Chemical Bonding, Structure and Properties for the ESAT

Formation of Ionic Compounds

When a metal reacts with a non-metal, electrons are transferred from the metal atoms to the non-metal atoms. This process forms ions: the metal loses electrons to become a positively charged cation, while the non-metal gains electrons to become a negatively charged anion. An ionic compound is the resulting giant lattice of these oppositely charged ions, held together by strong electrostatic attractions.

Ion Charges and Electron Configuration

Atoms typically form ions to obtain the stable electron configuration of a noble gas. The charge of an ion can be predicted from its Group in the Periodic Table:

1. Group 1: Elements have one outer electron and lose it to form $1+$ ions (e.g., $Li^+$, $Na^+$, $K^+$).
2. Group 2: Elements have two outer electrons and lose them to form $2+$ ions (e.g., $Mg^{2+}$, $Ca^{2+}$, $Ba^{2+}$).
3. Group 13 (Aluminium): Aluminium loses three electrons to form $Al^{3+}$.
4. Group 16: Elements have six outer electrons and gain two to form $2-$ ions (e.g., $O^{2-}$, $S^{2-}$).
5. Group 17: Elements have seven outer electrons and gain one to form $1-$ ions (e.g., $F^-$, $Cl^-$, $Br^-$, $I^-$).\n\nSome metals can exist in multiple oxidation states. In these cases, Roman numerals are used in the name to indicate the charge. For example, iron(II) oxide contains $Fe^{2+}$ ions, while iron(III) oxide contains $Fe^{3+}$ ions.

Compound Ions and Chemical Formulae

Groups of atoms can carry an overall charge, known as compound ions. You must memorise the following:

• Hydrogen: $H^+$
• Ammonium: $NH_4^+$
• Hydroxide: $OH^-$
• Nitrate: $NO_3^-$
• Sulfate: $SO_4^{2-}$
• Carbonate: $CO_3^{2-}$
• Phosphate: $PO_4^{3-}$\n\nTo determine the formula of an ionic compound, the total positive and negative charges must balance to zero. For example, in aluminium oxide, two $Al^{3+}$ ions (total $6+$) balance three $O^{2-}$ ions (total $6-$), giving the formula $Al_2O_3$. In magnesium hydroxide, one $Mg^{2+}$ balances two $OH^-$ ions, requiring brackets in the formula: $Mg(OH)_2$.

Properties of Ionic Compounds

Ionic compounds exist as a giant lattice, a regular, continuous structure of alternating positive and negative ions.

Their physical properties are a direct result of this structure:

1. Melting Points: They have high melting points because the strong electrostatic attractions between ions throughout the lattice require significant energy to overcome.
2. Conductivity: As solids, they are insulators because the ions are fixed in position. When melted (molten) or dissolved in water (aqueous), the ions are mobile and can carry an electrical charge.

Covalent Bonding and Molecular Substances

Covalent bonds form when non-metal atoms share one or more pairs of electrons. Most covalent substances exist as small molecules, such as methane ($CH_4$) or glucose ($C_6H_{12}O_6$). Within these molecules, atoms are held by strong covalent bonds, but the individual molecules are only held to each other by weak intermolecular forces.

Because these intermolecular forces are easy to overcome, molecular substances typically have low melting and boiling points. They do not conduct electricity because they lack mobile ions or delocalised electrons. Methane can be represented by 2D stick diagrams ($H-C-H$), 3D diagrams showing spatial arrangement, or space-filling models showing relative atom sizes.

Giant Covalent Structures

Some covalent substances form a continuous lattice of atoms. Examples include diamond (carbon), graphite (carbon), and silicon dioxide ($SiO_2$).

• Diamond and Silicon Dioxide: These have extremely high melting points and are very hard because every atom is linked to others by strong covalent bonds in a rigid 3D network.
• Graphite: While it has a high melting point, it is soft because it consists of layers that can slide over each other. It conducts electricity because it contains delocalised electrons that can move along these layers.

Metallic Bonding

Solid metals consist of a giant structure of positive metal ions surrounded by a "sea" of delocalised electrons. These are outer shell electrons that are free to move throughout the entire structure.

Metals have high melting points due to the strong attraction between the positive ions and the delocalised electron cloud. They are excellent conductors because the delocalised electrons carry charge. Metals are also malleable: the layers of ions can slide over each other without breaking the metallic bond, as the electron sea simply adjusts to the new positions.

Intermolecular Forces and Phase Changes

It is vital to distinguish between covalent bonds and intermolecular forces. When water ($H_2O$) boils to form steam, the intermolecular forces between molecules are overcome, but the covalent bonds between the hydrogen and oxygen atoms within the molecule remain intact.

Worked Examples: Identifying Structure from Properties

Consider a substance with a melting point of $801^{\circ}C$ that is an insulator as a solid but conducts when liquid. This fits the profile of an ionic structure. Conversely, a substance like diamond, with a melting point of $3550^{\circ}C$ that is an insulator in both states, is a giant covalent structure. Ethane ($C_2H_6$), which has a very low boiling point, has such properties because the forces between its molecules are weak, even though the internal $C-H$ bonds are strong.