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However, other elements don't have this stable arrangement of electrons. Instead, they must gain, lose or share electrons, in order to achieve that optimal configuration. Ionic bonding is such way of doing this.
- This article is about ionic bonding in physical chemistry.
- We'll define ionic bonding before looking at examples and diagrams of common ionic compounds.
- We'll then explore giant ionic lattices and their properties.
- After that, we'll consider the strength of ionic bonding, and ionic radius.
- Finally, we'll understand the evidence for ionic bonding.
Ionic bonding definition
There are multiple ways in which atoms can achieve their goal of a noble gas structure. Non-metals often join up in pairs, trios or larger groups, and share their outer shell electrons (take a look at Covalent Bond). A group of metal atoms of the same element will lose electrons to form positive ions in a sea of delocalised electrons (see Metallic Bonding). But when a metal and a non-metal come together, the easiest way for them to both obtain a full outer shell is for one of the species to lose electrons, and the other to gain them. The transfer of electrons forms ions, and oppositely charged ions bond ionically with each other.
An ionic bond is the electrostatic attraction between oppositely charged ions.
Let’s explore the process a little further.
Ions
Ions are atoms that have gained or lost one or more electrons to form a charged particle.
Ionic bonding always occurs between positively charged ions, called cations, and negatively charged ions, called anions. In both cases, the ions have the electron configuration of a noble gas.
- One element loses electrons. Because electrons are negative, this results in a cation.
- The other element gains these electrons. This results in an anion.
- Both ions end up with full outer shells of electrons.
- In ionic bonding, the cation is always a metal and the anion is always a non-metal. Overall, the charges on the ions cancel out, making a neutral compound.
- We call this transfer of electrons electrovalence.
Electrostatic attraction
Forming ions is only half the picture - By definition, ionic bonding doesn’t involve the transfer of electrons at all! Rather, it is about the interaction between these ions as a result of gaining or losing electrons.
When two oppositely charged species are close by, they attract each other. This is known as electrostatic attraction; you might also remember that this is the force that attracts electrons towards the nucleus in an atom. When mixed together, cations and anions are electrostatically attracted to one another, and ionic bonding is simply another term for this attraction.
Ionic bonding example and diagram
We now know what an ionic bond is: the electrostatic attraction between oppositely charged ions. Let's now consider some examples. We'll learn how to work out the charges of ions, the formula of an ionic compound, and represent the overall ionic bonding in a dot and cross diagram. Here's the process:
- First of all, we determine how many electrons each element needs to lose or gain to achieve a full outer shell. Metals lose electrons, whilst non-metals gain electrons.
- This also tells us the charge of the ion formed - remember that gaining electrons results in a negative charge, whilst losing electrons results in a positive charge.
- We then use the number of electrons lost or gained by each element to work out their ratio in the compound - this gives us its chemical formula.
- Finally, we draw the ions using dot and cross diagrams. These show their new electron configurations. We position the ions inside of square brackets and write the charge of the ion on the outside.
You might see dot and cross diagrams shown with just the outer shell of electrons. However, we've included the inner shells as well, to help you fully understand the ion's electron configuration.
Let's consider a few examples.
Represent the ionic bonding within sodium chloride using a dot and cross diagram. State the charge of each ion and give the chemical formula of the compound.
Sodium chloride is made up of positive sodium cations and negative chloride anions. We can use our knowledge of their electron configurations to work out the charges of the ions they form. Sodium has the electron configuration of 1s2 2s2 2p6 3s1. The easiest way for it to achieve a full outer shell is by losing one electron from its 3s sub-shell, so it has the arrangement 1s2 2s2 2p6. As you know, electrons are negatively charged, and so losing an electron, results in a positive ion with a charge of +1. We show this using a dot and cross diagram:
A dot and cross diagram of a sodium ion. StudySmarter Originals
Chlorine, however, has the structure 1s2 2s2 2p6 3s2 3p5. In order to have a full outer shell, it needs to gain one electron. In fact, it takes the electron that sodium loses. This forms a negative ion with the electron configuration 1s2 2s2 2p6 3s2 3p6:
A dot and cross diagram of a chloride ion. StudySmarter Originals
Each sodium atom loses one electron to form a positive sodium ion with a charge of 1+, whilst each chlorine atom accepts one electron to form a negative chloride ion with a charge of 1-. Therefore, the ions form a compound with a 1:1 ratio of sodium ions to chloride ions. This has the formula NaCl:
A dot and cross diagram of the ionic compound NaCl. StudySmarter Originals
Here's another example. This time, two electrons are transferred between the ions.
