Ionic Bond and Ionic Compound & Nomenclature

This is part of preliminary HSC Chemistry course under the topic of Bonding

HSC Chemistry Syllabus

    • Investigate the different chemical structures of atoms and elements, including but not limited to:
    – Ionic networks

        Structure, Properties and Nomenclature of Ionic Compounds

        Ion Formation & Octet Rule

        The octet rule is a fundamental concept in chemistry that refers to the tendency of atoms to have eight electrons in their valence (outermost) shell to achieve a stable electronic arrangement. This rule is based on the observation that atoms of most elements interact in such a way that they can achieve a valence shell with eight electrons, similar to the electron configuration of noble gases, which are inherently stable due to their complete valence shells.

        Atoms of some elements can achieve an octet by gaining or losing electrons to form ions. Metals, for example, tend to lose electrons to form positively charged cations, whereas non-metals tend to gain electrons to form negatively charged anions. For example, sodium (Na) can lose one electron to achieve an octet in its second shell, forming a Na⁺ ion, while chlorine (Cl) can gain an electron to complete its valence shell, forming a Cl⁻ ion.

        Ionic Bond

        When an electron is 'donated' from a metal atom to a non-metal atom to form a cation and anion respectively, an ionic bond is formed between the two ions. For example, Na and Cl atoms commonly form an ionic bond. 

        Ionic bond is typically formed between two atoms with a large difference in electronegativity. This is usually the case between metal atoms (low electronegativity) and non-metal atoms (high electronegativity). The electrons in an ionic bond are more attracted towards the ion that is more electronegative. When a bond contains electrons that are more attracted towards one atom/ion than the other, it is described as polarised or polar.

        The charge on the ions is influenced by the number of transferred electrons needed to achieve a stable noble gas configuration for the ions. 

        For example, let's take a look at calcium and chlorine atoms when they form ionic bonds.

        The electron configuration of calcium and chlorine atoms are:


        Ca: [Ar] `4s^2`

        Cl: [Ne] `3s^2``3p^5`

        To reach a noble gas configuration, the calcium atom has to lose its two valence electrons in the `4s` orbital. Conversely, chlorine atoms, having 7 valence electrons, need an additional electron to complete their outer shell. With two chlorine atoms, the calcium atom can distribute its 2 electrons to fulfil the octet rule for all involved atoms. Therefore, one calcium ion can form ionic bonds with two chlorine ions. This principle is applicable to any other halogens (group 17 element) that may replace chlorine. 

        Ionic compound for calcium chloride


        The diagram above illustrates how a calcium atom reacts with chlorine by donating its two electrons to help complete the octet of the chlorine atoms. This type of diagram is called an electron transfer diagram.

        Ionic Compounds

        Ionic compounds are compounds (two or more elements) consisting of cations and anions. These cations and anions are typically elements of metals and non-metals respectively. Ionic compounds can also be classified as either binary (containing one cation and an anion) or polyatomic (containing multiple cations and anions). 

        Ionic compounds form lattice network structures, where ions organise into a regular, repeating arrangement to generate a crystal. The diagram below illustrates the lattice structure of sodium chloride (compound formed from Na and Cl atoms/ions). It shares a likeness to scaffolding, regularly seen in construction sites.



        Ionic bonds, the "glue" holding ions in the ionic lattices together, form when valence electrons completely transfer from one atom (usually a metal) to another (typically a non-metal). Unlike covalent bonds where electrons are shared, this complete exchange leads to a more stable structure with noble-gas-like electron configurations for both participating atoms. This transfer of electrons results from the electrostatic attraction between oppositely charged ions, which is why ionic bonds are generally stronger than covalent bonds.

        Ionic substances, during crystal formation, align themselves in distinct arrangements known as close packing arrays. These arrays typically assume either a hexagonal or a cubic shape. However, for the scope of HSC Chemistry, we'll primarily focus on cubic structures. An example of a compound exhibiting a cubic packing structure is sodium chloride (NaCl).


        Empirical Formula

        The empirical formula of a compound denotes its atomic or ionic composition, expressed in the simplest whole number ratio.

        The chemical formulae of ionic compounds are usually expressed as empirical formulae. This is because ionic compounds typically form a crystal lattice structure, where ions are arranged in a highly ordered three-dimensional network. This lattice is not made up of discrete molecules but rather a continuous array of ions in fixed ratios, extending throughout the material. The empirical formula reflects this repeating unit within the lattice that maintains the compound's overall neutrality.

