The Science of Everything Podcast - Episode 15: Chemical Bonding

Episode Date: February 18, 2011

An overview of how atoms bond together to form different chemical substances, including a discussion of the tree main types of bonds (covalent, ionic, and metallic), and the relationship of these bonding types to the concept of electronegativity. I also discuss the difference between polar and non-polar bonds, and conclude with some interesting applications of bonding theory to understanding the properties of materials.

You're listening to The Science of Everything podcast, episode 15, chemical bonding. Today I'm going to talk about chemical bonding, which is the study of how atoms bind together to form different chemical substances and molecules and so on. First, I'll start off with an introduction to some of the basic concepts that you need to understand, including electro-negativity and the octet rule. And after that, I'll proceed to have a look at the three major types of chemical bonding, ionic bonds, covalent bonds, and metallic bonds. Then I'll also talk about polar bonds versus non-polar bonds and how all that works,

and I'll conclude with some interesting applications of bonding theory to understanding how different materials have different properties and why that is the case. So let's get into it, and I'll start off with having looked at some of the basic concepts of what is chemical bonding and how does it work. A chemical bond is simply an attraction between two atoms that allows the formation of chemical substances that contain two or more atoms. The bond is caused by the electromagnetic force of attraction between oppositely charged particles, specifically the protons in the atomic nucleus and the electrons surrounding the nucleus.

And if you're not familiar with the basics of atomic physics, I'd advise you to listen to episode 8 about the atom and possibly episode 14 about quantum mechanics. That'll help you understand, have a bit of a background for this podcast. Okay, so the fundamental cause of chemical bonds of why atoms combined together to form different molecules, is that electrons tend to move into lower, or more favorable, energy states. As you recall, electrons sit in different energy states, or orbitals, or shells, they're also called, about whatever atom they're in.

The further away these shells are, or these orbitals are, from the nucleus, the higher the energy state the electron is in. And this is because you can think of it as the height of a weight, above the ground. The higher it is above the ground, the further it has to fall, and therefore, the more gravitational potential energy it has. The concept of similar for electrons, the further away they are orbiting from the nucleus, the further they could fall, if you like, towards the nucleus, and therefore the more electromagnetic potential energy that they have. So the point is that orbitals or shells further away from the nucleus have more potential energy, and therefore

they are a higher energy state. Electrons, like everything else in the universe, tend to minimize their free energy or the amount of excess energy that they have. And so they will always tend to move from higher energy states to lower energy states if at all possible. However, because of the Pauley Exclusion Principle that we talked about in podcast 14, it's not possible for more than two electrons to occupy the same orbital around the nucleus. And so if closer orbitals, that is lower energy state orbitals, are all filled up with electrons, then additional electrons have to move to higher and higher. orbitals and therefore orbitals having more and more energy. However, electrons can sometimes get around

this problem if the atom that they are a part of comes close to another atom that has unoccupied orbitals closer to its nucleus. So you can think of it as if the electron moves from a higher energy level in atom one to a lower energy level in atom two. Or sometimes it doesn't move completely from one atom to the other, but it sort of forms a hybrid orbital between the two atoms. And I'll talk more about that in a bit later when I talk about the different types of bonds. But the basic point is that the fundamental cause of chemical bonding of why atoms clumped together to form different molecules and substances is because these electrons are seeking the lowest possible energy state. And if you move atoms close to each other, then that lowest possible energy state may be in the other atom or somewhere in between the two atoms

and not necessarily in the original atom that it was in. and when this happens, both atoms, or it could be more than two, three or four, but let's just say both atoms are now attracted to that negatively charged electron because of their positively charged nucleus, and their mutual attraction to that electron keeps the two atoms together. So you can imagine it as if there are a couple of electrons in the middle and one atom on either side of those electrons. So these electrons in the middle came from, let's say there are four electrons in the middle in between these two atoms, and these four electrons in the middle, and these four electrons in the middle, and these four electrons in the middle, electrons originally, two of them originally came from the left atom, two of them originally came from the right-hand atom. Having lost two electrons each, both of these atoms are now actually ions, they have a charge of positive two, because originally they were neutral, now they're lost two electrons, so they're positive two. So they're positively charged, each of these atoms. And so both of these atoms, or ions, are now attracted to the negatively charged electrons in the middle. And so because they're both attracted to the electrons in

between them, they're sort of both attracted to each other, or at least they're both clumped together, and they form what we call a chemical bond between each other. So we can see from this general principle that chemical bonds are fundamentally caused by electrons seeking more favorable energy states, which in turn leads to the formation of partial or total ions out of the atoms from which they came, from which the electrons came, and this formation of partial or total ions in turn leads to electromagnetic attractions between electrons and those positively negatively charged ions, and that causes the atoms to, or now the ions, to clump together and to form bonds with each other. Okay, so now I'll get into some more specifics.

