Lesson Video: Ionic Bonding Chemistry • 7th Grade

Video Transcript

In this video, we will learn about the electrostatic attraction between ions that we call ionic bonding. We’ll also look at how these bonds result in the distinctive properties of ionic compounds. An ion is the combination of a nucleus containing protons and usually some electrons. Nuclei can also contain neutrons. But since they only contribute mass, they don’t really affect ionic bonding, so we won’t be looking at them in this video.

A proton has a specific type of charge, which we call positive charge, and we say one proton has a charge of one plus. Meanwhile, electrons have equal and opposite charge to protons, so they’re negatively charged with a charge of one minus. Particles with opposite charges attract one another via the electrostatic force, while particles with the same type of charge repel. So protons will naturally attract electrons and form atoms and ions.

By definition, an atom is neutral overall, which means the number of protons in the nucleus must be equal to the number electrons in the electron cloud. When a nucleus and electrons come together, they can lose energy to their surroundings and become more stable. But that doesn’t mean that atoms cannot get even more stable. Atoms can lose electrons and end up with more protons than electrons, forming what we call cations. Or atoms can gain electrons and have more electrons than protons, making them negatively charged anions. In some cases, it’s more stable if we have cations and anions rather than atoms. To understand why, we need to zoom in.

When electrons are added to a nucleus, they start to fill out space and occupy what we call shells. The bigger the shell, the more electrons it can contain. Electrons can’t occupy the same space at the same time, so when electrons are added to the nucleus, the electrons fill up space like tiers at a stadium. The first electron shell can contain only two electrons, but the second shell can contain eight. And as shells get bigger, they can fit more electrons.

Now, as an example, let’s have a look at an atom of fluorine. Fluorine atoms have nine protons, which means the nucleus of a fluorine atom has a charge of nine plus. And since atoms must be neutral, we also have nine electrons. We can add the first two electrons to the inner shell using dots to represent the electrons. And the remaining seven go into the second shell, filling seven out of the available eight spaces. A lithium atom, on the other hand, has only three protons and therefore three electrons. The first two electrons go in the first shell, and the last one sits in the second.

But let’s examine this outer electron for a moment. This outer electron experiences a force of attraction to the nucleus, but it’s only three-plus charged. It also experiences a force of repulsion from the inner electrons. Now, what about an electron in the outer shell of the fluorine atom? This electron experiences a much greater force of attraction to the nine-positive nucleus of the fluorine atom, while still experiencing a similar degree of repulsion from the inner electrons and the other electrons in the same shell. The electrons of the fluorine atom aren’t going anywhere. But what about the lithium atom?

If the lithium and the fluorine atoms get close together, this outer electron of the lithium atom has two possible configurations: being on the lithium atom, being weakly attracted to the three-plus nucleus, or being on the fluorine atom forming a fluorine ion strongly attracted to the nine-plus nucleus. The most stable configuration happens when the electron hops over from the lithium to the fluorine forming F- and Li+. F- is generally known as fluoride. The F- and Li+ ions then attract each other through electrostatic forces, forming LiF or lithium fluoride. The bond between the ions is known as an ionic bond, and substances containing ionic bonds are known as ionic compounds.

In cases where the force of attraction between an outer electron and two different nuclei isn’t as clear, we get covalent bonding where electrons are shared between atoms. However, we won’t be looking at covalent bonding in this video. Instead, we’re going to look at how to predict whether an atom will produce a cation or an anion and what charge it will have.

Chemists have discovered that when atoms form ions, they generally only lose or gain electrons in certain places. Generally, electrons in the inner shells stay where they are. And we either see electrons being added to the outer shell until it’s full, or we see the atom lose electrons forming a cation. The inner shell, now exposed, is considered the outer shell because it’s the shell furthest from the nucleus with at least some electrons in it. If we like, we can remove the empty shell entirely.

