Chemical bonding is the process by which atoms combine to achieve a more stable outer electron configuration. When metals react with non-metals, they achieve this stability through the transfer of electrons, forming ions held together by strong electrostatic attractions.
🔑 Key Principle
An ionic bond is not a localized link between two specific atoms. Instead, it is the multidirectional, non-directional electrostatic attraction between oppositely charged ions. This attraction acts in all directions, binding the ions into a giant three-dimensional lattice.
What is an Ionic Bond?
During the formation of an ionic bond, metal atoms lose electrons to form positively charged ions (cations), while non-metal atoms gain these electrons to form negatively charged ions (anions). The transfer of electrons allows both atoms to achieve a stable, full outer shell of electrons (the noble gas configuration).
The electrostatic attraction between oppositely charged ions, acting in all directions throughout a giant ionic lattice.
Formulas of Common Ions
You must be able to deduce the formulas of ionic compounds from the charges of the constituent ions. Memorize the formulas and charges of these common polyatomic ions:
| Cations (+) | Formula | Anions (−) | Formula |
|---|---|---|---|
| Sodium | \( \text{Na}^+ \) | Chloride | \( \text{Cl}^- \) |
| Magnesium | \( \text{Mg}^{2+} \) | Oxide | \( \text{O}^{2-} \) |
| Ammonium | \( \text{NH}_4^+ \) | Hydroxide | \( \text{OH}^- \) |
| Aluminum | \( \text{Al}^{3+} \) | Nitrate | \( \text{NO}_3^- \) |
| Iron(II) | \( \text{Fe}^{2+} \) | Carbonate | \( \text{CO}_3^{2-} \) |
| Iron(III) | \( \text{Fe}^{3+} \) | Sulfate | \( \text{SO}_4^{2-} \) |
Dot-and-Cross Diagrams
Dot-and-cross diagrams are used to show the arrangement of electrons in the outer shell of atoms and ions. Dots represent electrons originating from one atom, while crosses represent electrons from the other.
For A-Level Chemistry, you must be able to draw dot-and-cross diagrams for simple ionic compounds like \( \text{NaCl} \), \( \text{MgO} \), and \( \text{MgCl}_2 \).
When drawing dot-and-cross diagrams for ionic compounds:
- Always put square brackets around each ion.
- Write the charge outside the top right corner of the brackets (e.g. \( 2+ \) or \( 2- \)).
- Draw the outer electrons only. Remember to use a different symbol (e.g. a cross) for the transferred electron(s) in the anion to show they came from the cation.
Solution:
1. Magnesium is in Group 2, so it has 2 outer shell electrons: \( \text{Mg} \) [2, 8, 2].
2. Chlorine is in Group 7, so it has 7 outer shell electrons: \( \text{Cl} \) [2, 8, 7].
3. To achieve a full outer shell, the magnesium atom must lose 2 electrons. Each chlorine atom needs to gain only 1 electron.
4. Therefore, one magnesium atom transfers 1 electron to each of two chlorine atoms.
5. The ions formed are a magnesium cation, \( \text{Mg}^{2+} \) [2, 8], and two chloride anions, \( \text{Cl}^- \) [2, 8, 8]. The compound formula is \( \text{MgCl}_2 \).
The Giant Ionic Lattice
In the solid state, ionic compounds do not exist as isolated pairs. They exist as a regular, repeating three-dimensional arrangement of alternating positive and negative ions, known as a giant ionic lattice.
In a sodium chloride lattice, each sodium ion is surrounded by six chloride ions, and each chloride ion is surrounded by six sodium ions. This is described as a 6,6-coordination.
Physical Properties of Ionic Compounds
The physical properties of ionic substances are a direct result of their giant lattice structure and the strong electrostatic forces holding the ions together.
| Property | Observation | Explanation in terms of Structure and Bonding |
|---|---|---|
| Melting & Boiling Points | High | Strong electrostatic attractions act throughout the giant lattice in all directions. A large amount of thermal energy is required to overcome these strong bonds. |
| Electrical Conductivity | Conducts when molten or dissolved; insulator when solid | In the solid state, ions are fixed in position within the lattice and cannot move to carry charge. When molten or dissolved in water, the lattice is broken, freeing the ions to move and carry charge. |
| Brittleness | Brittle | Applying a force shifts layers of ions. This brings ions of like charge next to each other, resulting in repulsion that shatters the crystal. |
| Solubility | Often soluble in polar solvents (e.g., water) | Polar water molecules attract the ions on the lattice surface, pulling them away from the lattice and stabilizing them in solution (hydration). |
Solution:
1. Identify the constituent ions in each compound:
- Sodium chloride consists of \( \text{Na}^+ \) and \( \text{Cl}^- \) ions.
- Magnesium oxide consists of \( \text{Mg}^{2+} \) and \( \text{O}^{2-} \) ions.
2. Compare ionic charges: The charges on the ions in magnesium oxide (\( 2+ \) and \( 2- \)) are higher than the charges on the ions in sodium chloride (\( 1+ \) and \( 1- \)).
3. Compare ionic radii: Magnesium ions (\( \text{Mg}^{2+} \)) are smaller than sodium ions (\( \text{Na}^+ \)) because they have more protons pulling the same number of shells. Oxide ions (\( \text{O}^{2-} \)) are also smaller than chloride ions (\( \text{Cl}^- \)).
4. Explain electrostatic attraction: The higher charges and smaller radii of \( \text{Mg}^{2+} \) and \( \text{O}^{2-} \) result in much stronger electrostatic attractions between oppositely charged ions in the lattice of \( \text{MgO} \) compared to \( \text{NaCl} \).
5. Relate to melting point: Consequently, significantly more thermal energy is required to overcome the electrostatic attractions and break the lattice of \( \text{MgO} \), resulting in a much higher melting point.
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