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AP®︎/College Biology
Course: AP®︎/College Biology > Unit 1
Lesson 3: Introduction to biological macromoleculesCovalent bonds
Covalent bonds involve the sharing of electron pairs between atoms. Electron pairs shared between atoms of equal or very similar electronegativity constitute a nonpolar covalent bond (e.g., H–H or C–H), while electrons shared between atoms of unequal electronegativity constitute a polar covalent bond (e.g., H–O). Created by Sal Khan.
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can someone tell me what a diatomic element is? whoever does thanks for doing so(20 votes)
Diatomic means a molecule of 2 atoms bonded together
Most commonly you’ll see it with H, N, O, F, Cl, Br and I which naturally all exist as diatomic molecules: H2, N2, O2, F2, Cl2, Br2 and I2(54 votes)
how do you know if the bond is ionic or covalent?(11 votes)- Look at the difference in electronegativity between the two atoms, if it’s more than about 1.7 then it’s more on the ionic side of things.(33 votes)
Why doesn't the oxygen take an electron from the hydrogen because it is more electronegitive?(10 votes)
That would mean there would be an ionic bond between the oxygen and the hydrogen. While the electronegativity difference between the two elements is large, it isn't large enough to be considered an ionic bond, rather just a polar covalent bond. Hope that helps.(13 votes)
Can someone please explain what the octet rule is?(6 votes)- Atoms tend to react until they have 8 electrons in their outermost shell. 8 = oct.(15 votes)
how do i determine when there is a covalent bond and when there is an ionic one?(3 votes)- a covalent bond is a bond between 2 non metals while an ionic bond is between a metal and a non metal(11 votes)
Can H2O be an ionic bond? In one of the previous videos Sal mentioned that a H atom can either lose or gain an electron. So if H would just give away its electrons to O then wouldn't it be an ionic bond. Or are there other things that make it specifically a covalent bond?
Thanks!(3 votes)- For two atoms to be in an ionic bond, they need to have a large enough electronegativity difference. Electronegativity being a measure of how strongly atoms attract electrons to themselves. If the electronegativity difference between the two bonding atoms is small, they each individually do not have enough strength to steal the other’s electrons and so instead they share it in a covalent bond. If the electronegativity difference is high, then one atom has enough strength to steal the other’s electrons in an ionic bond.
For water, H2O, we have hydrogen and oxygen. We can use the Pauling scale to know their electronegativity values. Oxygen is 3.44 while hydrogen is 2.20. This results in an electronegativity difference of 1.24 (3.44 – 2.20 = 1.24). For the bond to be ionic, the electronegativity difference needs to be around 2.0. A smaller difference like 1.24 indicates a polar covalent bond. So while the oxygen is considerably more electronegative than hydrogen, it isn’t so great a different to result in an ionic bond.
Hope that helps.(8 votes)
Please tell me the difference between single,double and triple covalent bonding(5 votes)- A single bond has 1 bond between the atoms
A double bond has 2 bonds between the atoms
A triple bond has 3 bonds between the atoms(6 votes)
- how is this in collage biology I am learning this in 8th grade(4 votes)
- I think it goes more into depth and gives you a better understanding of the subject.(7 votes)
What happens if Hydrogen wants to share electrons with Lv? Would the electrons spend more time with Hydrogen or Lv?(3 votes)- When you mention Lv, do you mean Livermorium?
Assuming you do, this would require knowing the electronegativities of both elements; hydrogen and livermorium. For hydrogen this is known to be 2.20 on the Pauling scale.
But livermorium is not known to us currently. And the reason for this is that electronegativity calculations for the Pauling scale involve the need of molecules and compounds of an element. For livermorium this is difficult because it is a synthetic element where all the isotopes have very short half-lifes.
Being a synthetic element means that it has only been created by humans in laboratory settings using specialized machines called particle accelerations; so this makes it rare. Additionally all the livermorium isotopes we've created are radioactive meaning they are unstable and decay into other elements. One way to measure how stable a radioactive isotope is using its half-life. The half-life of an isotope is the amount of time for half of the atoms to undergo a decay. The shorter the half-life, the less time the atoms persist in that element. For livermorium the half-lifes are less than a second (milliseconds long). So any atoms we do produce don't stay around long enough for us to do much chemistry on them. And this includes experimentally being able to measure its electronegativity.
So the best we can do is make predictions about livermorium's electronegativity. Elements on the periodic table in the same group (or column) have similar chemical properties. So we can extrapolate using the data from the other elements in the same group (group 16). Going from oxygen to polonium the electronegativity values decrease from 3.44 to 2.0. If we continue this trend we would expect livermorium to have an electronegativity value less than polonium's 2.0 value. This would mean it would be less electronegative compared to hydrogen. So therefore if livermorium did form a molecule with hydrogen (like LvH2), we would expect the electrons to spend more time around the hydrogens than the livermorium atom.