Represent the ionic bonding in magnesium oxide using a dot and cross diagram. Include the charge of each ion and the formula of the compound.
Magnesium has an electron configuration of 1s2 2s2 2p6 3s2. To achieve a full outer shell, each atom needs to lose two electrons from its 3s sub-shell. This forms a cation with a charge of 2+ and an electron configuration of 1s2 2s2 2p6.
Oxygen, however, has the electron configuration 1s2 2s2 2p4. Each atom needs to gain two electrons to form an anion with a charge of 2- and an electron configuration of 1s2 2s2 2p6.
Note that each magnesium atom loses two electrons, whilst each oxygen atom gains two electrons. The ratio of magnesium ions to oxygen ions is therefore 1:1, giving us the formula MgO. Here's the final dot and cross diagram:
A dot and cross diagram of the ionic compound MgO. StudySmarter Originals
However, some compounds don’t have a simple 1:1 ratio of cations to anions. An example is calcium fluoride.
Give the electron configuration of the ions in calcium fluoride. State how many electrons each atom loses or gains and give the chemical formula of the compound.
Calcium has the electron configuration [Ar] 4s2. To achieve a full outer shell, each calcium atom needs to lose two electrons from its 4s sub-shell, giving each calcium ion the electron configuration [Ar].
Fluorine has the electron configuration [He] 2s2 2p5. Each fluorine atom needs to gain one electron to form a fluoride ion with the electron configuration [He] 2s2 2p6, which we can write as [Ne].
Note that whilst each calcium atom loses two electrons, each fluorine atom gains just one. Therefore, we need twice as many fluorine atoms as calcium atoms. This gives calcium fluoride the formula CaF2.
Not familiar with electron configuration and its associated notation? Check out "Electron Configuration" for more information.
Giant ionic lattices
Ionic compounds don’t form molecules. Instead, they form structures known as giant ionic lattices. Giant simply means we don’t know exactly how many of each ion it has, just that it has a large number of each - the lattice could stretch on infinitely! However, we do know the ratio of the ions. In sodium chloride, as explored above, the ratio of sodium ions to chloride ions is 1:1. The compound forms a repeating lattice stretching in all directions. Each positively charged ion bonds ionically to all the negative ions surrounding it, not just to one particular ion, and vice versa, as shown below:
Part of NaCl's giant ionic lattice. StudySmarter Originals
Note that the lattice stretches infinitely in all directions, and that each ion bonds to up to six oppositely charged ions.
Properties of ionic bonding
Ionic bonds are very strong. This means that they require a lot of energy to overcome. We also know that all ionically-bound species form giant ionic compounds, made up of oppositely charged ions held together by strong ionic bonds in all directions. This gives giant ionic compounds certain properties:
- Giant ionic compounds have high melting and boiling points because the electrostatic attraction between ions is strong and requires a lot of energy to overcome. Because of this, they are generally solid at room temperature.
- The charged ions in giant ionic compounds can form bonds with polar water molecules. The energy released overcomes the ionic bonds holding the lattice together and dissolves the compound, meaning giant ionic compounds are soluble in water.
- When molten or in aqueous solution, ionic compounds can conduct electricity. This is because the ions are free to move and carry a charge.
- Giant ionic compounds are hard and strong due to the high strength of the electrostatic attraction between oppositely charged ions.
- Ionic compounds are fairly brittle. If you give them a sharp blow, you may distort the carefully positioned lattice structure. This results in two ions with the same charge adjacent to each other. These ions would repel each other and shatter the compound.
