Chemical Bonding
Ionic Bonding
5th Year · 6th Year (Leaving Cert)
- ✓By the end of this lesson students will be able to define ionic bonding and explain how it forms.
- ✓By the end of this lesson students will be able to describe the formation of cations and anions, including their electron configurations.
- ✓By the end of this lesson students will be able to explain the lattice structure of ionic compounds.
- ✓By the end of this lesson students will be able to relate the characteristic physical and chemical properties of ionic compounds to their bonding and structure.
- ✓By the end of this lesson students will be able to write balanced chemical formulae for simple ionic compounds.
Key concepts
Ionic bonding is a type of chemical bond formed by the electrostatic attraction between oppositely charged ions. It typically occurs between a metal atom and a non-metal atom. The metal atom transfers one or more valence electrons to the non-metal atom, resulting in the formation of positively charged cations and negatively charged anions. These oppositely charged ions are then held together by strong electrostatic forces of attraction.
Atoms form ions to achieve a stable electron configuration, typically a full outer electron shell (an octet, or a duet for hydrogen and helium). Metals tend to lose their valence electrons to form positive ions called cations. For example, Group 1 metals lose one electron to form a +1 ion, and Group 2 metals lose two electrons to form a +2 ion. Non-metals tend to gain electrons to fill their valence shell, forming negative ions called anions. For example, Group 17 (halogens) gain one electron to form a -1 ion, and Group 16 non-metals gain two electrons to form a -2 ion.
When an atom forms an ion, its electron configuration changes to reflect the loss or gain of electrons. For cations, electrons are removed from the highest energy level. For anions, electrons are added to the highest energy level until the outer shell is full. For example: Sodium atom (Na): 1s²2s²2p⁶3s¹ Sodium ion (Na⁺): 1s²2s²2p⁶ (loses 1 electron) Chlorine atom (Cl): 1s²2s²2p⁶3s²3p⁵ Chloride ion (Cl⁻): 1s²2s²2p⁶3s²3p⁶ (gains 1 electron)
Ionic compounds do not exist as discrete molecules but rather as an extended, repeating three-dimensional arrangement of cations and anions. This regular, ordered arrangement is called an ionic crystal lattice. Each ion in the lattice is surrounded by a specific number of oppositely charged ions, maximising the attractive forces and minimising the repulsive forces. The strong electrostatic forces of attraction extend throughout the entire lattice, holding the ions rigidly in fixed positions.
The strong electrostatic forces within the ionic lattice give rise to characteristic properties: 1. **High Melting and Boiling Points**: A large amount of energy is required to overcome the strong electrostatic forces of attraction between ions in the lattice to allow them to move freely. 2. **Hard and Brittle**: Ionic compounds are hard due to the strong forces holding the ions in place. However, they are brittle because a strong force can cause a layer of ions to shift, bringing like-charged ions into proximity. The resulting repulsion causes the crystal to cleave or shatter. 3. **Electrical Conductivity**: In the solid state, ionic compounds do not conduct electricity because the ions are fixed in the lattice and cannot move to carry charge. However, when molten (liquid) or dissolved in a polar solvent (like water), the ions become mobile and are free to move, allowing them to conduct electricity. 4. **Solubility**: Many ionic compounds are soluble in polar solvents (e.g., water) because the polar water molecules can surround and separate the individual ions from the lattice (a process called hydration), overcoming the electrostatic forces. They are generally insoluble in non-polar solvents.
Key facts to remember
- 1Ionic bonds form between metals and non-metals through the transfer of electrons.
- 2Metal atoms lose electrons to form positively charged cations.
- 3Non-metal atoms gain electrons to form negatively charged anions.
- 4Ions achieve a stable electron configuration, typically a full outer shell (octet).
- 5Ionic compounds exist as a regular, repeating three-dimensional ionic crystal lattice.
- 6Strong electrostatic forces of attraction hold the ions together in the lattice.
- 7Ionic compounds generally have high melting and boiling points due to these strong forces.
- 8Solid ionic compounds do not conduct electricity because their ions are fixed in the lattice; however, they conduct when molten or in solution because ions are mobile.
Worked examples
Example 1
Describe the formation of sodium chloride (NaCl) from its constituent atoms, including their electron configurations and the type of bonding involved.
Answer
Sodium (Na) is a metal with electron configuration 1s²2s²2p⁶3s¹, and Chlorine (Cl) is a non-metal with electron configuration 1s²2s²2p⁶3s²3p⁵. Sodium loses its single valence electron to become a stable Na⁺ ion (1s²2s²2p⁶). This electron is transferred to chlorine, which gains it to become a stable Cl⁻ ion (1s²2s²2p⁶3s²3p⁶). The strong electrostatic attraction between the oppositely charged Na⁺ and Cl⁻ ions forms an ionic bond, resulting in the ionic compound sodium chloride (NaCl).
Remember to show the electron transfer and the resulting charges on the ions.
Example 2
Explain why solid magnesium oxide (MgO) has a very high melting point, typically over 2800 °C, and does not conduct electricity, whereas molten magnesium oxide does.
Answer
Magnesium oxide is an ionic compound with strong electrostatic forces of attraction between Mg²⁺ and O²⁻ ions in its crystal lattice. A significant amount of energy is needed to overcome these forces, hence its very high melting point. In the solid state, the ions are fixed and immobile, so they cannot conduct electricity. However, when molten, the ions become free to move and act as mobile charge carriers, allowing molten magnesium oxide to conduct electricity.
Always link properties back to the nature of the bonding and the mobility (or lack thereof) of the ions.
Example 3
Write the chemical formula for the ionic compound formed between calcium and fluorine, and explain the electron transfer involved.
Answer
The chemical formula is CaF₂. Calcium (Ca) loses two electrons to form a Ca²⁺ ion. Each fluorine (F) atom gains one electron to form an F⁻ ion. Therefore, one calcium atom transfers one electron to each of two fluorine atoms, resulting in one Ca²⁺ ion and two F⁻ ions, which combine to form the neutral ionic compound CaF₂.
Ensure the overall charge of the compound is neutral. Use the 'criss-cross' method if helpful for balancing charges.
Common mistakes
- ✗Confusing ionic bonding (electron transfer) with covalent bonding (electron sharing).
- ✗Incorrectly determining the charge of an ion (e.g., Ca⁺ instead of Ca²⁺).
- ✗Assuming that all ionic compounds are soluble in water.
- ✗Believing that solid ionic compounds conduct electricity.
- ✗Not understanding why ionic compounds are brittle, despite being hard (due to repulsion upon shifting layers).
- ✗Failing to balance the charges correctly when writing chemical formulae for ionic compounds.
Exam tips
- ★When describing ionic bonding, always use the term 'transfer of electrons' and mention the formation of 'oppositely charged ions' held by 'electrostatic forces'.
- ★For properties questions, always link the property directly to the 'strong electrostatic forces' in the lattice and the 'mobility of ions'.
- ★Practise writing electron configurations for both atoms and their corresponding ions.
- ★Ensure you can correctly determine the charges of common ions based on their group number in the Periodic Table and use them to write balanced chemical formulae.
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