Ionic Bonding and Structure for the ESAT

How are ionic compounds formed?

When a metal reacts with a non-metal, there is usually the transfer of electrons. During this process, the metal atoms lose electrons to become positively charged ions (cations), while the non-metal atoms gain these electrons to become negatively charged ions (anions). The resulting ionic compound is not a collection of individual molecules but a giant lattice composed of positive and negative ions held together by strong electrostatic forces.

Predicting ion charges

When atoms react to form ions, they gain or lose electrons to obtain the electron configuration of a noble gas, which represents a stable state. Consequently, the charge of these ions depends on the original electron configuration of the atoms. For instance, atoms in Group 1 of the Periodic Table possess one electron in their outer shell. By losing this single electron, they form $1+$ ions.

The following trends apply to common elements in the Periodic Table:

1. Group 1: Form $1+$ ions (e.g. $Li^+$, $Na^+$, $K^+$).
2. Group 2: Form $2+$ ions (e.g. $Mg^{2+}$, $Ca^{2+}$, $Ba^{2+}$).
3. Group 13: Aluminium forms $Al^{3+}$.
4. Group 16: Form $2-$ ions (e.g. $O^{2-}$, $S^{2-}$).
5. Group 17: Form $1-$ ions (e.g. $F^-$, $Cl^-$, $Br^-$, $I^-$).

Some metals can form ions with different charges depending on the reaction conditions. In these cases, a Roman numeral is included in the name of the compound to identify the specific charge on the metal ion. For example, iron(II) oxide contains $Fe^{2+}$ ions, whereas iron(III) oxide contains $Fe^{3+}$ ions.

Common compound ions

Combinations of different atoms can exist as a single unit with an overall charge, known as 'compound ions' or polyatomic ions. You must learn the formulae and charges for the following common ions:

• Hydrogen: $H^+$
• Ammonium: $NH_4^+$
• Hydroxide: $OH^-$
• Nitrate: $NO_3^-$
• Sulfate: $SO_4^{2-}$
• Carbonate: $CO_3^{2-}$
• Phosphate: $PO_4^{3-}$

Determining the formulae of ionic compounds

In any ionic compound, the total number of positive and negative charges must balance out to zero. This principle allows us to determine the chemical formula by ensuring the total charge of the cations equals the total charge of the anions.

Worked Examples

• Iron(III) bromide: The iron(III) ion is $Fe^{3+}$ and the bromide ion is $Br^-$. To balance the $3+$ charge, three $Br^-$ ions are needed. Formula: $FeBr_3$.
• Calcium oxide: The calcium ion is $Ca^{2+}$ and the oxide ion is $O^{2-}$. One of each ion balances the charge. Formula: $CaO$.
• Potassium sulfide: The potassium ion is $K^+$ and the sulfide ion is $S^{2-}$. Two $K^+$ ions are needed to balance one $S^{2-}$ ion. Formula: $K_2S$.
• Aluminium oxide: The aluminium ion is $Al^{3+}$ and the oxide ion is $O^{2-}$. To balance the charges, we need two $Al^{3+}$ ions (total $6+$) and three $O^{2-}$ ions (total $6-$). Formula: $Al_2O_3$.
• Sodium carbonate: The sodium ion is $Na^+$ and the carbonate ion is $CO_3^{2-}$. Two $Na^+$ ions are required. Formula: $Na_2CO_3$.
• Magnesium hydroxide: The magnesium ion is $Mg^{2+}$ and the hydroxide ion is $OH^-$. Two $OH^-$ ions are required. Brackets are used to show the multiplier applies to the whole compound ion. Formula: $Mg(OH)_2$.
• Ammonium sulfate: The ammonium ion is $NH_4^+$ and the sulfate ion is $SO_4^{2-}$. Two $NH_4^+$ ions are required. Formula: $(NH_4)_2SO_4$.

Physical properties of ionic compounds

Ionic compounds are generally solids at room temperature. They exist as a giant lattice, which is a huge, continuous, and regular structure of repeating ions.

There is a strong overall electrostatic attraction between all the positive and negative ions throughout the lattice. This leads to several key properties:

1. High melting points: Significant energy is required to overcome the many strong attractions between the ions in the lattice.
2. Electrical conductivity: In the solid state, ionic compounds do not conduct electricity because the ions are held in fixed positions and cannot move. However, when melted (molten) or dissolved in water (aqueous solution), the lattice breaks down and the ions are free to move and carry the charge, making the substance a conductor.