I’ve shown this in one example, but bear in mind that most elements commonly exhibit either one or the other behavior: gain electrons to fill a shell or lose electrons to empty a shell. Since it’s electrons in the outer shell that exhibit this special behavior, they get a special name: valence electrons. And the outer shell of an atom or ion is known as the valence shell. However, there’s one more complication: not all the shells contain the same number of electrons; it’d be quite difficult to remember. So fortunately, there’s a little trick.

The first electron shell can fit two electrons, while the second can fit eight. But the third electron shell can actually fit 18 electrons. However, those last 10 slots only get filled when atoms get bigger. So sometimes we say that the third electron shell contains only eight. This is useful because it helps us formulate the octet rule, which tells us that an atom will tend to react to achieve eight electrons in its outer shell. And the final electron configuration will tend to mirror that of a noble gas.

When we pair up elements and use the octet rule, we can predict the chemical behavior of these elements very well although with lithium to boron, they tend to lose electrons to form the electron configuration of helium, which has two electrons in its outer shell. And we can use the octet rule to predict a significant proportion of the chemical behavior of these elements. It doesn’t make sense to apply the octet rule to the noble gases since their atoms already have full outer shells. And some synthetic elements are so rare and unstable that we don’t need to worry about how the rule applies. And the electronic behavior of d- and f-block elements is a little too complicated for the octet rule to be reliable.

We can look at a group of many of the elements and see how many electrons short of a full shell their atoms are or how many electrons they’re likely to lose. When they react, atoms of the alkali metals in group one tend to lose one electron. This gives them a full outer shell, an electron configuration mirroring a noble gas, and a charge of one plus. For the alkaline Earth metals, having two electrons in their outer shell, they’re likely to lose those two and form two-plus ions. Atoms of the elements in group 13, otherwise known as group three, tend to form three-plus ions, losing three electrons. Although the closer to the top of the periodic table the element is, the more likely it is to bond covalently instead of ionically.

Elements in group 14, like carbon and lead, will form four-plus ions under some circumstances and four-minus ions under much rarer ones. And again, elements closer to the top tend to bond covalently more than they do ionically. On the other hand, atoms of elements in group 15 tend to gain three electrons to form three-minus anions. Those of the group 16 elements tend to gain two electrons forming two-minus ions. And the halogens, the elements in group 17 also known as group seven, tend to be found as one-minus ions.

This still leaves us the question of which elements reliably bond ionically and which will bond covalently. We can crudely divide the elements on the periodic table into metallic elements and nonmetallic elements. And as a general rule, we only see ionic bonding when we see a nonmetal and a metal in the same compound. Nonmetals in pure form or paired with other nonmetals tend to bond covalently. And we see metallic bonding for pure metals and alloys. Sodium chloride and lithium fluoride are good examples of nonmetal–metal pairings which exhibit ionic bonding. Fluorine gas and hydrogen chloride exhibit covalent bonding. And we see metallic bonding in pure metals or in alloys like brass which is a mixture of copper and zinc.

The next thing we need to look at is how we represent ionic bonds. One of the most direct ways you can represent ionic bonding is using electron shell diagrams. There are many ways of doing these diagrams, so here are some of the variants. The nuclei may be represented using the respective element symbol or as a charge or a number of protons. The overall charge can be placed outside brackets or attached to the symbol in the middle. Electrons may be displayed with other symbols to indicate where they might have come from. And the inner or nonvalence electron shells may be left out entirely.

We can even simplify things further using Lewis dot or electron dot diagrams. In these diagrams, the expectation is the nuclei will be labeled with elements symbols. But the charges could be on the element symbol or outside some brackets. And valence electrons are dotted around elements’ symbols paired up where possible. Now, before we go on to any examples, we need to look at two more things. And that’s the structures that ions form and the properties that arise because of those structures.

Unlike in covalent bonding, ions attract each other in all directions, and a single ion of one charge can attract multiple ions of the opposite charge. When those ions arrive, then they start attracting positive ions again and the layers build up. And before long, you’ve got a crystal. This regular structure is what we call a lattice. But of course, this drawing is only two-dimensional, and atoms and ions are three-dimensional. This kind of three-dimensional structure is what we mean when we draw these two-dimensional diagrams. These lattices can conceivably go on for almost forever.

