Lesson Video: Covalent Bonding | Nagwa Lesson Video: Covalent Bonding | Nagwa

Lesson Video: Covalent Bonding Chemistry • 7th Grade

In this video, we will learn how the outer electrons of atoms can be shared to form covalent bonds. We’ll look at the forces between electrons and nuclei, that explain why these bonds form, and look at ways of representing them.


Video Transcript

In this video, we will learn how the outer electrons of atoms can be shared to form covalent bonds. We’ll look at the forces between electrons and nuclei that explain why these bonds form and look at ways of representing them.

A proton is a subatomic particle. Protons have a specific type of charge that we call positive charge. And we say that a single proton has a charge of one plus. An electron is a different type of subatomic particle. Each electron has a charge that is the same size but of the opposite type to the charge of a proton. So electrons are said to be negatively charged with a charge of one minus.

Particles with opposite charges attract each other, while particles with the same charge repel. Positively charged protons attract negatively charged electrons. So when a proton and an electron meet, they can form an atom. The electron doesn’t just stick to the proton for reasons that are far too complicated to go into in this video. Instead, we have the nucleus of the atom deep in the center, and the electron occupies a much larger electron cloud.

When electrons and a nucleus come together, they can lose energy to their surroundings and become more stable. But that doesn’t mean that atoms cannot get any more stable. Atoms can gain or lose electrons forming ions or they can come together to form what we call covalent bonds. Electrons around the nucleus can’t all be in the same space at the same time. When you add electrons to a nucleus, the electrons fill up the space. And like tiers at a stadium, the further from the nucleus you go, the more space there is.

Electrons are said to occupy shells around the nucleus. The first shell can only contain two electrons. The second electron shell can contain a maximum of eight electrons. The bigger the shell, the more electrons that can fit. An atom of hydrogen has one electron and one proton. The electron is quite tightly bound to the hydrogen nucleus with its one plus charge.

An atom of helium has just enough electrons to fill the first shell. These two electrons are held much more tightly to the nucleus, which has a two plus charge than in hydrogen. An atom of lithium has three electrons, so the first two are really tightly bound to the nucleus, which has a three plus charge. And the third electron, forced to be much further from the nucleus by the other two electrons, is much less tightly bound.

An atom of neon has 10 electrons, two in the first shell and eight in the second shell. So atoms of helium and atoms of neon both have full outer shells. But atoms of fluorine have only nine electrons, two in the first shell and seven in the second. This means in an atom of fluorine, there is space in the second shell for one more electron.

When we bring a lithium and fluorine atom together, the outer electron of lithium is in a competition attracted to the lithium nucleus and to the fluorine nucleus. The electron can hop over and the force of attraction between the electron and the nine-plus-charged fluorine nucleus is much much greater than the force of attraction it felt to the lithium nucleus.

While there are many other factors in the reaction of lithium and fluorine, this is a contributing factor. The ion arrangement is more stable than the atom arrangement, so we have a lithium ion, Li+, and a fluorine ion, F-. You’ll generally see this called a fluoride ion. These differently charged ions attract each other in an ionic bond, forming lithium fluoride. However, we’re not specifically looking at ionic bonds in this video. It’s just helpful to see what happens when you take things to extremes. In less extreme cases, rather than atoms becoming ions, atoms can share electrons.

What we have here are two atoms of fluorine. There’s no reason for electrons to jump from one atom to the other because they’re identical. But that doesn’t mean we couldn’t get a configuration that’s even more stable. If two fluorine atoms get close to each other, the outer electrons of each atom will be attracted to the nucleus of the other atom. As the forces balance out, the most stable arrangement emerges where two electrons are in between the atoms being equally shared. By sharing electrons, these two atoms have managed to fill their outer shells without gaining or losing electrons and becoming ions.

This arrangement is much more stable than when the atoms were separated. Chemists have discovered there’s only electrons in the outer shell of atoms that are shared in covalent bonds. We call electrons in the outer shell of an atom or ion the valence electrons. So we sometimes also refer to the outer shell as the valence shell because the valence shell contains valence electrons. So a simple definition of valence electron is an electron in the outer shell of an atom or ion. It’s generally also required that valence electrons be able to participate in bonding, but this simple definition will do.

Atoms of hydrogen have one valence electron but so do lithium atoms because the two electrons in the inner shell are not part of the valence shell. Broadly speaking, the number of valence electrons of an atom determines its chemical character. When atoms bond and share their electrons, they form what we call a covalent bond. Shared electrons are attracted to each nucleus, and these electrons help fill the outer shells of each atom. The co- in covalent means together or sharing or joint, and “valent” refers to the outer electrons involved.

The simplest type of covalent bond is the single bond. In a single covalent bond, two electrons are shared. In a double covalent bond, four valence electrons are shared. And in a triple covalent bond, six electrons are shared. Quadruple bonds where eight electrons are shared are possible, but only between much larger atoms or in highly exotic circumstances. In general, you’ll only need to worry about single, double, and triple bonds. But how do we know how many covalent bonds different atoms are likely to form?

Fluorine atoms have seven out of the maximum of eight valence electrons available because the second shell can only fit eight electrons. A fluorine atom can get that one extra electron from another atom, like another fluorine atom. An atom of oxygen has six valence electrons, so it needs to form a double bond or two single bonds to fill it. And an atom of nitrogen has only five out of the maximum of eight valence electrons. So it needs to form a triple bond, a double and a single bond, or three single bonds to fill its outer shell.

