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 ($IE$) is the minimum energy required to remove an electron from a gaseous atom in its ground state.
- The first ionization energy ($IE_1$) refers to the energy needed to remove the first electron.
- The second ionization energy ($IE_2$) is the energy required to remove the second electron, and so on.
These processes can be represented as follows:
- For the first ionization: $$\text{X(g)} + \text{energy} \rightarrow \text{X}^+ \text{(g)} + e^-$$
- For the second ionization: $$\text{X}^+ \text{(g)} + \text{energy} \rightarrow \text{X}^{2+} \text{(g)} + e^-$$
Successive ionization energies always increase because, as electrons are removed, the remaining electrons experience a stronger electrostatic attraction to the positively charged nucleus.
NoteIonization energy is measured for gaseous atoms to ensure that interactions between atoms don’t affect the energy values.
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:
- $IE_1 = 500 \, \text{kJ mol}^{-1}$
- $IE_2 = 4560 \, \text{kJ mol}^{-1}$
- $IE_3 = 6910 \, \text{kJ mol}^{-1}$
- A massive jump between $IE_1$ and $IE_2$ 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.
- Successive ionization energies for an unknown element:
- $IE_1 = 800 \, \text{kJ mol}^{-1}$
- $IE_2 = 2420 \, \text{kJ mol}^{-1}$
- $IE_3 = 3660 \, \text{kJ mol}^{-1}$
- $IE_4 = 25000 \, \text{kJ mol}^{-1}$
- The large jump occurs between $IE_3$ and $IE_4$, indicating the element has three valence electrons.
- This places the element in Group 13.
- 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.
