Ionisation energy is the energy required to remove an electron from a gaseous atom or ion. As one progresses across a period in the periodic table, there is a general trend of increasing first ionisation energy. However, there are specific points where this trend shows discontinuities. Such irregularities offer significant insights into the electronic structure of atoms and indicate the existence of different energy sublevels.
Ionisation Energy: A Recap
Ionisation energy is fundamentally linked to the electronic configuration of atoms. A quick recap:
- First Ionisation Energy: The energy needed to remove the outermost electron from a neutral atom.
- Subsequent Ionisation Energies: The energy required to remove additional electrons. Typically, these energies increase, as inner-shell electrons are closer to the nucleus and experience a stronger force of attraction.
Image courtesy of Watthana Tirahimonch
Trends in Ionisation Energy Across a Period
As one moves from left to right across a period:
- Number of protons in the nucleus increases.
- The atomic radius decreases because of an increased positive charge, pulling electrons closer to the nucleus.
- The electron shielding effect remains relatively constant, as electrons are added to the same principal energy level.
- Overall, these factors result in a general increase in ionisation energy.
Trends in ionization energy. Showing a change in the atomic radius because of an increased positive charge, pulling electrons closer to the nucleus.
Image courtesy of Surfguppy
Discontinuities in the Trend
Despite the general increase in ionisation energy, there are notable exceptions, particularly between groups 2 and 3, and between groups 15 and 16. To comprehend these, we must delve into the arrangement of electrons in various subshells.
Between Groups 2 and 3
Elements in Group 2: (e.g., beryllium and magnesium) have their outer electrons in the s-sublevel. Removing an electron from a filled s-sublevel requires considerable energy.
Ionization energies of group 2 elements.
Image courtesy of Albris
Elements in Group 3: (e.g., boron and aluminium) start filling the p-sublevel. The p-sublevel is slightly higher in energy than the s-sublevel of the same principal energy level. This means the first electron in the p-sublevel is somewhat easier to remove than the last electron in the s-sublevel.
Between Groups 15 and 16
Elements in Group 15: (e.g., nitrogen and phosphorus) have half-filled p-sublevels. There's a certain stability associated with half-filled sublevels.
Elements in Group 16: (e.g., oxygen and sulfur) introduce an additional electron to a previously half-filled p-sublevel. This results in electron-electron repulsion, making it slightly easier to remove the first electron than anticipated.
Difference between the first and second ionization energies of different elements.
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Focus on the Electron Removed
While it's tempting to attribute the aforementioned discontinuities solely to the "special stability" of filled and half-filled sublevels, the crux lies in the energy of the electron removed. As electrons fill up different sublevels, their energy, and consequently the energy required to remove them, varies.
- s-sublevel electrons: Are closer to the nucleus and are more strongly attracted, so they require more energy to remove.
- p-sublevel electrons: Are further away and experience slightly more shielding, so they may require less energy to remove, especially when starting to fill up the sublevel.
Implications for Electronic Configurations
These discontinuities in ionisation energies corroborate the existence of sublevels in electron configurations. They further affirm the arrangement of sublevels and the order in which they are filled.
- The transition from a filled s-sublevel to the beginning of a p-sublevel marks a distinct change in ionisation energy, reflecting the reality of the s and p sublevels' different energies.
- Similarly, the relative stability of a half-filled p-sublevel is evident from the ionisation energy trend.
To conclude, ionisation energy is not merely a reflection of atomic size or nuclear charge. It is intricately tied to the nuances of electron configurations and sublevel energies. Discontinuities in the trend across periods in the periodic table underscore the layered complexity of atomic structure and serve as a testament to the detailed choreography of electrons in various orbitals.
FAQ
The concept of "special stability" is most pronounced in the s and p sublevels due to their symmetrical electron distributions. For d sublevels, while there's some stability associated with half-filled (d5) and fully-filled (d10) configurations, the effect is less pronounced. This is because d orbitals are more diffuse, experience more shielding from inner electrons, and the energy difference between successive d orbitals is relatively small. Thus, while there's some added stability in specific d sublevel configurations, the effect on ionisation energy is not as marked as in s and p sublevels.
It can, but it's essential to consider the broader context. While "special stability" associated with half-filled or fully-filled sublevels can lead to higher ionisation energies, other factors also influence ionisation energy, such as atomic radius, shielding effect, and nuclear charge. A higher than expected ionisation energy might indicate "special stability," but it's crucial to examine the electron configuration, position in the periodic table, and other atomic properties before drawing conclusions.
Yes, discontinuities in ionisation energy can be used as indirect evidence for electron configurations. For instance, the noticeable drop in ionisation energy between Group 2 and Group 3 elements supports the idea that after filling the 2s sublevel, electrons start populating the 2p sublevel. As electron configurations dictate the distribution of electrons among various sublevels and shells, the trends and discontinuities observed in ionisation energies can be correlated to the arrangement of electrons in an atom, reinforcing our understanding of electron configurations.
The discontinuities in ionisation energy, especially between s and p sublevels, can give insights into the chemical reactivity of elements. Elements with lower ionisation energies tend to lose electrons more easily, indicating they can act as reducing agents or metals. In contrast, those with higher ionisation energies are more likely to accept electrons, displaying non-metallic or oxidising properties. The sudden drop in ionisation energy between Group 2 and Group 3 elements signals a transition in electron filling, which is linked to a shift in chemical properties from metallic to metalloid behaviour.
The d and f sublevels are both part of the higher principal energy levels. By the time electrons start filling the d and f sublevels, the atom already has several inner electron shells. These inner electron shells provide significant shielding, reducing the effective nuclear charge experienced by the outermost electrons. Additionally, the f sublevel is more diffused and further from the nucleus, and electrons in the f-sublevel generally experience lesser nuclear attraction than those in the d-sublevel. Due to these factors, the ionisation energies of d and f sublevels do not show pronounced discontinuities as observed between s and p sublevels.
Practice Questions
The drop in ionisation energy between elements in Group 2 and Group 3 is attributed to the transition in electron filling from the s-sublevel to the p-sublevel. Elements in Group 2, like beryllium and magnesium, have a filled s-sublevel. Removing an electron from this full sublevel requires considerable energy. In contrast, elements in Group 3, such as boron and aluminium, begin filling the p-sublevel, which is slightly higher in energy and more shielded. Consequently, the first electron in the p-sublevel is easier to remove than the last electron in the s-sublevel, leading to the observed drop in ionisation energy.
It's crucial to focus on the energy of the electron being removed because ionisation energy is directly related to how tightly an electron is held in an atom. While the "special stability" of filled and half-filled sublevels does play a role in ionisation energy trends, the primary factor is the energy of the specific electron in question. The s-sublevel electrons are generally closer to the nucleus and experience a stronger attraction, making them harder to remove. In contrast, the p-sublevel electrons, especially when starting to fill the sublevel, are more shielded and slightly easier to remove. This nuance in electron energy is fundamental to understanding the observed discontinuities.
