Electronegativity & Ionisation Energy – HSC Chemistry


This is part of preliminary HSC Chemistry course under the topic of Periodicity.

HSC Chemistry Syllabus

  • Demonstrate, explain, and predict the relationships in the observable trends in the physical and chemical properties of elements in periods and groups in the periodic table, including but not limited to:
– State of matter at room temperature
– Electron configurations and atomic radii
– First ionisation energy and electronegativity
– Reactivity with water


      For Year 11 HSC Chemistry students, understanding the concepts of electronegativity and ionisation energy, along with their periodic trends, is fundamental in mastering the principles of chemical bonding and reactivity. These properties are not just abstract numbers; they are indicators of an element's behaviour in chemical processes. 


      Electronegativity is a measure of an atom's ability to attract and hold onto electrons in a chemical bond.

      It's a dimensionless quantity, often quantified on the Pauling scale, where fluorine (F) is assigned the highest value of 4.0, being the most electronegative element.

      Electronegativity Trends in the Periodic Table

      Electronegativity Trend in Periodic Table


      • Across a Period: Electronegativity increases from left to right across a period. This is because atoms have more protons (increasing the nuclear charge) and the same number of core electrons (electrons in shells closer to the nucleus), pulling the valence electrons closer and making the atoms more effective at attracting bonding electrons.
      • Down a Group: Electronegativity decreases down a group. As the atomic number increases, so does the number of electron shells. This increases the distance between the nucleus and the valence electrons, reducing the nuclear attraction experienced by bonding electrons.

      Electronegativity in Chemical Bonds

      The difference in electronegativity between atoms in a bond determines the bond's character, and therefore is an important trend to understand.

      • Small Difference: Covalent bonding, where electrons are shared.
      • Large Difference: Ionic bonding, where electrons are transferred from one atom to another.

      Electronegativity vs Electron Affinity

      While both electronegativity and electron affinity describe an atom's attraction for electrons, they are not the same:

      • Electronegativity is a measure of an atom's ability to attract electrons when it is part of a compound.
      • Electron Affinity is the energy change that occurs when an atom in the gas phase accepts an electron to form an anion. It's a more specific measure, often with values expressed in kJ/mol. High electron affinity means the atom releases more energy (more negative value) upon gaining an electron, indicating a strong attraction for additional electrons.

      Noble Gases

      Noble gases are a unique group in the periodic table due to their complete valence electron shells, which confer remarkable stability. This stability significantly influences their electronegativity and electron affinity values.

      In general, noble gases have relatively high electronegativity values compared to other groups, especially when considering their positions on the far right of the periodic table. However, it's important to note that electronegativity is most relevant in the context of chemical bonding, and since noble gases rarely form bonds due to their full valence shells, the concept of electronegativity is less applicable to them than to other elements. In fact, noble gases are often not assigned electronegativity values in many scales, including the Pauling scale. 

      Electron affinity measures the energy change when an electron is added to a neutral atom in the gas phase to form an anion. Noble gases have low (often considered negligible or even positive) electron affinities. This is because their valence shells are full, meaning they do not have a tendency to gain electrons. Adding an electron to a noble gas would disrupt its stable electronic configuration, requiring energy rather than releasing it. For this reason, noble gases are not inclined to form anions, and their electron affinities do not follow the same trend as seen in other groups of the periodic table where a higher electron affinity indicates a stronger attraction for electrons.

      Ionisation Energy

      Ionisation energy is the energy required to remove an electron from an atom in the gas phase. It is a critical indicator of an element's reactivity and is usually expressed in kilojoules per mole (kJ/mol).

      • First ionisation energy is the energy needed to remove the first electron from a neutral atom. It provides insight into how tightly an atom holds its outermost electron.
      • Second ionisation energy is the energy required to remove a second electron after the first has been removed. The second ionisation energy is always higher than the first because, after losing an electron, the ion's effective nuclear charge on the remaining electrons increases, making them more difficult to remove.

      Ionisation Energy Trends in the Periodic Table

      Ionisation energy trend in periodic table


      • Across a Period: Ionisation energy increases from left to right. Atomic radius decreases, and electrons are more strongly attracted to the nucleus, causing them to require more energy to remove. However, there are exceptions to this rule (see below for examples).
      • Down a Group: Ionisation energy decreases down a group. The valence electrons become farther from the nucleus and are more easily removed due to less effective nuclear charge.

      Example – First & Second Ionisation of Sodium

      The first and second ionisation energies of sodium (Na) are 496 and 4560 kJ/mol respectively. Why is there such a substantial difference?

      The first ionisation energy of sodium is the energy required to remove the outermost, or highest-energy, electron from a neutral sodium atom in the gas phase. Sodium's electron configuration is `[Ne] 3s^1`, with one electron in the outermost `3s` orbital. The process can be represented as:


      $$Na \rightarrow Na^+ + e^-$$


      This first ionisation energy is relatively low compared to many other elements, which is consistent with sodium's position in Group 1 of the periodic table. Elements in this group have a single electron in their outermost shell, which is relatively easy to remove due to the electron being further from the nucleus and experiencing less electrostatic attraction from the positively charged nucleus, especially because of the shielding effect of the inner electrons.

      The second ionisation energy of sodium is the energy required to remove a second electron after the first has already been removed, turning the Na⁺ ion into Na²⁺. This process involves removing an electron from the completed `2p` shell, which is much closer to the nucleus and fully shielded by only the `1s` and `2s` electrons.

      The process can be represented as:


      The second ionisation energy is significantly higher than the first because:

      • The electron being removed from a positively charged ion feels a stronger effective nuclear charge since there are fewer electrons shielding the positive charge of the nucleus.
      • The electron is also being removed from a completely filled `2p` orbital, which would disrupt a stable electron configuration.
      • The `2p` orbital is closer to the nucleus than the `3s` orbital where the first electron was removed. This means the electron is more strongly attracted to the nucleus.

      Example – Ionisation Energy of Magnesium and Aluminium


      The ionisation energy of an element is influenced by several factors, including atomic size, nuclear charge, and electron shielding. When comparing aluminium (Al) to magnesium (Mg), it might initially seem counterintuitive for aluminium to have a lower first ionisation energy because both elements are in the same period (period 3 of the periodic table), and aluminium has more protons in its nucleus, suggesting a stronger attraction to its electrons.

      Magnesium has an electron configuration of `[Ne] 3s^2`, with its outermost electrons in the `3s` orbital. Aluminium has an electron configuration of . The outermost electron of aluminium is in the `3p` orbital, which is at a higher energy level than the `3s` orbital.

      The `3s` electrons in magnesium are closer to the nucleus and experience less shielding from the inner electrons than the `3p` electron in aluminium. This makes the `3s` electrons more tightly bound to the nucleus.

      In aluminium, the `3p` electron is not only further from the nucleus due to being in a higher subshell but also experiences more shielding from the `3s^2` electrons. This reduces the effective nuclear charge experienced by the `3p` electron, making it easier to remove compared to the `3s` electrons in magnesium.


      Previous Section: Physical Properties & Types of Elements

      Next Section: Compound Structure