Chemistry: Atoms First 2e · Chemical Bonding and Molecular Geometry
Ionic Bonding
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An ionic bond Electrostatic attraction between oppositely charged ions is the electrostatic attraction between oppositely charged ions. It forms when a metal transfers one or more electrons to a nonmetal: the metal becomes a cation Positive ion formed when a metal loses electrons, the nonmetal an anion Negative ion formed when a nonmetal gains electrons Full entry →, and opposite charges hold them together. Sodium chloride is the classic example — Na⁺ and Cl⁻ pack into an extended crystal lattice Repeating 3-D ion arrangement Full entry → rather than discrete pairs.
This topic explains why electron transfer happens, what determines bond strength (Coulomb's law E = kQ1Q2/d Full entry →), and why ionic solids behave as they do — high melting points, brittleness, and conductivity only when molten or dissolved. It also notes that real bonds lie on a continuum, not simply "ionic" or "covalent."
Why this matters
Ionic bonding explains materials you rely on every day:
- Body chemistry. Nerve and muscle function depend on Na⁺ and K⁺ moving across membranes; bone mineral is an ionic solid; rehydration salts deliver Na⁺, K⁺, Cl⁻.
- Materials. Ceramics and glass are hard and heat-resistant due to strong ion–ion attractions; MgO lines furnaces (melts above 2800 °C).
- Safety and handling. Knowing ionic solids conduct only when molten or dissolved prevents dangerous assumptions about wires and batteries; brittleness explains why salts shatter rather than bend.
- Exams. Lattice-energy trends — which compound melts higher, which ion pair attracts more strongly — are a favorite way to test Coulomb's law reasoning.
The college version
Core Concepts
How an ionic bond forms
A sodium atom, [Ne]3s1, can achieve the stable neon configuration by losing its one valence electron; a chlorine atom, [Ne]3s2 3p5, can achieve the argon configuration by gaining one:
Na → Na+ + e- and Cl + e- → Cl-
The electrostatic attraction between the resulting Na⁺ and Cl⁻, repeated in all directions through the crystal, is the ionic bond. This "electron transfer" picture is a useful model — no bond is 100% ionic, and real bonds carry some covalent character (next topic).
Coulomb's law and bond strength
The attraction between two ions is given by Coulomb's law:
E = kQ1 Q2d
where Q1 and Q2 are the ion charges, d is the distance between centers, and k is a constant. Charge enters directly (2+/2− attract about four times as strongly as 1+/1−) and distance matters because smaller ions sit closer. Since E is negative for attraction, larger magnitudes mean stronger bonds and higher lattice energies.
Lattice energy
lattice energy Energy released when gaseous ions form a solid Full entry → is the energy released when gaseous ions assemble into one mole of an ionic solid. It is the thermodynamic payoff that makes ionic compounds stable: though ionizing sodium costs energy, the assembled crystal's attraction more than repays it. Lattice energy follows Coulomb's law:
- Higher charges → larger lattice energy. MgO (2+/2−) has a lattice energy near 3795 kJ/mol — about five times NaCl's ≈ 787 kJ/mol.
- Smaller ions → larger lattice energy. LiF beats LiCl because F⁻ is smaller than Cl⁻, so the ions sit closer.
What lattice energy predicts
Melting an ionic solid means tearing ions apart, so melting point tracks lattice energy: MgO melts above 2800 °C, NaCl ≈ 801 °C, KCl lower. Hardness follows the same reasoning. Solubility does not — it depends on how well water surrounds each ion (hydration energy).
Properties of ionic solids
All signature properties trace to strong, non-directional ion–ion attractions:
- Crystalline structure. Ions pack into repeating 3-D lattices (NaCl's face-centered cubic arrangement).
- High melting and boiling points. Breaking the lattice takes enormous energy.
- Hard but brittle. A sharp blow shifts a layer so like charges align; repulsion shatters the crystal instead of bending it.
- Poor solid conductivity, good molten/dissolved conductivity. Ions locked in the solid cannot move; melted or dissolved, they carry current. Saltwater conducts; dry salt does not.
Lewis symbols for ions
A Lewis symbol Symbol plus dots for valence electrons Full entry → shows an element's symbol plus dots for valence electrons. Ions are written in brackets with the charge: Na+ has no dots (it lost its only valence electron), while Cl- shows eight — a completed octet. These symbols build the Lewis structures in Topic 4.
Worked Example: Applying Coulomb's Law to Compare Lattice Energies
Problem. Which has the larger lattice energy and higher melting point: NaCl or MgO?
Solution. Write Coulomb's law and compare charge products:
E = kQ1 Q2d
For NaCl, Q1 Q2 = (+1)(-1) = -1; for MgO, (+2)(-2) = -4, four times larger in magnitude, so attraction per unit distance is about four times stronger. Measured lattice energies confirm: NaCl ≈ 787 kJ/mol versus MgO ≈ 3795 kJ/mol; melting points follow (≈ 801 °C vs above 2800 °C). Conclusion: the 2+/2− pair makes MgO far more strongly bound.