There are three principles, three main principles that I've identified, and I think it's a good way of organizing the field. Three principles that help us to understand when and why chemical bonds. form. These are electro-negativity, paired versus unpaired electrons, and the octet rule. So I'll go through these one at a time. Electro-negativity is a very important concept. And electro-negativity is a chemical property that describes the tendency of an atom to attract electrons towards itself, or equivalently the tendency to form negative ions. So obviously if the atom is attracting electrons, it's going to get extra electrons and so it will become negatively charged and negatively charged ion. That's called an

an atom's electronegativity is affected by its atomic number and also the distance of its valence electrons, the valence electrons are the electrons in the atom shell, and also the distance of those electrons from the nucleus of the atom. So that might then be complicated, but basically as you move from left to right across the periodic table, electro negativity increases because atomic number is increasing, but as you move down the periodic table, electron negativity decreases because the electrons are getting further and further, the valent electrons, excuse me, are getting further and further away from the nucleus because they're moving out into higher energy orbitals and therefore further away from the nucleus. So if you take those two principles,

electron negativity is lowest at the bottom left-hand side of the periodic table and highest at the top right side of the periodic table. Fluoride is actually the element with the highest electron negativity and oxygen is about number two. That's why oxygen is very reactive because it has very high electronegativity. Okay, but first I should explain why electronegativity is important. I've said it's the detentancy of an atom to attract electrons towards itself. Well, why does this matter, and why does it vary from one element to another? The reason this is important is because electro-negativity plays a massive role in most chemical bonding,

as you would perhaps have imagined. Because suppose you have an element with a low electronegativity, and then an element with a high-electric negativity, And suppose you bring those two atoms of those two elements together. In many circumstances, the higher electronegativity element will attract electrons away from the low electronegativity element, therefore forming two ions. The high electrolyte negativity element will form an anion

because it will be now negatively charged, and the other element will form a cation because it's positively charged. And the positive and negatively charged ions will now be attracted towards each other, and the two atoms, or now ions, will have formed a bond with each other. That's one method of chemical bonding. Not all chemical bonds are exactly like that, but often the exchanges of electrons, which internally to chemical bondings,

are produced originally by differences in electronegativity, basically different pulling powers of electrons. You can bring multiple different atoms together, and if one of those atoms has greater electronegativity, or in other words, great attractive power for electrons, it can pull electrons towards itself away from the other atoms, thereby forming partial or total differences in charge distribution, which then lead to attractive forces between the atoms and a chemical bond forming. Now, before moving on, I just want to explain why the electronegativity varies in the way I described.

The reason electronegativity increases across the right of the periodic table is because every time you move one element to the right along the periodic table, you're increasing the atomic number, and that means you're putting. one extra proton in the nucleus. And for every extra proton in the nucleus, that's an extra additional increment of higher attractive force of the nucleus on electrons. So more protons in the nucleus means a greater pulling power on electrons. And so, of course, electronegativity is going to go up. However, that only works along one period of the periodic table. Down the groups, or down the columns of the periodic table, it works the other way. Because as you move down to a a new row, what happens is, just because the way the periodic table is set up, a new shell,

a new electron shell, begins to be filled. So at the end of each row of the periodic table, the valence shell of electrons becomes filled, and then at the start of a new row, electrons have to go into that new shell and start filling it up. And when this happens, the electrons into the new shell are further away from the nucleus than the electrons in the previous shell. Previously, the additional electrons in that previous shell were kind of just going, you can think of it as in different spots around the sphere that formed that shell. You know, one went on top, one down the bottom, one, this side, one that side. It's a bit more complicated than that because these are probability orbitals and so on, as I talked about in the quantum

mechanics podcast. But you can just imagine that each shell, each electron shell, is sort of like a sphere around the nucleus at a certain distance away from the nucleus, and different electrons occupy different positions within that sphere, but once all those positions are filled, electrons need to start moving into a higher sphere that's further away from the nucleus. And so the point is that when you move into a new shell, the electrons in shells that are closer to the nucleus, so innermost shells, will sort of shield the electrons in the outermost shell from feeling the attractive force of the protons in the nucleus.

