Chances are you sprinkle table salt on your meals, use baking soda in your cooking, or encounter calcium carbonate in chalk. Ionic bonding explains how these substances form. Understanding ionic bonds helps us see how different materials are made and why they have certain properties.
When forming bonds, atoms attempt to achieve a stable noble gas configuration by sharing, losing or gaining electrons for each of the atoms associated. In the case of ionic bonds, atoms lose or gain electrons to form ions, which then form electrostatic interactions.
Ions
Ions are charged species formed when electrons are transferred from one atom to another. The charge of an ion depends on:
whether the atom gains electrons to form negatively-charged ions (or anions), or loses electrons to form positively-charged ions (or cations)
the number of electrons gained or lost.
For example, \(\ce{Na}^{+}\) is an \(\ce{Na}\) atom that has lost one electron. \(\ce{O}^{2-}\) is an \(\ce{O}\) atom that has gained two electrons.
Monatomic ions
Monatomic (or monoatomic) ions are ions consisting of a single atom. Ions formed from main group elements are named using the name of the element, followed by the word "ion". The chemical symbol for monatomic ions is written using the chemical symbol for the element, with the charge of the ion written afterwards as a superscript. For instance, the lithium ion is written as \(\ce{Li}^{+}\), the magnesium ion is \(\ce{Mg}^{2+}\) and the aluminium ion is \(\ce{Al}^{3+}\).
Transition metals (\(d\)-block and \(f\)-block) and some \(p\)-block elements can form multiple cations such as \(\ce{Cu}^{+}\) and \(\ce{Cu}^{2+}\), and \(\ce{Pb}^{2+}\) and \(\ce{Pb}^{4+}\). When naming these ions, the ion with the smaller charge is written with an "–ous" suffix, and the larger charge is written with an "–ic". For example, copper would form the cuprous ion (\(\ce{Cu}^{+}\)) and the cupric ion (\(\ce{Cu}^{2+}\)). The charge of the ion can also be written in Roman numerals within parenthesis next to the element name. For example, \(\ce{Cu}^{+}\) is copper(I) ion, and \(\ce{Cu}^{2+}\) is copper(II). There is no space between the element name and the parenthesis.
Anions are named by removing the ending of the element name and replacing it with "–ide". Let's consider the anion formed when chlorine (\(\ce{Cl}\)) gains an electron. The anion formed would be the chloride ion (\(\ce{Cl}^{-}\)).
Polyatomic ions
Polyatomic ions consist of more than one atom. Some example are the \(\ce{SO}_{4}^{2-}\) (sulfate), \(\ce{NO}_{3}^{-}\) (nitrate), \(\ce{PO}_{4}^{3-}\) (phosphate) and \(\ce{NH}_{4}^{+}\) (ammonium) ions. Atoms in polyatomic ions are held together by covalent bonds.
The octet rule
The valence electron configuration of noble gases is considered the most stable. Except for helium, all noble gases have eight electrons in their valence shell (completely filled \(s\) and \(p\) subshells). Main group elements gain or lose electrons to achieve a full outer valence shell containing eight electrons. This is the octet rule.
The main group metals attempt to achieve the electron configuration of the noble gas just before them in the periodic table by losing electrons. For instance, potassium's electron configuration is \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\,4s^{1}\). the potassium atom will lose one \(4s\) electron to become a potassium cation (\(\ce{K}^{+}\)) to achieve a noble gas electron configuration: \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\). This is the same as the electron configuration for argon.
Main group non-metals acquire the electron configuration of the noble gas just after them in the periodic table by gaining electrons. Consider the electron configuration of chlorine: \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{5}\). To achieve a noble gas electron configuration, the chlorine atom gains one \(3p\) electron to become a chloride anion (\(\ce{Cl}^{-}\)). The electron configuration for the chloride ion is \(1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{6}\), which is also the same as the electron configuration of argon.
In general, metals form cations and non-metals form anions.
