Chemistry notes, chapter 5
Chemical Bond
Textbook Chemistry Form 4 KSSM (DLP), pages 108–133
What this chapter covers
- Explain why atoms form chemical bonds, using the stable duplet and octet electron arrangements (5.1)
- Describe the formation of ions and ionic bonds, and write half-equations (5.2)
- Describe single, double and triple covalent bonds and compare ionic and covalent bonds (5.3)
- Explain hydrogen bonds and their effect on boiling point and solubility (5.4)
- Explain the formation of a dative bond, for example in H3O+ and NH4+ (5.5)
- Explain the metallic bond and why metals conduct electricity (5.6)
- Compare the properties and uses of ionic and covalent compounds (5.7)
5.1Basics of compound formation
Chemical bond: a force that holds atoms or ions together when a compound is formed. Chemical bonds involve only the valence electrons.
Noble gases exist as single atoms (monoatomic) and are unreactive because they already have a stable duplet (helium) or octet electron arrangement. Atoms of other elements become stable by transferring or sharing electrons.
| Bond | How electrons are used | Between |
|---|---|---|
| Ionic bond | Transfer of electrons | Metal atom and non-metal atom |
| Covalent bond | Sharing of electrons | Non-metal atoms |
Textbook Figures 5.1 and 5.2, pp. 110-111
Why does no compound form when sodium is heated with neon?
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Neon already has a stable octet, so its atoms do not gain, lose or share electrons.
5.2Ionic bond
Ionic bond: the strong electrostatic attraction force between positive and negative ions, formed after a metal atom transfers electrons to a non-metal atom.
- A metal atom donates valence electrons and becomes a positive ion (cation). Na → Na+ + e−
- A non-metal atom accepts electrons and becomes a negative ion (anion). F + e− → F−
- Oppositely charged ions attract one another and form an ionic compound.
Based on textbook Figure 5.6, p. 113
Example: magnesium chloride, MgCl2
Mg is 2.8.2 and donates 2 electrons: Mg → Mg2+ + 2e− (2.8).Each Cl is 2.8.7 and accepts 1 electron: Cl + e− → Cl− (2.8.8).
One Mg gives electrons to two Cl, so the formula is MgCl2.
Ionic charge: the number of electrons lost or gained. Mg loses 2, so Mg2+; Al loses 3, so Al3+; O gains 2, so O2−.
Write the half-equation for the formation of the aluminium ion (proton number 13).
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Al is 2.8.3, so it donates 3 electrons: Al → Al3+ + 3e−.
5.3Covalent bond
Covalent bond: formed when non-metal atoms share pairs of electrons to achieve a stable duplet or octet arrangement.
| Type | Pairs shared | Example | Why |
|---|---|---|---|
| Single bond | 1 | Chlorine, Cl2; HCl; H2O | Cl (2.8.7) needs 1 electron |
| Double bond | 2 | Oxygen, O2; CO2 | O (2.6) needs 2 electrons |
| Triple bond | 3 | Nitrogen, N2 | N (2.5) needs 3 electrons |
Based on textbook Figures 5.7 and 5.8, pp. 114-115
A Lewis structure shows only the valence electrons, and a line between two atoms stands for one shared pair (Cl–Cl, O=O, N≡N).
Ionic bond and covalent bond
| Ionic bond | Covalent bond | |
|---|---|---|
| Electrons | Transferred | Shared |
| Between | Metal and non-metal | Non-metal and non-metal |
| Particles formed | Positive and negative ions | Molecules |
| Similarity | Only valence electrons are involved and atoms achieve a stable duplet or octet. | |
Showing electron transfer for a compound of two non-metals, such as CO2 or HCl. Two non-metals share electrons.
Carbon (proton number 6) bonds with hydrogen. How many hydrogen atoms bond with one carbon atom?
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C is 2.4, so it needs 4 more electrons and shares with 4 hydrogen atoms: CH4.
