The study of successive ionisation energies provides critical insights into the electron configurations of elements and their positions within the periodic table. Through the analysis of this data, we can identify patterns that offer an understanding of various elemental properties and behaviours.
Recognising Successive Ionisation Energies and Electron Configuration
- Successive Ionisation Energies: Refers to the energy required to remove each electron from an atom, one after the other.
- The first ionisation energy (1st IE) is the energy needed to remove the first, or outermost, electron.
- The second ionisation energy (2nd IE) is the energy to remove the second electron, and so on.
- Electron Configuration Insights:
- The big jumps in successive ionisation energies indicate the start of a new shell. For instance, if the 3rd IE is significantly larger than the 2nd IE, it suggests that the atom originally had 2 electrons in its outermost shell.
Deducting the Group of an Element
- By analysing the number of electrons in the outer shell (determined from where the large jumps in IEs occur), we can deduce an element's group in the periodic table.
- Elements with one outer electron are in Group 1.
- Two outer electrons indicate Group 2, and so forth.
- For instance, magnesium has the electron configuration of 2,8,2. The first two ionisation energies of magnesium are relatively low. However, the third ionisation energy is significantly higher. This indicates the removal of an electron from a closer and more tightly bound shell. This pattern suggests that magnesium is in Group 2.
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Databases and Graphs of Trends in Ionisation Energies
- Importance of Databases:
- Databases allow us to compile and compare ionisation energies of various elements, providing a holistic view of trends across the periodic table.
- These trends can be graphically represented for better clarity and ease of understanding.
- Graphical Representations:
- Graphs plotting successive ionisation energies against the number of electrons removed provide clear visuals of the "jumps" in energy, signifying the commencement of ionising a new electron shell.
- Such graphical trends are paramount for students to visualise and comprehend the abstract concept of electron shells and energy levels.
Successive Ionisation Energies of Transition Elements
- Variable Oxidation States:
- Transition elements display a variety of oxidation states in their compounds.
- The study of their successive ionisation energies offers an explanation for this phenomenon.
- Patterns in Successive IEs:
- For transition metals, electrons are removed from the 4s subshell before the 3d subshell, despite 4s filling before 3d. This is due to the proximity in energy levels between 3d and 4s.
- The relatively small increases in successive ionisation energies, especially when removing 3d electrons, demonstrate the multiple stable charged states of transition metals. This results in their variable oxidation states.
- For example, manganese, with an electron configuration of [Ar] 4s² 3d⁵, can exhibit oxidation states from +2 to +7 due to the relatively small energy differences in successive ionisation energies when removing 3d electrons.
Trends of ionization energy for 1st, 2nd and 3rd-row transition metals.
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FAQ
Transition metals have a wide range of oxidation states because of their unique electron configuration. They possess electrons in both the s and d subshells, which are close in energy. This means that electrons from both of these subshells can be involved in bond formation. As a result, transition metals can lose varying numbers of electrons, leading to multiple possible charges or oxidation states. Additionally, the subtle energy differences between the 3d and 4s orbitals allow for the removal of either s or d electrons, or both, further contributing to the variable oxidation states.
The first ionisation energy is always lower than the second because, in the first ionisation, an electron is removed from a neutral atom, while in the second ionisation, an electron is removed from a positively charged ion. This means that during the second ionisation, the remaining electrons are held more tightly to the nucleus due to the increased positive charge, and thus, it requires more energy to remove another electron. Essentially, with each successive ionisation, the atom becomes more positively charged, increasing the difficulty of removing additional electrons.
Ionisation energy and atomic size are inversely proportional. As the atomic size (or atomic radius) increases, ionisation energy generally decreases. This is because, in larger atoms, the outermost electrons are farther from the nucleus and are more shielded by inner-shell electrons. Both factors reduce the force of attraction between the valence electron and the nucleus, making it easier to remove the electron and thus decreasing the ionisation energy. Conversely, smaller atoms have their outermost electrons closer to the nucleus, which results in a stronger attraction and higher ionisation energies.
The significant jump in successive ionisation energies when moving from valence electrons to an inner shell is primarily due to three reasons. Firstly, electrons in inner shells are closer to the nucleus, thus experiencing a stronger attractive force. Secondly, these inner shell electrons are less shielded from the nuclear charge than valence electrons, making them harder to remove. Finally, removing an electron from a completely filled or half-filled stable shell requires extra energy. These combined effects mean that ionising an electron from an inner shell necessitates considerably more energy than from an outer shell.
As you move across a period from left to right, ionisation energy generally increases. This is due to an increasing nuclear charge which pulls the electrons closer to the nucleus, making them harder to remove. Additionally, there's minimal change in shielding effect across a period as electrons are added to the same energy level. Conversely, as you move down a group, ionisation energy typically decreases. This is because of the addition of energy levels, increasing the atomic size, and the increased shielding by the inner electrons. The outermost electron, in this case, is farther from the nucleus and is more easily ionised.
Practice Questions
The group of an element in the periodic table can be deduced from its successive ionisation energies by identifying where a significant jump in energy occurs. The significant increase in ionisation energy indicates the start of ionising electrons from a new, inner shell. The number of electrons removed before this jump corresponds to the number of valence electrons, thus determining its group in the periodic table. For instance, if the third ionisation energy is notably higher than the second, it suggests the atom originally had two electrons in its outermost shell, placing it in Group 2.
The successive ionisation energies of transition elements reveal unique patterns due to the removal of electrons from both the 4s and 3d subshells. Electrons are first removed from the 4s subshell before the 3d, even though 4s fills before 3d. This is attributed to the proximity of energy levels between 3d and 4s. The small increases in successive ionisation energies when removing 3d electrons highlight the multiple stable charged states of transition metals. Consequently, this results in variable oxidation states for transition elements. For example, manganese can exhibit oxidation states ranging from +2 to +7 due to the subtle energy differences in its successive ionisation energies when extracting 3d electrons.