So, unlike covalent compounds, ionic compounds do not have a molecular formula. Instead, we find the repeat unit that describes the whole lattice. And the unit formula for the ionic compound is the same as its empirical formula, the simplest ratio of the elements in the compound.

Lastly, we come to the properties of ionic compounds. There’s a huge variety of ionic compounds from simple ones, like sodium chloride, to more complicated ones, like ammonium nitrate. But generally speaking, they’re solids at room temperature and have high melting points due to the strong ionic bonding. And many of them dissolve in water. Water, being polar, has slightly positive and slightly negative bits, which helps stabilize ions in solution. Now, let’s have some practice.

Which of these statements explains why atoms of neon do not form ionic bonds? (A) Atoms of neon have delocalized electrons. (B) Atoms of neon must gain or lose four electrons to form a stable octet in their outer shell. (C) Atoms of neon already have a stable octet in their outer shell. (D) Atoms of neon form dense metallic structures instead. Or (E) Atoms of neon share electrons to form covalent structures.

Neon is an element to be found in group 18 otherwise known as group eight of the periodic table. The atomic number of neon is 10. This means that an atom of neon will contain 10 protons in its nucleus. And since atoms are by definition neutral, we’ll need 10 electrons as well. So our atoms of neon consist of a nucleus with a 10-plus charge surrounded by 10 electrons in an electron cloud. Ionic bonds are electrostatic attractions between positively and negatively charged ions. Atoms can lose or gain electrons to form positive or negative ions.

We need to look at the five statements and find the one that accounts for the fact that atoms of neon do not form ionic bonds. The first thing it will be helpful to remember is that neon is one of the noble gases, which are well-known for not forming any bonds at all, with some exceptions. However, generally speaking, neon goes around as single atoms. Statement (A) suggests that atoms of neon have delocalized electrons. When considering bonding, we generally think about metallic bonding when we consider delocalized electrons. But neon, being a monatomic gas, doesn’t have delocalized electrons, so we can dismiss this answer.

The next statement says that atoms of neon must gain or lose four electrons to form a stable octet in their outer shell. The outer shell is sometimes also known as the valence shell. The word octet in this statement should hint towards the octet rule, which tells us that atoms tend to react to acquire eight electrons in their outer shell. To continue with this, we need to put the electrons in our neon atom into their shells. The first electron shell can fit a maximum of two electrons, leaving us with eight. And the second electron shell can fit all those eight electrons and no more. So we’ve accounted for the locations of all 10 electrons.

Eight electrons is an octet, and the second electron shell for a neon atom is the outer shell. So we don’t need to lose or gain electrons in this shell to form an octet. However, if, for instance, we were dealing with atoms of silicon with electron configuration 2, 8, 4, we could lose four electrons to have 2, 8 or gain four electrons to have 2, 8, 8. Statement (C) says that atoms of neon already have a stable octet in their outer shell. This is true. The octet of electrons in the neon atom is particularly stable, so we don’t see neon reacting to lose or gain electrons and form ionic bonds.

However, just in case, let’s have a look at the other two answers. We’ve already identified that neon is nonmetallic. It’s a monatomic gas generally, so the fourth answer is not correct. And atoms of neon won’t form covalent structures and share electrons because they already have a full outer shell. So the statement that explains why atoms of neon do not form ionic bonds is that atoms of neon already have a stable octet in their outer shell.

To finish off, let’s have a look at the key points. An ionic bond is the electrostatic attraction between two oppositely charged ions. A positively charged ions is known as a cation, and a negatively charged ion is known as an anion. When atoms react, they may gain or lose electrons to form anions or cations. We can use the octet rule to predict whether an atom is likely to form an anion or a cation since atoms tend to react to acquire eight valence electrons.

Ionic structures, structures produced by ionic bonds, consist of lattices of cations and anions in three dimensions. And for ionic compounds, the unit formula is the same as the empirical formula, the simplest ratio of elements in the compound.