Hope that helps.(7 votes)
- How does that work in phosphate? Phosphor has 5 electrons on its outer shell but somehow in phosphate it has a double bond to one and a single bond to three more oxygens. Isn't that two bonds too many? Or do I get something wrong?(3 votes)
- Phosphorus doesn’t need to follow the octet rule(6 votes)
Video transcript
- [Instructor] In a previous video, we introduced ourselves
to the idea of bonds and the idea of ionic bonds, where one atom essentially is able to take electrons from another atom. But then because one
becomes positively charged and the other becomes negatively charged, they get attracted to each other. Now we're going to go to the other end of the bonding spectrum, where instead of stealing
electrons from each other, we're going to share them. Let's say we're dealing
with two oxygen atoms. So let me draw one oxygen here. A neutral oxygen has
eight electrons total, but six of them are in its outer shell. So it has one, two, three, four, five, six valence electrons. And the way that I arrange
them is I pair them up last. So you have these two valence electrons that are not paired with another electron. And now let me draw another oxygen, and I'm going do it
with a different color, so we can keep track of the electrons. So another oxygen right over there, also has six valence
electrons, one, two, three, four, five, six valence electrons. Now this oxygen on the left, in order to become more stable, it would love to somehow gain or maybe share two more electrons. And of course, this oxygen on
the right, it's still oxygen. It also would love to gain or share two more valence electrons. So how could it do it? Well, what if the oxygen on
the left shared this electron and this electron with
the oxygen on the right, and the oxygen on the
right shared this electron and this electron with
the oxygen on the left? Well, if they did that, you would have something
that looks like this. You have your oxygen on the left. You have the oxygen on the right. And the way we show two
electrons that are being shared, let's say these two
electrons are being shared, is just a line like this. This shows that there are two electrons that are being shared
by these two oxygens. And let's say that these two electrons are also being shared. You would do that with a line like this. And then we could draw the remainder of the valence electrons. This oxygen on the left had, outside of the electrons
that are being shared, it had four more valence electrons. And then the oxygen on the right had four more valence electrons, one, two, three, four. Now what's interesting here
is the shared electrons, these are going to cause these
oxygens to stick together. If they don't stick together, these electrons aren't going to be shared. So what we have formed here
is known as a covalent bond, covalent bond. And what's interesting is it
allows both of these oxygens in some ways to be more stable. From the left oxygen's point of view, it had six valence electrons, but now it's able to share two more. Remember, each of these bonds, each of these lines
represent two electrons. So this oxygen could say, hey, I get to have one, two, three, four, six, eight electrons
that I'm dealing with, and the same thing is going to be true of this oxygen on the right. Now there are some covalent bonds that are between not-so-equals. So for example, if we're
talking about water and if we're talking about how
oxygen bonds with hydrogen. So if we have oxygen right over here, once again, I can draw its six
valence electrons, one, two, three, four, five, and let me just draw the
sixth one right over there. And if I have hydrogen, hydrogen has one valence electron. So let's say that's a
hydrogen right over there with one valence electron, maybe another hydrogen right over there with one valence electron. Oxygen and hydrogen form covalent bonds. In fact, that is how water is formed. And so what would that look like? Well, it would look like this. You have oxygen right over here. You have these two pairs of
electrons that I keep drawing. And then this electron right over here could be shared with the hydrogen, and that hydrogen's
electron could be shared with the oxygen. So that forms a covalent
bond with this hydrogen. And then this electron from the oxygen can be
shared with the hydrogen, and that electron from
the hydrogen can be shared with the oxygen. And so that would form a covalent bond with that other hydrogen. And now here, once again,
oxygen can kinda pretend like it has eight valence electrons, two, four, six, eight. And the hydrogens can kind of pretend that it has two valence electrons. But the one difference here is that oxygen is a lot more electronegative
than hydrogen. It's to the right of hydrogen. It's in this top-right corner, outside of, other than the noble gases, that really like to hog electrons. So what do you think is
going to happen here? Well, the electrons in each
of these covalent bonds are going to hang out around the oxygen more often than around the hydrogen. So if the electrons spend
more time around the oxygen, you're going to have, in general, more negative charge around the oxygen. And so you're going to have
a partial negative charge on the oxygen end of the water molecule, and then you're going to
have partial positive charges on the hydrogen ends of the molecules. And in case you're curious, that little symbol I'm using for partial, that's the lowercase Greek letter delta, which is just the convention in chemistry. And so this type of covalent bond, because there is some polarity, one side has more charge than the other, this is known as a polar covalent bond, polar covalent bond.