And so that reduces the electronegativity of the element, because even though it's got a lot of protons in the nucleus, or more protons in the nucleus than the previous element back in the previous period of the periodic table, it's got another proton, but now it's got all that electrons feeling all that extra shielding of the electrons in the previous shell. And so electronegativity drops off substantially

when you move from one row to another of the periodic table. And so that's why, as you move down the periodic table, electro negativity decreases because you've got more shielding of inner shells, but as you move to the right-hand side, electro-negativity increases because you've got more protons providing more pulling power to electrons. Okay, so that's the concept of electronegativity.

Now, moving on to the second of our three basic concepts that have paired versus unpaired electrons. If you recall from episode 14, each orbital about the nucleus can house two electrons, one spin up, one spin down. Don't worry about what that means. Spin is just a property that they can. can have one and there are two different states up and down. Once there are two electrons in a given

orbital, it's full and additional electrons have to go into another orbital. It is possible, however, for an electron to occupy an orbital by itself, so have only one electron in that orbital instead of the maximum two. And when that happens, that that's called an unpaired electron. Equivalently, when you have two in the orbital, that's called a paired electron. Now, it turns out that when you do all the quantum mechanical calculations and whatever else and work it all out, paired valence electrons are much more stable than unpaired valence electrons. So when you have unpaired valence electrons, they tend to become paired or to pair up with other unpaired valence electrons in other atoms. They can't pair with valence electrons in the same atom because that would disrupt the states

of the different electron orbitals, but they can pair up with valence electrons in other atoms. that are close enough by. And so when they do that, both of these electrons sort of pair up into, in the same orbital, and now they're paired valance electrons, so it's more stable, and so they've reached a lower energy state by doing so,

but by pairing up in that way, they've sort of bound the two atoms, of which they were previously members, they've bound them together, because they're sort of in the same orbital. What they actually do is they form a sort of a hybrid orbital that goes around both of the atoms that were previously separate.

And so because both of these, each of the electrons is now associated with both atoms, not just one electron with one atom and one electron with the other atom as it was previously. And so because both atoms are now attracted to both of the electrons, the atoms are sort of both attracted to each other as well. And so that forms a chemical bond

between the two elements. And so, electro-negativity, we saw as an important cause of many bonds because one element sort of steals electrons completely or partially from another because of high electronegativity. And now we see that having unpaired electrons is another important cause of chemical bonds because electrons tend to pair up,

having two in any given orbital instead of only one. And when they do that, when they pair up, the prospective atoms that they're in tend to form a bond with each other. And the final principle that I'd like to talk about is called the octet rule. Octet just refers to eight. So it's kind of the eight rule, or the rule of eights. And this rule simply states that atoms or elements tend to be stable, most stable, when they have eight electrons in their valence shell or in their outermost electron shell.

Remember, a shell is comprised of a number of different orbitals. So if we have eight electrons in an atom, a shell, we have four orbitals in it. Now, why you might ask is eight a special number? Well, the answer is it's not really a special number. This octet rule is only a rule of thumb. It only works for some elements, but the elements that it works for are like the first 20 or 30 or so, and those include elements like oxygen and hydrogen and magnesium, calcium, fluorine, carbon. Many of the elements that we are concerned with for a wide variety of applications.

So for many of the transition metals and the larger or more exotic elements to the end of the periodic table, this octet rule doesn't work, but it does work for many elements we care about. about, and so we can use it. So there's no intrinsic reason to why this rule is, why this rule applies. It's just sort of a simplified, abstracted version of the results you get when you do all the quantum mechanical calculations of where electron orbitals will be most energetically favourable to be placed and so on. So it's not a law of nature in itself, it's just a rule of thumb, which is pretty good for at least beginning to understand chemical bonds. So basically the idea is that if the valence shell of the atom has eight electrons in it, then that atom is stable and probably

won't form bonds. If it has only six or seven, then it will tend to form bonds by gaining electrons to acquire the eight needed to fill up its outer shell. If it only has one or two electrons in its valence shell, it will tend to lose those electrons, and thus, you know, its next in a most shell will then become its valence shell, and that will have, will then become its valence shell, and that will have eight electrons in it. The row on the far right-hand side of the periodic table contains what are called the noble gases. These are helium, krypton, argon, radon, and xenon.

