice the general increase, with slight dips at boron and oxygen due to electron repulsion effects (discussed below).

Trends Down a Group

As you move down a group in the periodic table, the first ionization energy decreases:

The outermost electrons are farther from the nucleus due to additional energy levels.
Increased shielding by inner electrons reduces the effective nuclear charge felt by the valence electrons.

Tip

When analyzing trends, always consider both nuclear charge and shielding effects to explain the observed patterns.

Graphical Representation of Trends

When ionization energy is graphed against atomic number, the periodic nature of the elements becomes clear.
Peaks correspond to noble gases (stable configurations with high $I E_{1}$ ), while troughs correspond to alkali metals (low $I E_{1}$ ).

Graph of ionization energies against atomic number.

Discontinuities in Trends

Two notable exceptions occur in the general trend across a period:

Group 2 to Group 13 (e.g., Be to B):
- Beryllium $B e$ : $1 s^{2} 2 s^{2}$
- Boron $B$ : $1 s^{2} 2 s^{2} 2 p^{1}$
- The 2p electron in boron is higher in energy and shielded by the 2s electrons, making it easier to remove than the 2s electron in beryllium.
Group 15 to Group 16 (e.g., N to O):
- Nitrogen $N$ : $1 s^{2} 2 s^{2} 2 p^{3}$ (half-filled p sublevel, stable)
- Oxygen $O$ : $1 s^{2} 2 s^{2} 2 p^{4}$ (paired electrons in the 2p sublevel experience repulsion)
- The increased repulsion in oxygen’s 2p sublevel makes it slightly easier to remove an electron compared to nitrogen.

Analogy

Think of nitrogen’s half-filled sublevel like a balanced seesaw—it’s inherently stable. Adding another electron (as in oxygen) disrupts the balance, making it easier to remove one of the paired electrons.

Reflection and Practice

Self review

What is the relationship between successive ionization energy and an element’s group number?
Why does ionization energy generally increase across a period but decrease down a group?
How can you identify a discontinuity in ionization energy trends, and what does it signify about electron configurations?

Theory of Knowledge

How does the periodic trend in ionization energy relate to the reactivity of elements in everyday materials, such as the metals used in batteries or the non-metals used in fertilizers?

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