Exceptions to the octet rule
In some molecules, atoms do not follow the octet rule. Atoms such as sulfur, phosphorus and chlorine can deviate from the octet rule by forming molecules in which there are more than eight electrons in the valence shell. This is due to the availability of a vacant \(d\) orbital in the third shell.
Ionic compounds
Ionic compounds are held together by electrostatic interactions between cations and anions involved. The cations and anions are arranged in such a way that gaps between ions are filled efficiently and permit maximum electrostatic interactions. Ionic compounds are polar materials.
Formation of ionic compounds
Ionic bonds forms between metals and non-metals through the transfer of electrons from metals to non-metals. Metals provide the cation (by losing electrons), and non-metals provide the anion (by gaining electrons).
Diagram showing the transfer of electrons from Atom 1 (metal) to Atom 2 (non-metal). Electron transfer results in Atom A becoming positively charged and Atom B negatively charged. They interact via an ionic bond to form an ionic compound.
Generally, chemical bonds formed between atoms with larger electronegativity (e.g. elements in groups 16 and 17 of the periodic table) and smaller electronegativity (e.g. elements in groups 1 and 2 of the periodic table) are considered ionic. As an example, \(\ce{NaCl}\) consists of \(\ce{Na}\) which has an electronegativity value of \(0.93\) and \(\ce{Cl}\) which has an electronegativity of 3.16. The difference between these values is \(2.23\).
For an ionic compound to be formed, the difference between electronegativities of two participating atoms must be larger than or equal to \(2\). This permits strong electrostatic interactions between the electrons of one atom and the nucleus of the other atom.
Writing chemical formulas for ionic compounds
Ionic compounds are neutral species. Ions involved in an ionic compound combine in such a ratio that the net charge becomes zero. The ratio of ions in an ionic compound can be represented using a chemical formula, using subscripts after each chemical symbol. This displays the lowest possible ratio in which ions can combine to form the compound. Examples are \(\ce{MgO}\) and \(\ce{CaCl2}\).
Parentheses are used to enclose polyatomic ions when more than one polyatomic ion is involved in an ionic compound. In this case, the subscript is written outside the parentheses. Examples are \(\ce{Fe2(SO4)3}\) and \(\ce{(NH4)2CO3}\).
Naming ionic compounds
Ionic compounds can be named by following some easy steps. Let's look at how we can name the compound \(\ce{CaCl2}\).
Name the cation by writing the name of the metal element. The cation in \(\ce{CaCl2}\) would simply be "calcium".
Name the anion by writing the name of the non-metal element. Remove the ending of the element name and replacing it with "–ide". The anion in \(\ce{CaCl2}\) would be "chloride".
Transitional metals and some of the \(p\) block elements have more than one charge. Therefore, if the cation is from those groups, you must include the charge within parentheses next to its name in Roman numerals or use "–ous" and "–ic" to indicate smaller charge and larger charge, respectively. For instance, \(\ce{CuCl}\) is written as copper(I) chloride or cuprous chloride, \(\ce{CuCl2}\) is written as copper(II) chloride or cupric chloride.
Physical properties of ionic compounds
Ionic compounds exhibit a range of distinctive properties that make them essential in various applications. The table below lists some properties of ionic compounds and examples of how the properties are applied.
Property
Example application
Durable and beautiful as crystalline solids
Potassium aluminium sulfate is used in traditional bookbinding to increase the toughness and lustre of decorative paste papers.
Conduct electricity in solution
Lithium salts in lithium–ion batteries are used to conduct electricity by facilitating the flow of charge.
High melting and boiling points
Magnesium oxide is used in furnace linings because it can withstand extreme temperatures.
Soluble in polar solvents
Potassium chloride is a medication used in the management and treatment of hypokalemia. Its solubility in bodily fluids allows it to be absorbed more effectively.
Example 1 – determining the number of electrons gained or lost
Determine how many electrons an aluminium atom has to lose or gain to obtain a noble gas electron configuration.