I think that's all of them. There may be one more. But anyway, these are called noble gases because they tend not to interact with anything. They're very stable. And why is that? It's because they have full outer shells. As you go from one noble gas

to the next element, which will be at the start of the next row down. The fact that it's at the start of a new row is indicating that it's opening up a new shell and it's got one electron in that new shell, so it's no longer got a full valence shell. But if you take that one electron out, move back up, upper row, and to the end of that row,

you're back at the noble gas, which has a full valence shell, and so it's generally unreactive. And so that octet rule is also very powerful and helping to understand when and why bonds form. Okay, so remember we've got electro-negativity, which is pulling power for electrons, that differs between different elements based on their position in the periodic table.

We've got paired versus unpaired electrons. Unpaired electrons tend to combine with other unpaired electrons to form, so it become paired electrons, thereby forming bonds between those two atoms from which the electrons came, and the octet rule tells us that elements are stable and therefore unreactive when they have eight electrons in the atom of shell, and they will tend to react with other atoms, either giving away or accepting or sometimes sharing electrons,

to get eight electrons in the atom of shell, shell. All right, so let's use those three principles to look at the different types of chemical bonds and how they work. The three main types that I'm going to talk about here are ionic bonding, covalent bonding, and metallic bonding. Before I start, I want to emphasize that these are intra-molecular bonds, not inter-molecular bonds. The distinction is that intra-molecular bonds, that's I-N-T-R-A-intra, meaning within, intra-molecular bonds, apply or act within a molecule. atoms within a molecule to form that molecule. So, for example, the attractive forces between hydrogen and oxygen within the water atom, H-2-O, are intra-molecular bonds. There are also inter-molecular bonds,

that's I-N-T-E-R, inter-molecular bonds, that act between molecules, such as the hydrogen bonding, something you may have heard of, that acts between different water molecules to keep an ice cube, for example, of water molecules together. These intra-molecular bondings, bonding forces are different, and generally intra-molecular bonds are stronger than inter-molecular bonds. But the main reason I brought this up is because I want to emphasize that ionic, covalent, and metallic are all intramolecular bonds, and that's all I'm going to be talking about today. Inter-molecular bonds will be the subject of a future podcast. Okay, so first we'll start with ionic bonding. This occurs when an atom loses, completely loses, or completely gains electrons to fill up its

atomose shell. Now, this usually happens between metals and non-metals, so that that is elements from the left-hand side of the periodic table, those are metals, and elements on the right-hand side of the periodic table, those are the non-metals. And the reason it occurs between those two is because metals, to the left of the periodic table, tend to have one or two, maybe three, extra electrons in their valence shell, which by getting rid of those electrons, they could become more stable, and so they'll tend to get rid of those electrons, give them off to some other atom, whereas the electrons on the right-hand side of the periodic table, excuse me, the atoms to the right-hand side of the periodic table,

the non-metals are maybe one, two or three electrons short of having a full out-of-shell. So they'll tend to accept those electrons being donated from the metals. And so what happens? A classic example is sodium and chlorine. Sodium has one electron in its valence shell. Chlorine requires one additional electron to fill up its valet shell. The sodium donates one electron to the chlorine, forming a positively charged sodium ion and a negatively charged chlorine ion. And these oppositely charged ions then attract each other and form sodium chloride, which is common table salt. Now, ionic compounds like sodium chloride.

Sodium chloride is just one example, by the way. There are many ionic compounds which form in pretty much exactly the same way. you have a metal or non-metal, the metal loses electrons, the non-metal gains electrons, and then the resulting ions attract each other, forming a compound. These ionate compounds tend to form big lattices, because what you'll have is, think about it, suppose you had a sodium ion, which is positively charged in the middle, and it will be attracted to a chlorine ion, which is negatively charged, say on the right-hand side. But then on the left-hand side of the sodium, there'll also be another chlorine ion,

and also on the top and the bottom of the sodium and in the front and back of the sodium. So on all six sides of the sodium ion, you'll have a negative-d-charged chlorine ion, because each side of the sodium, if you like, is attracted to the negative-fid-charged chlorine. But the same can be said from the chlorine ions perspective, it is also surrounded by six positively charged sodium ions. And so this lattice-like structure just sort of expands over a very large space, at least on atomic dimensions. And so the lattice is just a repeating pattern.