Strength of ionic bonding
We know that an ionic bond is the electrostatic attraction between oppositely charged ions. The strength of ionic bonding all depends on the strength of this attraction. This means that not all ionic compounds are created equally - some are stronger than others:
- Ions with a greater charge experience stronger ionic bonding. This is because the attraction between them and oppositely charged ions is much stronger.
- Smaller ions experience stronger ionic bonding. This is because there is less distance between the nucleus and the outer shell electrons and so the attraction between them is stronger.
Compare the strength of the ionic bonds formed by:
- Aluminium and magnesium.
- Sodium and potassium.
Aluminium forms 3+ ions whereas magnesium forms 2+ ions. As a result, aluminium forms much stronger ionic bonds than magnesium.
Sodium and potassium both form ions with a charge of 1+. However, sodium is a much smaller ion than potassium and so forms stronger ionic bonds.
Ionic radius
Ionic radius does not only depend on the ion's number of electron shells, but also on the ion's charge.
Firstly, ions with more electron shells have larger ionic radii than ions with fewer electron shells. This means that as you move down a group in the periodic table, ionic radius increases.
But ions with the same number of electron shells can have different ionic radii. In fact, when an atom turns into an ion, its radius changes. This is all thanks to the gain or loss of electrons:
- Positive cations have a smaller radius than their original atoms. This is for two reasons. Firstly, to achieve a full outer shell of electrons, they need to lose their entire outer shell. This immediately decreases their radius. But because they have lost electrons, they also now have a higher proton:electron ratio. This increases the attraction between the positive protons in the nucleus and the negative electron shells, pulling the electrons closer to the centre of the ion.
- Negative anions have a larger radius than their original atoms. This is because they have gained electrons, which lowers their proton:electron ratio. This decreases the attraction between the positive protons in the nucleus and the negative electron shells, so the electrons drift further away from the centre of the ion. There is also increased repulsion between the electrons due to the presence of the gained electrons.
Describe the trend in ionic radius for the series of isoelectronic ions from N3- to Al3+.
Isoelectronic ions are ions that have the same electron configuration. This means that they all have the same number of electron shells. If they were all atoms, you'd expect them to have the same radii. However, these are ions with different charges, and so they have different ionic radii. N3- is the most negative ion and so has the largest ionic radius. In contrast, Al3+ is the most positive ion and so has the smallest ionic radius.
Evidence for ionic bonding
To conclude this article, we'll consider the evidence for ionic bonding.
We know that ions exist thanks to electrolysis. When you apply an electrical current to an ionic solution, the ions move to the oppositely charged electrode. If the ions are coloured, it is easy to see this movement.
Electrolysis of NaCl. StudySmarter Originals
Ionic Bonding - Key takeaways
- Ionic bonding is the electrostatic attraction between oppositely charged ions.
- Ionic compounds are formed from metal cations and non-metal anions. The metal gives electrons to the non-metal so that both achieve a noble gas configuration.
- Examples of ionic compounds are NaCl, MgO and CaF2.
- Ionic compounds form giant ionic lattices, not molecules.
- Ionic bonds are strong and require a lot of energy to overcome. This makes ionic compounds hard, strong and brittle. They're also soluble in water and good conductors of heat when molten or in aqueous solution.
- The strength of ionic bonding depends on the size of the ion and the charge of the ion.
- Ionic radius depends on the number of electron shells and the charge of the ion.
- Electrolysis provides evidence for ionic bonding and the existence of ions.
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Frequently Asked Questions about Ionic Bonding
What is an ionic bond?
An ionic bond is the electrostatic attraction between oppositely charged ions.
How are ionic bonds formed?
Ionic bonds are formed when a metal donates electrons to a non-metal, forming charged ions, which are then attracted to each other.
What happens to electrons in an ionic bond?
An ionic bond does not involve electrons. Rather, the bond is the electrostatic attraction between charged ions. These ions are formed through the movement of electrons.
What is the difference between ionic and covalent bonds?
A covalent bond is a shared pair of electrons, whilst an ionic bond is the electrostatic attraction between oppositely charged ions.
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