Ionic bonding
Section: 3 Chemical Bonding | Syllabus: Cambridge AS Level Physics 9702
Electronegativity Electronegativity The power of an atom to attract electrons to itself (specifically, the bonding pair of electrons in a covalent bond). Electronegativity is measured on the Pauling scale , where fluorine (the most electronegative element) is assigned a value of 4.0.
Factors Affecting Electronegativity Nuclear charge: More protons → greater attraction on bonding electrons → higher electronegativity. Atomic radius: Larger atom → bonding electrons are further from the nucleus → weaker attraction → lower electronegativity.
Shielding by inner shells and sub-shells: More inner electrons shield the nucleus → bonding electrons experience a smaller effective nuclear charge → lower electronegativity. Trends in Electronegativity Across a period (left → right): Electronegativity increases .
Nuclear charge increases while shielding remains roughly constant, so the effective nuclear charge experienced by bonding electrons increases. Down a group (top → bottom): Electronegativity decreases .
Each element has an extra shell, increasing atomic radius and shielding - the nucleus attracts bonding electrons less strongly. Noble gases are not assigned electronegativity values as they do not normally form bonds.
Figure 3.1: Electronegativity Trends on the Periodic Table (Periodic table heatmap showing electronegativity values. Colour gradient: low values (blue) at bottom-left (Fr, Cs) increasing to high values (red/orange) at top-right (F, O, N).
Arrow annotations: "increases across period →" and "increases up group ↑".) Predicting Bond Type from Electronegativity The difference in Pauling electronegativity values (ΔEN) between two bonded atoms indicates the character of the bond: ΔEN (Pauling) Bond character Example 0 Pure covalent (equal sharing) Cl–Cl, H–H 0 – 1.7 Polar covalent H–Cl (ΔEN = 0.9), H–F (ΔEN = 1.9 - borderline) > 1.7 Ionic Na–Cl (ΔEN = 2.1), Mg–O (ΔEN = 2.4) Key Point The boundary between ionic and covalent is not sharp - it is a continuum.
As ΔEN increases, ionic character increases. Apart from homonuclear diatomic molecules (e.g. Cl₂), all bonds have some degree of polarity. Ionic Bonding Ionic Bonding The electrostatic attraction between oppositely charged ions - positively charged cations and negatively charged anions.
Ionic bonds form when electrons are transferred from a metal atom to a non-metal atom. The metal loses electrons to form a cation; the non-metal gains electrons to form an anion. The resulting ions attract each other electrostatically.
Example 1: Sodium Chloride, NaCl Na (2,8,1) loses 1 electron → Na⁺ (2,8) - noble gas configuration Cl (2,8,7) gains 1 electron → Cl⁻ (2,8,8) - noble gas configuration ΔEN = 3.0 − 0.9 = 2.1 → ionic bond 1 Na⁺ : 1 Cl⁻ → formula NaCl Figure 3.2: Dot-and-Cross Diagram - NaCl (Na atom losing its outer electron (shown as a dot) to Cl.
Na⁺ shown with empty outer shell in square brackets with charge +. Cl⁻ shown with full outer shell of 8 electrons (dots and crosses) in square brackets with charge −. Only outer shells need be shown.) Example 2: Magnesium Oxide, MgO Mg (2,8,2) loses 2 electrons → Mg²⁺ (2,8) O (2,6) gains 2 electrons → O²⁻ (2,8) ΔEN = 3.4 − 1.2 = 2.2 → ionic bond 1 Mg²⁺ : 1 O²⁻ → formula MgO MgO has a higher melting point than NaCl because of the greater charge on both ions (2+ and 2−), giving stronger electrostatic attraction.
Figure 3.3: Dot-and-Cross Diagram - MgO (Mg²⁺ with empty outer shell in square brackets [Mg]²⁺. O²⁻ with 8 electrons (4 pairs) in outer shell in square brackets [O]²⁻. Label the 2 electrons transferred from Mg to O.) Example 3: Calcium Fluoride, CaF₂ Ca (2,8,8,2) loses 2 electrons → Ca²⁺ (2,8,8) Each F (2,7) gains 1 electron → F⁻ (2,8) 1 Ca²⁺ requires 2 F⁻ to balance charge → formula CaF₂ Figure 3.4: Dot-and-Cross Diagram - CaF₂ (Ca²⁺ in centre with square brackets [Ca]²⁺.
Two F⁻ ions on either side, each with 8 electrons in outer shell in square brackets [F]²×−. Show electrons transferred from Ca to each F atom.) Giant Ionic Lattice Structure In ionic compounds, ions arrange into a giant ionic lattice - a regular 3D arrangement where each ion is surrounded by oppositely charged ions.
In NaCl: each Na⁺ is surrounded by 6 Cl⁻ ions, and each Cl⁻ is surrounded by 6 Na⁺ ions. The lattice is held together by strong electrostatic attractions in all directions. Figure 3.5: Giant Ionic Lattice of NaCl (3D lattice diagram with Na⁺ (small, green) and Cl⁻ (large, grey) alternating in a cubic arrangement.
Label: each Na⁺ surrounded by 6 Cl⁻; each Cl⁻ surrounded by 6 Na⁺. Show the repeating unit.) Exam Tip In dot-and-cross diagrams for ionic compounds, show only the outer (valence) shell electrons. Use dots for one atom and crosses for the other.
Put each ion in square brackets with the charge written outside.
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