Formation of Ionic Compounds (VCE SSCE Chemistry): Revision Notes

Formation of Ionic Compounds

Introduction to ion formation

When metal atoms react with non-metal atoms, they form ionic compounds through a process of electron transfer. These reactions can be highly energetic - for example, the reaction between sodium and chlorine to produce sodium chloride releases substantial heat and energy. During these reactions, atoms reorganise their electrons to form stable ions that arrange themselves into three-dimensional lattice structures.

Understanding how individual ions form from atoms and how to determine the ratio of different ions in a compound is fundamental to working with ionic compounds in chemistry.

Forming ions

When metal and non-metal atoms react to form ionic compounds, two key processes occur:

• Metal atoms lose electrons to form positively charged ions called cations
• Non-metal atoms gain electrons to form negatively charged ions called anions

This electron transfer happens because metals and non-metals have different electronegativities. Metals typically have low electronegativity values, meaning they have a weak attraction for electrons. In contrast, non-metals are more electronegative, meaning they have a stronger attraction for electrons. During ionic compound formation, non-metal atoms take one or more electrons from the outermost shell of metal atoms.

The octet rule and ionic compounds

Noble gases (group 18 elements) already possess this stable configuration, which is why they rarely react. Another way to understand the octet rule is that atoms tend to gain or lose electrons until they achieve an electron configuration identical to that of the nearest noble gas on the periodic table.

For example, when sodium reacts with chlorine:

• Each sodium atom loses one electron
• Each chlorine atom gains one electron
• The sodium ion achieves the electron configuration $2,8$ (same as neon)
• The chloride ion achieves the electron configuration $2,8,8$ (same as argon)

When atoms lose electrons, they become more positively charged because the number of protons now exceeds the number of electrons. This positive charge is indicated with a superscript $+$ sign. If two electrons are lost, the charge is written as $2+$. Similarly, when atoms gain electrons, they become negatively charged, indicated with a superscript $-$ sign. If three electrons are gained, the charge is written as $3-$.

Electron transfer diagrams

An electron transfer diagram is a visual representation showing how electrons move from metal atoms to non-metal atoms during ionic compound formation. These diagrams use Bohr models to show the electron shells before and after the transfer.

Sodium and chlorine reaction

Lithium and oxygen reaction

Magnesium and oxygen reaction

Magnesium and chlorine reaction

Writing equations for ionic reactions

To ensure you understand ion formation, it can be helpful to write equations showing the electronic configurations of both reactants and products. Here's an example for the reaction between lithium and nitrogen:

Chemical formulas of ionic compounds

Ionic compounds contain oppositely charged ions arranged in three-dimensional lattices. Different ions can have different charges - for example, aluminium forms ions with a $3+$ charge, whilst oxygen forms ions with a $2-$ charge. Understanding how to use ion charges to write the overall formula for an ionic compound is an essential skill.

Writing formulas of simple ionic compounds

For sodium chloride:

• A sodium ion ($\text{Na}^+$) has a $1+$ charge
• A chloride ion ($\text{Cl}^-$) has a $1-$ charge
• Therefore, the ratio is $1:1$ and the formula is $\text{NaCl}$

For magnesium chloride:

• A magnesium ion ($\text{Mg}^{2+}$) has a $2+$ charge
• A chloride ion ($\text{Cl}^-$) has a $1-$ charge
• Therefore, two chloride ions are needed to balance the $2+$ charge
• The ratio is $1:2$ and the formula is $\text{MgCl}_2$

The following diagram illustrates how formulas are determined for other ionic compounds:

Common ions reference

Here are some frequently encountered cations:

Charge	1+	2+	3+	4+
Ions	caesium ($\text{Cs}^+$)	barium ($\text{Ba}^{2+}$)	aluminium ($\text{Al}^{3+}$)	lead(IV) ($\text{Pb}^{4+}$)
	copper(I) ($\text{Cu}^+$)	cadmium(II) ($\text{Cd}^{2+}$)	chromium(III) ($\text{Cr}^{3+}$)	tin(IV) ($\text{Sn}^{4+}$)
	lithium ($\text{Li}^+$)	calcium ($\text{Ca}^{2+}$)	gold(III) ($\text{Au}^{3+}$)
	potassium ($\text{K}^+$)	cobalt(II) ($\text{Co}^{2+}$)	iron(III) ($\text{Fe}^{3+}$)
	silver ($\text{Ag}^+$)	copper(II) ($\text{Cu}^{2+}$)
	sodium ($\text{Na}^+$)	iron(II) ($\text{Fe}^{2+}$)
		lead(II) ($\text{Pb}^{2+}$)
		magnesium ($\text{Mg}^{2+}$)
		manganese(II) ($\text{Mn}^{2+}$)
		mercury(II) ($\text{Hg}^{2+}$)
		nickel ($\text{Ni}^{2+}$)
		strontium ($\text{Sr}^{2+}$)
		tin(II) ($\text{Sn}^{2+}$)
		zinc ($\text{Zn}^{2+}$)

Common anions include:

Charge	1-	2-	3-
Ions	bromide ($\text{Br}^-$)	oxide ($\text{O}^{2-}$)	nitride ($\text{N}^{3-}$)
	chloride ($\text{Cl}^-$)	sulfide ($\text{S}^{2-}$)
	fluoride ($\text{F}^-$)
	iodide ($\text{I}^-$)

Rules for writing chemical formulas

Worked example: determining chemical formulas

Writing formulas of more complex ionic compounds

Polyatomic ions

Up to this point, we have dealt with simple ions containing only one atom of an element. However, many ions contain two or more atoms, which may be of different elements. These are called polyatomic ions.

Characteristics of polyatomic ions:

• If different elements are present, they are combined in a fixed ratio
• The group of atoms behaves as a single unit with a specific charge
• Subscripts within the ion formula indicate the number of each type of atom

Common polyatomic ions

Charge	1+	1-	2-	3-
Ions	ammonium ($\text{NH}_4^+$)	cyanide ($\text{CN}^-$)	carbonate ($\text{CO}_3^{2-}$)	phosphate ($\text{PO}_4^{3-}$)
		hydrogen phosphate ($\text{H}_2\text{PO}_4^-$)	chromate ($\text{CrO}_4^{2-}$)
		ethanoate ($\text{CH}_3\text{COO}^-$)	dichromate ($\text{Cr}_2\text{O}_7^{2-}$)
		hydrogen carbonate ($\text{HCO}_3^-$)	hydrogen phosphate ($\text{HPO}_4^{2-}$)
		hydrogen sulfide ($\text{HS}^-$)	oxalate ($\text{C}_2\text{O}_4^{2-}$)
		hydrogen sulfite ($\text{HSO}_3^-$)	sulfite ($\text{SO}_3^{2-}$)
		hydrogen sulfate ($\text{HSO}_4^-$)	sulfate ($\text{SO}_4^{2-}$)
		hydroxide ($\text{OH}^-$)
		nitrite ($\text{NO}_2^-$)
		nitrate ($\text{NO}_3^-$)
		permanganate ($\text{MnO}_4^-$)

A formula expressed in terms of the simplest whole-number ratio of particles is called an empirical formula.

The diagram shows how charge balance works with polyatomic ions in magnesium nitrate and iron(III) sulfate.

Formulas involving elements with multiple electrovalencies

The electrovalency of an ion refers to the charge on that ion. Some transition metals, including copper, lead, iron and tin, can form ions with several different electrovalencies. For example, copper can form $\text{Cu}^+$ ions (charge of $1+$) and $\text{Cu}^{2+}$ ions (charge of $2+$).

Other metals with variable electrovalency include:

• Lead: $\text{Pb}^{2+}$ and $\text{Pb}^{4+}$
• Iron: $\text{Fe}^{2+}$ and $\text{Fe}^{3+}$
• Tin: $\text{Sn}^{2+}$ and $\text{Sn}^{4+}$

Naming ionic compounds

Standard naming rules for ionic compounds are: