Chapter map: this chapter goes in seven steps, and you’re at step 1. The first three steps are about what holds atoms together; the last four are about what covalent molecules look like and how their structure decides their properties.
- Ionic bonds 2.4.1
- Covalent bonds 2.4.2
- Metallic bonds 2.4.3
- Drawing structures 2.4.4–2.4.5
- Molecular shape 2.4.6
- Polarity 2.4.7
- Intermolecular forces 2.4.8
Why is table salt so tough?
Table salt doesn’t melt until 801 °C. Its crystals are hard but brittle, shattering when struck. Solid salt doesn’t conduct electricity, but melted or dissolved in water, it does. Lesson 1.3.6 showed that sodium chloride is made of Na⁺ and Cl⁻ ions. This lesson looks at what holds those ions together, and how strongly.
Ionic bonds
An ionic bond离子键 is the attraction between oppositely charged ions. It is mainly the electrostatic attraction between cations and anions; if the ions get too close, the repulsion between their nuclei and between their electrons grows, and the ions settle where attraction and repulsion balance.
- Which elements form ionic bonds: usually a reactive metal (groups 1 and 2) and a reactive nonmetal (groups 16 and 17). In the language of lesson 2.3.5, when the electronegativity difference is large (usually more than 1.7), electrons are essentially transferred and an ionic bond forms.
- A compound held together by ionic bonds is an ionic compound离子化合物. Besides metal oxides and halides such as NaCl and MgO, strong bases (NaOH, KOH) and most salts, including ammonium salts such as NH₄Cl that contain no metal, are ionic compounds.
Showing how an ionic compound forms
You can show the electron transfer with electron-dot symbols. On the left, draw the atoms’ Lewis symbols, with curved arrows showing which electrons move where; on the right, draw the ions that result.
- A cation’s symbol has no dots left, just the charge: Na⁺, Mg²⁺. An anion has a full octet of 8 dots, drawn in square brackets with the charge at the upper right.
- Each ion is drawn separately: MgCl₂ is shown as one Mg²⁺ with a bracketed Cl⁻ on each side.
- To make the transfer easy to follow, the metal’s electrons are often drawn as × and the nonmetal’s as dots. Chinese textbooks require this “formation diagram”; American textbooks use the same idea with Lewis symbols.
Choose a metal and a nonmetal and watch where the electrons go and how many atoms of each you need. Then switch to the second tab to compare lattice energies:
The metal's electrons are drawn as × and the nonmetal's own valence electrons as dots, so you can follow each one; the curved arrows show where each electron goes. This dot-and-cross style is standard in Chinese and British textbooks.
Lattice energy is the energy needed to pull 1 mol of an ionic solid completely apart into gaseous ions far from one another. The larger the lattice energy, the stronger the ionic bonding, and usually the higher the melting point.
Data from the CRC Handbook. Some textbooks instead define lattice energy as the energy released when gaseous ions form 1 mol of the solid, a negative number such as −786 kJ/mol for NaCl.
Structure and properties of ionic compounds
In an ionic compound, the cations and anions are arranged in a regular three-dimensional pattern, an ionic crystal离子晶体 (or crystal lattice). In sodium chloride, each Na⁺ is surrounded by 6 Cl⁻ ions and each Cl⁻ by 6 Na⁺ ions. There are no separate “NaCl molecules”; the formula NaCl, called a formula unit, only gives the 1 : 1 ratio of the ions.
The properties of ionic compounds all follow from this structure:
- High melting and boiling points: melting means overcoming ionic bonds that run through the whole crystal, which takes a lot of energy.
- Hard but brittle: a sharp blow shifts one layer of ions so that like charges face each other, repel, and split the crystal (lesson 2.4.3 animates this).
- Conduct only when melted or dissolved: in the solid, the ions are locked in place; melted or dissolved, they can move.
How strong is an ionic bond? Coulomb’s law
You met Coulomb’s law in lesson 2.3.2. The energy of attraction between two ions is roughly:
Here q₁ and q₂ are the ion charges and r is the distance between them (about the sum of the two ionic radii). So:
- The larger the charges, the stronger the ionic bonding. Both ions in MgO carry a charge of 2, so the product of the charges is 4 times that in NaF.
- The smaller the ions, the closer they get and the stronger the bonding. From NaF to NaI, the anion gets larger and the ionic bonding gets weaker.
One measure of ionic-bond strength is the lattice energy晶格能: the energy needed to pull 1 mol of an ionic solid (for NaCl, 6.022 × 10²³ Na⁺ ions and as many Cl⁻ ions; see lesson 2.2.2 for the mole) completely apart into gaseous ions far from one another. The larger the lattice energy, the stronger the bonding and, usually, the higher the melting point. NaCl’s lattice energy is 786 kJ/mol; MgO’s is 3791 kJ/mol. (Some books write lattice energy as the energy released when the solid forms, with a negative sign.)