sodium chlorine, sodium chlorine, sodium chlorine, one after the other, in three dimensions. And so this entire big structure is kind of like all one molecule. There's no place where one ends and the other begins. It's all sort of one big cubic lattice, which is sort of like a molecule. And many gemstones are like this. So these ionic compounds are held together by the electromagnetic force,

and ionic bonds are quite a strong form of bonding. That's why generally these substances, like gemstones, for example, are quite strong, but they're also brittle. The reason they're brittle is because there's no ability or little ability for the atoms to slip or slide, bend relative to each other. If you can do that, if the atoms can slip or slide or bend a bit, then you've got what's called a malleable material. You can hit it, it bends a bit, and then you remove the force, and either it stays or it bends back, but it doesn't shatter, it doesn't break, it's still all right. Ionic compounds don't have that ability, because if you offset one of the rows of sodiums and chlorines relative to the other one,

what you'll find is instead of them alternating sodium chlorine, sodium chlorine, you'll line them up. So you've got two sodium next to each other, two chlorine next to each other, two sodium next to each other, and then they'll repel, because light charges repel. And so if you sort of hit part of the ionic compound lattice, that'll dislodge two of these sections, causing them to repel, and the whole lattice, or at least part of the lattice will shatter. So the bonds are strong, so the lattice is strong, so long as the crystal lattice is strong, so long as it's still whole, but if you hit it hard enough or in the right spot,

it can shatter, and so it's very brittle. This is sort of like a diamond, although diamonds are actually not ionic bonding. That's actually a different form, but the general properties of a diamond, strong and brittle, are characteristic of ionic bonds. And so, for example, substances that are not brittle, or particularly strong, like metals, which are malleable,

or wood, or other things like that, or even biological flesh, these are not formed, for the most part, these are not formed from ionic bonds. That's why they do not have those properties. Okay, so that's ionic bonding. The next form of bonding that I want to talk about is covalent bonding. Now, covalent bond is different to ionic bonding because instead of one electron, sorry, one atom donating electrons to another, one adding up electrons, the other one taking them, the two atoms, or two or more atoms, can share electrons, and the atoms become attracted to each other because they're mutually attracted to the same pair, or pairs of valence electrons.

And so this occurs, for example, in water, H2O. There is one oxygen atom and two hydrogen atoms. Each of the hydrogen atoms has one valence electron, and so the hydrogens will be more stable if they are able to gain an additional electron. Not only will this fill up their atomose shell, but it will also eliminate the unpaired electron, which will then become a paired electron.

Both of those will increase the stability of the hydrogen. an atom. I should just point out in parentheses that the octet rule only applies to, the octet rule does not apply to the innermost shell. There are many different shells that electrons can fit into. The innermost shell, the very first shell, only has an occupation, a capacity of two electrons, or one orbital. So that's kind of the exception, but the rule still applies that it's more, the atom will be more stable with two electrons filling up that innermost shell than with only one or with three. So it's still the basic same idea. It's just the, I guess, the duet rule, that are the octet rule for the innermost shell.

Anyway, so that's why the hydrogen atoms will tend to bond. The oxygen atom has six electrons in its valence shell, and so it will be more stable with an additional two electrons, bringing up to the total eight. That's the octet rule. It has, it may have a couple of unpaired electrons, or it may not. It depends upon exactly how they're arranged into the orbitals.

But more importantly, oxygen is right on the, right to the top right-hand side of the pure, table, it has a very high electro-negativity. In fact, the second highest there is. And so it tends to pull electrons to it with a great force, even from other atoms. And so its electronegativity is much higher than hydrogens, and so it tends to pull the electrons from the two hydrogen atoms towards itself and kind of take them into its own outermost shell. But it doesn't completely take them, it doesn't completely strip them away from the hydrogens, because the hydrogens don't want, the hydrogens will be most stable when they have two electrons in their atom of shell, not no electrons.

So hydrogen stability is maximized when they have two electrons in their own shell, and oxygen stability is maximized when they have eight electrons in their animal shell. So all three of these atoms are, in effect, trying to take the electrons from the other atoms there, but obviously they can't all take them completely. So instead, they share them in effect. And that's what a covalent bond is. the electrons go into form a more complicated orbital, which kind of goes throughout the whole H2O molecule,

and kind of spreads around both atoms, or even all three atoms, depending on the circumstance. And so the electrons sort of spend some time near the hydrogen, sometime near the oxygen. And so they're sort of part of both of the atoms. And so in that circumstance, the electrons count as being in both the hydrogen's valence shell and the oxygen valence shell.

