ionic bond

ypes of Chemical Bonds and Ionic Bond Formation

• Metals tend to form cations by losing electrons due to their low ionization energy, while non-metals tend to form anions by gaining electrons because of their high electron affinity. This difference explains why ionic bonds form between metals and non-metals, whereas non-metals form covalent bonds by sharing electrons, and metals form metallic bonds by pooling electrons. There is a gradual change from metallic to non-metallic character moving left to right across a period and bottom to top in groups on the periodic table

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• An ionic bond is the electrostatic attraction between positively charged cations and negatively charged anions formed by the transfer of electrons from one atom to another. For example, sodium transfers one electron to chlorine, forming Na+ and Cl− ions that attract each other to form NaCl

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• Ionic solids form large, regular crystalline lattices where each ion is surrounded by multiple ions of opposite charge to maximize electrostatic attraction and stability

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Lewis Electron Dot Symbols and Structures

• Lewis symbols represent atoms with their chemical symbol surrounded by dots indicating valence electrons, facilitating visualization of electron transfer or sharing in bonding. For example, chlorine's symbol is Cl with seven dots around it representing its 7 valence electrons

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• Lewis structures combine Lewis symbols to show the formation of ionic or covalent bonds, such as Na donating an electron to Cl to form Na+ and Cl− ions, represented symbolically by their electron dots

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Energy Considerations in Ionic Bond Formation

• The electron transfer process to form ions requires energy (endothermic), such as ionization energy for metals losing electrons and electron affinity for non-metals gaining electrons. For example, Li losing an electron requires +520 kJ/mol, while F gaining an electron releases −328 kJ/mol, resulting in a net energy input for ion formation

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• The formation of the ionic solid lattice releases a large amount of energy (exothermic), called lattice energy, which compensates for the energy absorbed in ion formation, making the overall process energetically favorable and the ionic compound stable

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• The Born–Haber cycle uses Hess’s law to calculate lattice energy indirectly by summing enthalpy changes in steps such as atomization, ionization, electron affinity, and lattice formation, as demonstrated for NaCl, whose lattice energy is −786 kJ/mol

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Factors Affecting Ionic Bond Formation and Lattice Energy

• Ionization energy (IE): Lower IE of metals favors cation formation, promoting ionic bonding.

• Electron affinity (EA): Higher EA of non-metals favors anion formation, facilitating ionic bonding.

• Lattice energy (U): The energy released when gaseous ions form an ionic solid; it depends directly on the product of ionic charges and inversely on the distance between ions:

  U∝q1×q2rU∝rq1​×q2​​

  where q1q1​ and q2q2​ are ion charges and rr is the internuclear distance. Smaller ions with higher charges have greater lattice energies, leading to stronger ionic bonds

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• Ionic compounds with higher lattice energies are more stable, have higher melting points, and are less soluble. For example, MgO (with ions of similar size to LiF but higher charges) has stronger ionic bonding than LiF

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Exceptions and Extensions to the Octet Rule in Ionic Compounds

• Some atoms, especially near helium in the periodic table (H, Li, Be, B), form ions with less than an octet (e.g., Li+, Be2+), achieving stable configurations with two electrons like He, though they may form additional bonds when possible (e.g., BF3 forms BF4−)

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• Transition and post-transition metals often do not obey the octet rule but follow the 18-electron rule due to the involvement of d orbitals. For example, Fe2+ and Fe3+ ions have valence configurations not isoelectronic with noble gases, but Zn2+ and Ga3+ ions have filled d10 subshells and noble gas cores

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• Electrons are lost first from orbitals with the highest principal quantum number (n) when forming positive ions in transition and post-transition metals, explaining deviations from the octet rule

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Summary Table: Key Concepts of Ionic Bonding

Illustrative Example: Formation of NaCl via Born–Haber Cycle

Using Hess’s law:

U=ΔHf∘−(ΔHstep1+ΔHstep2+ΔHstep3+ΔHstep4)U=ΔHf∘​−(ΔHstep1​+ΔHstep2​+ΔHstep3​+ΔHstep4​)

Calculates lattice energy U=−786 kJ/molU=−786kJ/mol

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