Example 1: How does MgCl₂ form?
Use electron-dot symbols to show how magnesium chloride forms.
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Draw the atoms. Magnesium has 2 valence electrons, one dot on each of two sides; chlorine has 7, one short of an octet.
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How many atoms: electrons lost must equal electrons gained. Each magnesium atom loses 2 electrons, but each chlorine atom can take only 1, so one magnesium needs two chlorines:
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Draw the arrows. Draw one curved arrow from each of magnesium’s two electrons to a different chlorine atom, showing where each electron goes.
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Write the product. Magnesium becomes Mg²⁺, and each chlorine becomes a bracketed Cl⁻ with a full octet, drawn on either side of the Mg²⁺. Choose Mg and Cl in the tool above to compare.
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A common mistake. Writing the two chloride ions together as “2Cl⁻” on one side, or leaving off the brackets and charges. Draw each ion separately.
Example 2: Ranking melting points
Without looking up data, rank NaCl, MgO, and KCl from highest to lowest melting point.
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The basis. Melting an ionic solid overcomes ionic bonding, whose strength follows Coulomb’s law: higher charges and smaller distances mean stronger attraction and a higher melting point.
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Compare charges first. The charge products:
MgO’s bonding is by far the strongest.
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Same charges? Compare sizes. NaCl and KCl share the same anion, but K⁺ (138 pm) is larger than Na⁺ (102 pm), so K⁺ and Cl⁻ sit farther apart and attract less strongly.
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Conclusion. Melting point: MgO > NaCl > KCl.
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Check with data. MgO melts above 2800 °C, NaCl at 801 °C, and KCl at 771 °C. Compare charges first; they matter much more than size.
- NaCl
Sodium is a reactive metal and chlorine a reactive nonmetal; the electronegativity difference is 2.23.
- HCl
Hydrogen and chlorine are both nonmetals and share electrons: HCl is a covalent compound (lesson 2.4.2).
- MgO
It is made of Mg²⁺ and O²⁻ ions, with very strong ionic bonding.
- CO₂
Carbon and oxygen are both nonmetals; CO₂ is made of molecules with covalent bonds only.
- KOH
K⁺ and OH⁻ are held by ionic bonds; inside the OH⁻ ion, the O–H bond is covalent.
- NH₄Cl
It contains no metal, but it is made of NH₄⁺ and Cl⁻ ions, so it is ionic.
- H₂O
Water is made of molecules with covalent O–H bonds only.
- CaF₂
It is made of Ca²⁺ and F⁻ ions: an ionic compound.
- AlCl₃
The electronegativity difference is only 3.16 − 1.61 = 1.55, and AlCl₃ is a covalent compound, a classic exception.
Common mistakes
Mistake 1: “Any compound containing a metal is ionic.” AlCl₃ contains aluminum but is covalent. In the other direction, NH₄Cl contains no metal but is ionic. What matters is whether there are ionic bonds, not whether there is a metal.
Mistake 2: “An ionic compound contains only ionic bonds.” In NaOH, Na⁺ and OH⁻ are held by an ionic bond, but the O and H inside OH⁻ share a covalent bond. Any compound with ionic bonds is ionic, even if it also has covalent bonds.
Mistake 3: “The formula NaCl stands for one sodium chloride molecule.” An ionic crystal has no molecules; every ion is surrounded by several ions of opposite charge. NaCl is a formula unit that gives the 1 : 1 ratio of Na⁺ to Cl⁻.
Quick check
Which of these substances contains ionic bonds?
CaCl₂ is made of Ca²⁺ and Cl⁻ ions. HCl, CO₂, and H₂O are covalent compounds made of molecules.
Which ionic compound has the highest melting point?
Both ions in MgO carry a charge of 2, and they are about the same size as the ions in NaF, so MgO has the strongest ionic bonding and the highest melting point (above 2800 °C).
Summary
- An ionic bond is the attraction between oppositely charged ions, usually formed between a reactive metal and a reactive nonmetal; a compound with ionic bonds is an ionic compound.
- Electron-dot formation diagrams use curved arrows for the transferred electrons; anions go in brackets with their charge, and each ion is drawn separately.
- Ionic crystals contain no molecules. Ionic compounds have high melting points, are hard but brittle, and conduct only when melted or dissolved. The larger the ion charges and the smaller the ions, the stronger the ionic bonding and the larger the lattice energy.
On to AP (for later)
AP U2.1 and U2.3 both use Coulomb’s law to compare ionic compounds. A classic question asks why MgO melts at a higher temperature than NaF; a full-credit answer says the ions have larger charges and similar sizes, so the Coulombic attraction is stronger. AP also uses particle diagrams of ionic solids, with cations and anions alternating and the cations usually drawn smaller.