Step 1: Write the electron configuration for aluminium.
\[1s^{2}\,2s^{2}\,2p^{6}\,3s^{2}\,3p^{1}\]
Step 2: Identify the nearest noble gas. Write its electron configuration.
The nearest noble gas is \(\ce{Ne}\), which has an electron configuration of \(1s^{2}\,2s^{2}\,2p^{6}\).
Step 3: The purpose of losing or gaining an electron is to acquire an octet in the valence shell. Identify the number of electrons in the valence shell and deduce whether aluminium will gain or lose electrons to achieve an octet.
The valence shell contains three electrons: two in the \(3s\) orbital and one in the \(3p\) orbital. It can gain the electron configuration of \(\ce{Ne}\) by losing these three electrons.
The neutral \(\ce{O}\) atom has eight electrons and eights protons. Explain how \(\ce{O}\) atom acquires valence shell octet.
Step 1: Write the electron configuration for oxygen.
\[1s^{2}2s^{2}2p^{4}\]
Step 2: Identify the nearest noble gas. Write its electron configuration.
The nearest noble gas is \(\ce{Ne}\), which has an electron configuration of \(1s^{2}\,2s^{2}\,2p^{6}\).
Step 3: The purpose of losing or gaining an electron is to acquire an octet in the valence shell. Identify the number of electrons in the valence shell and deduce whether aluminium will gain or lose electrons to achieve an octet.
The valence shell contains six electrons: two in the \(2s\) orbital and four in the \(2p\) orbital. It can gain the electron configuration of \(\ce{Ne}\) by gaining two electrons.
Example 3 – writing chemical formulas
Determine the chemical formula of the compound produced when \(\ce{Ca}^{2+}\) and \(\ce{F}^{-}\) ions become electrostatically attracted.
Step 1: Write the symbol of the cation with its charge, followed by the anion with its charge.
\[\ce{Ca}^{2+}\ce{F}^{-}\]
Step 2: Cross each ion's charge (only the numerical value, not the sign) and write it down as subscripts. The calcium ion's charge will be written next to fluoride ion, and the fluoride ion's charge will be written next to calcium ion, both as subscripts. Where the subscript is a \(1\), simply exclude it.
\(\ce{Ca}_{1}\ce{F}_{2}\), which becomes \(\ce{CaF2}\)
Write the chemical formula for a compound composed of \(\ce{Al}^{3+}\) and \(\ce{O}^{2-}\).
Step 1: Write the symbol of the cation with its charge, followed by the anion with its charge.
\[\ce{Al}^{3+}\ce{O}^{2-}\]
Step 2: Cross each ion's charge (only the numerical value, not the sign) and write it down as subscripts. The aluminium ion's charge will be written next to oxygen ion, and the oxygen ion's charge will be written next to aluminium ion, both as subscripts.
\[\ce{Al}_{2}\ce{O}_{3}\]
Write the chemical formula of the compound formed by copper(II) ions and sulfate ions.
Step 1: Write the symbol of the cation with its charge, followed by the anion with its charge.
\[\ce{Cu}^{2+}\ce{SO4}^{2-}\]
Step 2: Cross each ion's charge (only the numerical value, not the sign) and write it down as subscripts. The copper(II) ion's charge will be written next to sulfate ion, and the sulfate ion's charge will be written next to copper(II) ion, both as subscripts. Since the sulfate ion is a polyatomic ion, brackets must be put around its chemical formula.
\[\ce{Cu}_{2}(\ce{SO4})_{2}\]
Step 3: Since the chemical formula is the lowest possible ratio of the elements in a compound, we need to simplify. Here, we can cancel out the subscripts as they are the same.
\[\ce{CuSO}_{4}\]
Your turn – ionic bonding
Understanding ionic bonding is essential for grasping how elements interact to form compounds with unique properties. Test your understanding of ionic bonding with this quiz.