        Considering the illustration above, you can observe that in a single cubic unit there are:

        • One entire Na ion located at the centre of the cube (1 Na ion)
        • 12 Na ions, each occupying a quarter of the 12 edges of the cube (`12 \times \frac{1}{4} = 3` Na ions)
        • 8 Cl ions, each taking up an eighth of the cube's corners (`8 \times \frac{1}{8} = 1` Cl ion)
        • 6 chloride ions, each occupying half of the centres of each face of the cell (`6 \times \frac{1}{2} = 3` Cl ions)

        Thus, for each cubic unit of sodium chloride lattice structure, there are 4 sodium ions and 4 chloride ions. This leads to an empirical formula of NaCl for sodium chloride.

        Ionic Compound Nomenclature

        The nomenclature for ionic compounds is straightforward, involving the naming of the cation (positive ion) followed by the anion (negative ion). Here are the foundational principles:

        1. Cation Naming: The cation is named first and is identified by the name of the element. For metals that can form more than one type of cation (transition metals), the charge on the cation is indicated by a Roman numeral in parentheses immediately following the element name.

        2. Anion Naming: The anion is named by taking the root name of the element and adding the suffix "-ide." For example, chlorine becomes chloride, and oxygen becomes oxide.

        Specific Rules for Naming Ionic Compounds

        1. Monatomic Ions: These are ions formed from single atoms. For monatomic cations, use the element name. For monatomic anions, use the root of the element name with the "-ide" suffix.

        2. Polyatomic Ions: These ions consist of two or more atoms bonded together that carry a charge. Polyatomic ions have specific names that must be memorised.

        3. Transition Metals and Variable Charges: For metals that can form cations with different charges, the charge is specified by a Roman numeral in parentheses. For example, iron(II) indicates a Fe2+ ion, while iron(III) indicates a Fe3+ ion.

        4. Naming Compounds with Polyatomic Ions: When naming ionic compounds that include polyatomic ions, the name of the cation comes first, followed by the name of the polyatomic ion. There's no need to change the ending to "-ide" unless the polyatomic ion is a simple anion like hydroxide (OH)


        • NaCl: Sodium chloride, where "sodium" is the name of the cation and "chloride" comes from chlorine with an "-ide" suffix.
        • CuSO4: Copper(II) sulfate, indicating copper with a 2+ charge and the polyatomic sulfate ion.
        • Fe2O3: Iron(III) oxide, where "iron(III)" indicates the Fe3+ cation, and "oxide" comes from oxygen with an "-ide" suffix.

        Properties of Ionic Compounds

        • High Melting Points and Hardness: Ionic crystals, structured by a continuous 3D arrangement of cations and anions, are held together by these robust ionic bonds. Due to the strong lattice structures, ionic solids tend to have high melting points and hardness.


        Diagram showing the consequence of shearing force applied on an ionic lattice structure, leading to like charged ions repelling each other, accounting for the brittleness of ionic substances.


        • Brittle: Ionic compounds are typically brittle. When a force is applied, it can cause the ions in the crystal lattice to shift, aligning ions with like charges next to each other. This repulsion between like charges leads to the fracturing of the crystal.
        • High Density: Ionic compounds tend to have high densities due to the close packing of ions in their crystal lattice. The density is influenced by the sizes of the ions and the type of crystal structure they form.
        • Electrical Conductivity: Ionic lattices, unlike metallic ones, lack free electrons or ions to carry a charge when a voltage is applied. Therefore, in their solid state, they do not conduct electricity. However, when melted or dissolved in an appropriate solvent such as water, the ions detach from the lattice and become mobile charge carriers. Consequently, solutions of ionic substances or molten ionic compounds can conduct electricity. The solubility of ionic substances can vary greatly, depending on the substance's polarity.
        • Thermal Conductivity: The structure of an ionic compound consists of ions held in place by strong electrostatic forces in a crystal lattice. This rigid lattice structure restricts the movement of ions, and since the ions cannot move freely, they cannot transfer kinetic energy (heat) efficiently through the material. As a result, the ability of ionic compounds to conduct heat is relatively poor compared to metals. However, when ionic compounds are melted, their thermal conductivity can increase because the ions become more mobile in the liquid state, allowing for better transfer of kinetic energy. Yet, even in the liquid state, ionic compounds are generally less thermally conductive than metals.
        • Solubility in Water: Many ionic compounds are soluble in water and other polar solvents. This is because the positive and negative poles of water molecules can stabilise the ions, overcoming the ionic bonds and allowing the compound to dissolve. The solubility varies widely among ionic compounds, depending on the specific ion interactions.