In simple terms
A friendly intro before the formal notes — no formulas yet.
The Cost of Removing an Electron
Ionisation energy is the energy needed to pull an electron away from an atom. This 'cost' changes predictably, telling us a lot about an atom's structure.
Imagine trying to pull a small magnet off a large magnetic block. It's harder to pull it off if the main block is stronger (higher nuclear charge), and easier if the small magnet is further away (larger atomic radius) or has other magnets in between blocking the pull (shielding).
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First IE is the energy needed to remove 1 mole of the outermost electrons from 1 mole of gaseous atoms, forming 1+ ions.
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Across a period, IE generally increases. The nuclear charge gets stronger, pulling the outer shell in closer, making electrons harder to remove.
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Down a group, IE decreases. The atom gets bigger with more shells, so the outer electron is further away and better shielded from the nucleus.
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Small dips occur from Group 2 to 3 (a new, higher-energy p-subshell) and Group 5 to 6 (repulsion from paired electrons in a p-orbital).
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Defining First Ionisation Energy
The first ionisation energy (often abbreviated as 1st IE or ) is defined with precision. It is the minimum energy required to remove one mole of electrons from one mole of gaseous atoms to produce one mole of gaseous ions with a +1 charge. The state symbols are crucial; the process must occur in the gas phase to eliminate intermolecular forces.
For a generic element M, the equation is: \ M(g) → M⁺(g) + e⁻ ΔH = 1st IE
Energy is always required, so ionisation is an endothermic process (ΔH is positive).
Measured in kilojoules per mole (kJ mol⁻¹).
The lower the ionisation energy, the more easily an atom forms a positive ion.
Factors Influencing Ionisation Energy
The magnitude of the ionisation energy is determined by the strength of the electrostatic attraction between the positive nucleus and the outermost electron. Three key factors govern this attraction:
Nuclear Charge: The number of protons in the nucleus. A higher nuclear charge results in a stronger attraction for electrons, increasing the ionisation energy.
Atomic Radius: The distance from the nucleus to the outermost electron. A larger atomic radius means the electron is further away, the attraction is weaker, and the ionisation energy is lower.
Electron Shielding: The repulsion experienced by an outer electron from electrons in inner shells. More inner shells lead to greater shielding, which reduces the effective nuclear charge felt by the outer electron, lowering the ionisation energy.
Trends in First Ionisation Energy
Across a Period (e.g., Na to Ar): There is a general increase in first ionisation energy. As you move across a period, the number of protons (nuclear charge) increases, pulling the electrons in the same principal shell closer to the nucleus. Although the number of electrons also increases, they are added to the same shell, so shielding does not change significantly. The dominant effect is the increased nuclear charge, leading to a stronger attraction that requires more energy to overcome.
Down a Group (e.g., Li to Cs): The first ionisation energy decreases. Moving down a group, the nuclear charge increases, but this is outweighed by two other factors. Firstly, the number of electron shells increases, so the atomic radius becomes larger. Secondly, the increased number of inner shells provides more shielding. Both effects weaken the attraction between the nucleus and the outermost electron, making it easier to remove.
Anomalies in the Trend Across a Period
The general increase across a period is not perfectly smooth. There are two characteristic 'dips' or anomalies.
Drop from Group 2 to Group 13 (e.g., Be to B): Beryllium (1s²2s²) has a higher 1st IE than Boron (1s²2s²2p¹). This is because the outermost electron in Boron is in a 2p orbital, which is at a slightly higher energy level and further from the nucleus than the 2s orbital. The 2s electrons also provide some shielding for the 2p electron. Both factors make the 2p electron easier to remove.
Drop from Group 15 to Group 16 (e.g., N to O): Nitrogen (1s²2s²2p³) has a higher 1st IE than Oxygen (1s²2s²2p⁴). In Nitrogen, the three 2p electrons are in separate orbitals (Hund's rule). In Oxygen, the fourth 2p electron must pair up in one of the orbitals. This pairing causes electron-electron repulsion (spin-pair repulsion), making it energetically easier to remove one of these paired electrons compared to removing an electron from a half-filled orbital in nitrogen.
Successive Ionisation Energies
Successive ionisation energies refer to the energy required to remove each subsequent electron. The second ionisation energy is the energy to remove an electron from a 1+ ion, the third from a 2+ ion, and so on. These values always increase for a given element because each electron is being removed from an increasingly positive particle, which exerts a stronger electrostatic pull. Large jumps in the values of successive ionisation energies provide powerful evidence for the existence of electron shells. A jump occurs when an electron is removed from an inner shell, which is closer to the nucleus and experiences less shielding.
2nd IE: M⁺(g) → M²⁺(g) + e⁻ \ 3rd IE: M²⁺(g) → M³⁺(g) + e⁻
Worked examples
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Explain why the first ionisation energy of magnesium (12Mg) is greater than that of sodium (11Na).
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Identify electron configurations: Na is [Ne] 3s¹ and Mg is [Ne] 3s². Both elements are in Period 3.
The first five successive ionisation energies of an element X are 738, 1451, 7733, 10543, and 13630 kJ mol⁻¹. To which group of the periodic table does element X belong? Explain your answer.
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Analyse the data: Look for large jumps between consecutive values.
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Define First Ionisation Energy.
The energy required to remove one mole of electrons from one mole of gaseous atoms to form one mole of gaseous 1+ ions. Units: kJ mol⁻¹.
Key takeaways
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Energy is always required, so ionisation is an endothermic process (ΔH is positive).
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Measured in kilojoules per mole (kJ mol⁻¹).
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The lower the ionisation energy, the more easily an atom forms a positive ion.
Practice — then mark it
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9701/22 · Q2(b)(i)
Complete Table 2.1 to give details of the reactions of some Period 3 oxides with H2O.
9701/22 · Q3(d)
Complete Table 3.1 to show the properties of nitrogen and phosphorus in their standard states.
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