5.4Hydrogen bond
Hydrogen bond: an attraction force between a hydrogen atom bonded to N, O or F in one molecule and an N, O or F atom in another molecule.
- Water molecules form hydrogen bonds with one another. Ice is less dense than water, so ice floats.
- Boiling point: besides weak Van der Waals forces, hydrogen bonds must be broken, so more heat is needed. Ethanol (C2H5OH) boils at a higher temperature than Cl2.
- Solubility: ethanol dissolves in water because it forms hydrogen bonds with water molecules.
- Daily life: wet hair sticks together and paper sticks to a wet finger because of hydrogen bonds with water.
Textbook Figures 5.12 to 5.16, pp. 117-119
Hydrogen bonds are not formed in HCl, because chlorine is not N, O or F.
Why does hydrogen fluoride have a much higher boiling point than hydrogen chloride?
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HF molecules form hydrogen bonds between them. Extra heat is needed to overcome these strong forces, so the boiling point is high.
5.5Dative bond
Dative (coordinate) bond: a covalent bond in which the shared pair of electrons comes from one atom only.
In H3O+, a water molecule uses a lone pair on oxygen to bond with a hydrogen ion, H+, which has no electrons. The same happens in NH4+, where nitrogen in NH3 supplies the pair. After bonding, oxygen has an octet and each hydrogen has a duplet.
Textbook Figure 5.18, p. 120
Look for a lone pair on one atom and an empty shell on the other. A boron atom in BF3 has only six valence electrons, so it can accept a pair from NH3.
Which atom supplies the electron pair in the bond between NH3 and H+?
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The nitrogen atom, using its lone pair. H+ has no electrons to give.
5.6Metallic bond
Metallic bond: the electrostatic attraction between positive metal ions and the sea of delocalised electrons.
- Metal atoms are packed closely in an orderly way. Their valence electrons are donated and become delocalised (free to move, owned by no particular atom).
- When electricity is supplied, the delocalised electrons move freely and carry charge from the negative terminal to the positive terminal, so metals conduct electricity.
Textbook Figures 5.19 and 5.20, p. 121
Explain how aluminium conducts electricity.
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Aluminium has a sea of delocalised electrons. They move freely through the metal and carry charge when a voltage is applied.
5.7Properties of ionic and covalent compounds
| Property | Ionic compounds | Covalent compounds (simple molecules) |
|---|---|---|
| Melting and boiling point | High: much heat needed to overcome strong electrostatic forces | Low: little heat needed to overcome weak Van der Waals forces |
| Conductivity | None as solid; conducts when molten or in aqueous solution (ions move freely) | None in any state (neutral molecules) |
| Solubility | Soluble in water; insoluble in organic solvents | Insoluble in water (mostly); soluble in organic solvents |
Textbook Experiment 5.1 and Figures 5.24 to 5.26, pp. 123-127
Water is a polar solvent: oxygen has partial negative charge and hydrogen partial positive charge. Water molecules attract the ions and break up the lattice, so ionic compounds dissolve.
Simple and giant molecules
| Simple molecule | Giant molecule | |
|---|---|---|
| Example | H2O, CO2, CH4 | SiO2, diamond |
| Forces overcome on melting | Weak Van der Waals forces between molecules | Strong covalent bonds throughout the structure |
| Melting and boiling point | Low | High |
Textbook Figure 5.27, p. 128
Uses
- Ionic: LiI in batteries, KCl and NH4NO3 in fertilisers, NaHCO3 in antacids, NaClO3 in detergents.
- Covalent: pigment and turpentine in paint, bromoethane in pesticides, paracetamol for fever, glycerol in skincare.
Saying that melting or boiling ice (a simple molecule) breaks covalent bonds. Only the weak forces between molecules are overcome; the covalent bonds inside each molecule stay.
A white solid dissolves in water and conducts electricity when molten. What type of compound is it?
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An ionic compound. Its ions are free to move in the molten state, and water molecules can pull the ions apart.