the ionic model
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What is an ion?
An ion is a charged particle formed when an atom or group of atoms loses or gains one or more electrons. The number of protons stays the same, so the charge comes from an imbalance between protons and electrons.
What is a cation, and how is it formed?
A cation is a positive ion formed when an atom loses electrons. The number of protons is then greater than the number of electrons. Metals typically form cations.
What is an anion, and how is it formed?
An anion is a negative ion formed when an atom gains electrons. The number of electrons is then greater than the number of protons. Non-metals typically form anions.
How is the charge on an ion related to electrons lost or gained?
Charge = number of electrons lost (positive) Charge = −(number of electrons gained) Example: losing 2 e⁻ → 2⁺; gaining 3 e⁻ → 3⁻.
Why don’t valence electrons experience the full nuclear charge?
Because inner-shell electrons shield them from the full positive charge of the nucleus and repel them. As a result, valence electrons feel a reduced attractive force, called effective nuclear charge (Z_eff).
Define effective nuclear charge (Z_eff) qualitatively.
Effective nuclear charge is the net positive charge “felt” by valence electrons, after accounting for shielding by inner electrons. It is less than the full nuclear charge (atomic number).
How can effective nuclear charge for valence electrons be approximated in simple cases?
For main-group elements, a simple approximation is: Z_eff ≈ (number of protons) − (number of inner/core electrons).
Use Na as an example to estimate Z_eff for the valence electron.
Sodium: Z = 11, electron configuration 1s²2s²2p⁶3s¹. 10 inner electrons (Ne core) ≈ full shielding → Z_eff ≈ 11 − 10 = +1 for the 3s valence electron.
How does Z_eff change across Period 3 from Na to Cl, using the textbook table?
Core is always [Ne] (10 electrons). Approximate Z_eff: - Na: 11−10 = +1 - Mg: 12−10 = +2 - Al: 13−10 = +3 - Si: 14−10 = +4 - P: 15−10 = +5 - S: 16−10 = +6 - Cl: 17−10 = +7 So Z_eff increases steadily across the period
Why are metal atoms more likely to lose electrons and form positive ions?
Metals (left side of the periodic table) have: - Lower effective nuclear charge on valence electrons - Large atomic radii - Low ionisation energies So valence electrons are held weakly and are easily removed, making cation formation energetically feasible.
How does ionisation energy trend support metals forming cations?
Ionisation energy decreases down a group and is lower on the left of a period; metals, especially at the bottom left, lose electrons most readily and so are the most likely to form cations and ionic compounds.
Why are non-metal atoms more likely to gain electrons and form anions?
Non-metals on the top right (excluding noble gases) have: - High effective nuclear charge - Small atomic radii They therefore attract electrons strongly, making electron gain energetically favourable and leading to anion formation.
Why are halogens especially good at forming X⁻ ions?
Halogens (Group 17): - Have one vacancy (the number of electrons needed to fill the outer electron shell to reach a noble gas configuration.) in the outer shell - Have high Z_eff and small radius (especially F) They easily gain one electron to form X⁻, achieving a noble gas configuration.
Describe, in terms of electron configuration, how sodium and chlorine form NaCl.
Na: 1s²2s²2p⁶3s¹ → loses 1 electron → Na⁺ : 1s²2s²2p⁶ (like Ne) Cl: 1s²2s²2p⁶3s²3p⁵ → gains 1 electron → Cl⁻ : 1s²2s²2p⁶3s²3p⁶ (like Ar) - The 3s electron from Na is “harpooned” by Cl, and then Na⁺ and Cl⁻ are held together by electrostatic attraction in the lattice.
Why is NaCl formation described as a redox reaction (linked to Reactivity 3.2)?
Sodium is oxidised: Na → Na⁺ + e⁻ Chlorine is reduced: Cl₂ + 2 e⁻ → 2 Cl⁻ Overall, electrons are transferred from Na to Cl → redox process.
Why is the reaction between Cs (bottom of Group 1) and F (top of Group 17) the “most ionic”?
Cs has the lowest ionisation energy in Group 1 and loses electrons most easily. F has the highest effective nuclear charge and smallest radius in Group 17 and attracts electrons most strongly. Their combination gives maximum electron transfer and strongest ionic character.
What are successive ionisation energies, and what general pattern do they show?
Successive ionisation energies are the energies required to remove electrons one by one from the same atom/ion. They increase each time, and show a large jump when an electron is removed from a new, inner energy level.
Why does sodium form Na⁺ but not Na²⁺?
Removing the first electron (3s¹) is relatively easy, but the second ionisation energy involves removing an electron from a full inner shell (Ne core). This requires a huge increase in energy, making Na²⁺ formation energetically prohibitive under normal conditions.
Why does aluminium commonly form Al³⁺ but not Al⁴⁺?
The first three electrons removed are valence electrons (3s²3p¹). Removing a fourth would again mean taking an electron from an inner, more stable shell, causing a massive jump in ionisation energy. So Al³⁺ is stable; Al⁴⁺ is not formed in simple ionic compounds.
Why does silicon not form a simple Si⁴⁺ ion in typical chemical reactions?
Its first ionisation energy is already high, and removing four electrons would involve very high successive energies, including removal from inner shells. There is no sufficient energy compensation (e.g. lattice enthalpy of formation) to make Si⁴⁺ formation feasible in simple ionic solids.
Why does increasing ionic charge strengthen the ionic bond but also increase formation energy cost?
• Higher charge → stronger electrostatic attraction → higher lattice enthalpy. • But forming higher-charged cations requires removing more electrons, so successive ionisation energies increase sharply. Feasibility is a balance between ion formation energy and lattice (or hydration) energy.
How do Group 17 elements form X⁻ ions, and why is it favourable?
They have one vacancy in their outer shell and high Z_eff, so gaining 1 e⁻ gives a stable noble gas configuration (ns²np⁶). The process is strongly favoured by their high electron affinity.
Why can Group 16 elements form X²⁻ ions but less readily than halides form X⁻?
Group 16 elements (e.g. O) have two vacancies but lower Z_eff than halogens. They can gain 2 e⁻ to form X²⁻, but they attract electrons less strongly than the halogens, so formation is somewhat less favourable.
Why is P³⁻ feasible, but Si⁴⁻ is not?
Phosphorus can gain 3 electrons to achieve an Ar-like configuration, but the growing negative charge produces increasing electron–electron repulsion. For Si⁴⁻, adding four electrons to a neutral atom creates very strong repulsion, and the size/charge combination makes it energetically unrealistic. So Si⁴⁻ does not form.