Worked Example: The Size Effect on Lattice Energy
Problem. LiF and LiCl both pair a +1 cation with a −1 anion. Which has the larger lattice energy, and why?
Solution. Charges are identical, so distance decides:
E = kQ1 Q2d ⇒ smaller d gives larger |E|
Fluoride (radius ≈ 133 pm) is much smaller than chloride (≈ 181 pm), so Li⁺ and F⁻ sit closer and attract more strongly. LiF has the larger lattice energy (≈ 1049 vs ≈ 853 kJ/mol for LiCl) and the higher melting point (≈ 845 vs ≈ 605 °C). Rule: for equal charges, smaller ions give stronger bonds.
Worked Example: Predicting Properties from Evidence
Problem. An unknown white solid dissolves in water and the solution conducts. The solid does not conduct when dry and melts above 700 °C. What class is it, and why?
Solution. The evidence points to an ionic solid. In the solid, ions are locked in a fixed lattice and cannot move, so no current flows; once dissolved, the lattice breaks apart and mobile ions carry charge. A molecular solid like sugar would dissolve without forming ions and would not conduct.
Common Confusions
| Do Not Confuse | With | Difference |
|---|---|---|
| Ionic solids | Conductors of electricity | Solid salts are insulators; conduction needs molten or dissolved ions |
| Ionic bond | A molecule | Ionic compounds are extended lattices of formula units, not discrete molecular pairs |
| Lattice energy | Ionization energy | Lattice energy assembles a crystal from ions; ionization energy removes an electron from an atom |
| "NaCl is 100% ionic" | Partial covalent character | Real bonds are a spectrum; the ionic model is an approximation |
| Brittleness | Weakness | Ionic crystals are hard; brittleness comes from like-charge repulsion under stress |
| Higher lattice energy | Higher solubility | Solubility depends on hydration energy, not lattice energy alone |
| Electron transfer | Electron sharing | Opposite ends of one bonding continuum |
Eli explains
The same idea, in plain words
Explain it like I’m 10
Imagine two friends on a seesaw: one atom hands an electron to another, and now one is a plus magnet and the other a minus magnet. They stick together tightly, and many pairs stack into a 3-D brick wall that is hard and melts only at very high temperatures — but hit it just wrong and it cracks, because like charges end up pushing each other apart. That wall is a salt crystal.
Key takeaways
- An ionic bond = electrostatic attraction between oppositely charged ions; the transfer model applies to metal + nonmetal.
- Coulomb's law: E = kQ1Q2/d — attraction scales with charge product, inversely with distance.
- Lattice energy is released when gaseous ions form a solid; it is why ionic compounds exist.
- Trends: higher charges and smaller ions → larger lattice energy → higher melting point (MgO > NaCl > KCl).
- Ionic solids: crystalline, hard, brittle, high melting, insulators as solids, conductors when molten or dissolved.
- No discrete molecules: formula units in a lattice.
- Real bonds are a continuum; even "ionic" NaCl has some covalent character.
Check yourself
5 review questions from the chapter. Try each one, then open the answer.
State Coulomb's law and name the two factors controlling ionic bond strength.
Show answer
E = kQ1Q2/d: attraction strengthens with larger ion charges and smaller interionic distance.
Why does MgO melt far above the melting point of NaCl?
Show answer
MgO pairs 2+/2− ions (charge product 4× NaCl's), giving much larger lattice energy and thus a much higher melting point.
True or false: solid table salt conducts electricity. Explain.
Show answer
False as stated — solid NaCl is an insulator because ions are locked in the lattice; it conducts when molten or dissolved, when ions move freely.
Arrange LiF, KCl, and MgO in order of increasing lattice energy.
Show answer
KCl < LiF < MgO (size: Cl⁻ > F⁻ → KCl < LiF; charge makes MgO largest).
What is lattice energy, and why is it essential for the stability of ionic compounds?
Show answer
Lattice energy is the energy released when gaseous ions form a solid; it repays the energy cost of forming ions, which is why ionic compounds form at all.
Study tools & related lessonsKey vocabulary · Related
Key vocabulary
- ionic bond
- Electrostatic attraction between oppositely charged ions
- cation
- Positive ion formed when a metal loses electrons
- anion
- Negative ion formed when a nonmetal gains electrons
- Coulomb's law
- E = kQ1Q2/d
- lattice energy
- Energy released when gaseous ions form a solid
- crystal lattice
- Repeating 3-D ion arrangement
- electrolyte
- Substance that conducts when molten or dissolved
- ionization energy
- Energy to remove an electron from a gaseous atom
- electron affinity
- Energy change when a gaseous atom gains an electron
- Lewis symbol
- Symbol plus dots for valence electrons
Sources & references
This lesson was adapted from the open educational references above; their licenses and attributions are preserved. See Copyright & Licensing.
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