Meanwhile, a carbon atom has four electrons in its outer shell. So we’d expect a carbon atom to form a total of four bonds in order to fill its outer shell. However, carbon generally doesn’t form quadruple bonds. In fact, you can assume that a triple bond is the best a carbon atom can do. Instead, you’ll commonly see carbon atoms sharing electrons with multiple other atoms, for instance, forming four single bonds, two single bonds and one double bond, two double bonds, or one single bond and one triple bond.

If we were reacting pure carbon to start with, the marked electrons would be from the carbon and the unmarked electrons would be from the other atoms. We can look at the first 10 elements on the periodic table and look at their number of valence electrons. This is what we get. What’s more interesting is the number of covalent bonds these atoms will form or the charge of the ions they’ll form when they react. Atoms of hydrogen tend to form one single bond, while helium atoms don’t form bonds at all.

Atoms of lithium and beryllium tend to react and give up their electrons to form Li+ and B2+ ions. And the combining power of these atoms increases up to carbon and then decreases again until we get to neon, which doesn’t form bonds at all. We call these numbers the valency, the combining power of that type of atom. But it will be very annoying if we have to go around and remember the valency for every single atom and deduce it from first principles.

Instead, there’s a rule of thumb that will help us out. And it’s called the octet rule. The first electron shell can fit two electrons. The second can fit eight. The third shell can fit 18 electrons. However, those last 10 slots are a little different, and they are only filled in larger atoms. So as a simplification, we often say that the third shell also contains only eight electrons. This is handy 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. This will tend to mirror the electron configuration of one of the noble gases.

Atoms of hydrogen and helium are exceptions because the first electron shell only contains two electrons. However, the octet rule does account for the most common bonding behavior of these elements, lithium to silicon, and a significant proportion of the behavior of these elements, between phosphorus and radon. It doesn’t make sense to apply the octet rule to the noble gases since they already have full outer shells. And some synthetic elements are simply too unstable to worry about how they’ll react. And things are a little too complicated for the d-block and f-block elements for the octet rule to apply.

The takeaway from this is that atoms will tend to react so they have eight electrons in their outer shell. But if we look at some of these elements, we see that not all of them bond covalently. We can roughly divide the periodic table into metals and nonmetals, and we tend to only see covalent bonding between nonmetal elements.

And we tend to see ionic bonding between metallic and nonmetallic elements. And between metals and metals, we see metallic bonding. Fluorine molecules and hydrogen chloride are example of nonmetal–nonmetal pairings with covalent bonds. And for metal–nonmetal, we have the examples of sodium chloride and lithium fluoride. And we see metallic bonding in alloys like brass, which is a mixture of copper and zinc.

The last thing we need to look at is how we represent covalent bonds because sometimes we can’t draw all the electrons or all the shells. Here are a few of the ways you might see covalent bonds represented. Here are some of the features you might see in various electron shell diagrams. The symbols for the nuclei may be written as element symbols or as charges or numbers of protons. And electrons can be drawn with different shapes and colors in order to indicate where they might’ve come from, for instance, dots and crosses. And for complicated diagrams, it’s sometimes easier if the valence shell is the only shell drawn and the inner shells are left out.

Meanwhile, electron dot or Lewis dot diagrams are useful for condensing all this information into a much easier-to-draw form. Element symbols are used for the nuclei. Dots are used for the electrons. The electrons are arranged in pairs on sides of the element’s symbol resembling the valence shell. And you can recognize bonding pairs by the dots between element symbols. And the other pairs of dots are known as lone pairs.

Double bonds and triple bonds simply require more bonding pairs. And it’s quite common to substitute the bonding pairs with the right number of lines, one line per two electrons. But the most common way to see covalent bonds is as lines without any lone pairs on the atoms. Each type of diagram is useful for a different purpose. Now it’s about time we had some practice.

How many electrons are shared in a double bond between two oxygen atoms?

Oxygen is an element, and we can find information about the element on the periodic table. The atomic number of the element oxygen is eight. This means that oxygen atoms contain eight protons. And since atoms are by definition neutral, we have eight electrons as well, eight electrons to balance out the charge of the eight protons. The question asks about electrons shared in a double bond between two oxygen atoms. Now, oxygen is a nonmetal, so we’d expect a certain type of bonding. And the word “shared” gives us the start of it, co- in covalent.

Now the “valent” part in covalent refers to the valence electrons. Valence electrons are simply those electrons in the outer or valence shell of an atom or ion. An oxygen atom has eight electrons and the first two fill the first electron shell. And the remaining six occupy the second electron shell, but the second electron shell can fit a maximum of eight electrons. We can find more electrons to fill that available space in our second oxygen atom. Since the valence shells are the only ones of interest here, I’m going to remove the inner shell.

If the atoms get closer together, some of the electrons are shared between the two nuclei, helping to produce a more stable configuration. Since this involves sharing of valence electrons, we have a covalent bond. And since there are four electrons involved, we’re dealing with a double covalent bond. The quick way round is just to remember that a double covalent bond contains four electrons, a single contains two, and a triple contains six. So how many electrons are shared in the double bond between two oxygen atoms? Four.

Now let’s have a look at the key points. A covalent bond is a chemical bond involving shared valence electrons. We can predict the number of covalent bonds that would form between different atoms using the octet rule of thumb, which tells us that atoms tend to react to acquire eight valence electrons in total. And nonmetals, either with themselves or other nonmetals, tend to bond covalently. And the main types of covalent bond are single bonds involving two shared electrons, double bonds involving four, and triple bonds involving six.

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