And so both the hydrogen and the oxygen have, full valent shells and so have increased stability. And when it's an increased stability, that's just another way of saying lower free energy. The electrons have moved into a lower energy state, which is preferable, and so that would tend to do that. And this kind of covalent bonding happens, it doesn't happen just in water, it happens in a wide variety of different substances. Particularly organic substances, carbon is king of covalent bonding because it has, because

carbon has four valence electrons, which means it can either lose four electrons or more commonly gain four electrons in order to fill up its valence shell. Now four electrons is a large amount to lose all gain, so it can form lots of bonds more than most other elements can. And because it can form lots of bonds, it can make lots of different materials and interesting stuff, and that's why carbon is so important. But I'll talk more about that later in another podcast. The main point is that covalent bonding occurs mostly between non-metals, because when more than one different type of non-metal, try to are brought together and form

a bond. And this is because the stability of both of the atoms or of all of the atoms will tend to be increased by gaining electrons, but they can't all gain electrons outright at the same time, because there's not enough of them. And so, what they tend to do is share them in this covalent bond.

Now, unlike ionic bonding, covalent bonding does not form big lattices, because there are no distinct ions created in covalent bonding and it is often not a symmetrical process. So in the H2O molecule, you've got the oxygen sort of in the middle and two hydrogens sort of on either side. There's no way where you could have an additional sort of oxygen

on the other side of one of the hydrogens and then another hydrogen there. It just wouldn't work. The orbitals would sort of interfere with each other, and that would not be a stable arrangement or an energy minimizing arrangement. So it's not symmetrical in the way that an ionic bond is where you can just have a big lattice with six neighbors of the opposite charge and six neighbours of the opposite charge on that and so on. So covalent bonds form distinct molecules,

which are separate entities. And the water molecule is an example of that, H2O. You can write a distinct chemical formula for that. There are always two hydrogen atoms and one oxygen atom, no more, no less. Whereas in, say, a sodium chloride compound, I mean, there's always one sodium atom for an ion for every one chlorine atom,

but the total number of atoms could be any, could be four to four billion. and there's, or many more, there's no relumin on that, it's still the same substance, but in water it's always H2O, and that's the same for all covalent substances. They form distinct molecules with unique chemical formulae. And I should point out that this sharing of electrons

doesn't have to happen with only one electron being shared, or two electrons being shared between the two atoms. Four or even six electrons can be shared between the two atoms, forming what are called single, double or even triple bonds between the atoms. Obviously, the more bonds you have, the more shared electrons between two atoms, the stronger is the bonding force. And so the more strongly bound is the molecule. Okay, the final type of bonding I'd like to talk about is metallic bonding.

And this, as you might gather, forms between metals, often the same metal or different metals as well. Now, remember that I said ionic bonds generally form between metals and non-metals, covalence, generally between non-metals. So it makes sense that there's also a bonding type between metals, and that is metallic bonding. Now, once again, the nature of metallic bonding can be understood by the tendency of metals to give up or lose their outermost valence electrons, because they only have a couple, maybe one, two, or three, valence electrons, and if they just lose those few electrons, they can achieve a full valence

shell. Theoretically, they could gain six or seven valence electrons to fill up that animal shell, but that's much more difficult than just losing two or one or two or three. So they tend to do that much more often. And so what happens is all of the metal atoms will sort of be trying to lose electrons at the same time, but none of the other metal atoms obviously will be accepting them because they're all trying to lose them. And so what happens is the electrons become delocalized, which means they're no longer associated

with any particular atom or even a particular con. covalent bond, but they just sort of have an orbital that extends over many adjacent atoms together. So it's not even a specific covalent bond between, say, two iron atoms. The electron just sort of spreads itself out and moves around all throughout this, the lattice or the group of iron or whatever other metal it is, of iron, we'll say it's iron, iron molecules, sorry, iron atoms. And this is called delocalization. The electron becomes delocalized because it's not at any particular location.

And because, and so when the metal atoms lose their outermost valence electrons, they obviously become positively charged. And because they're all positively charged, they are all attracted to this sort of sea, as it's often called, the sea of delocalized electrons. Obviously the electrons are negatively charged. So all of these metal ions are attracted to the sea of electrons. And so they're all sort of kind of, they're all sort of attracted to each other as well at the same time. You can think of it as like the metal ions are embedded in a sea of negative charge. and so they're all kind of stick together in that sense. And so metallic bonding explains why metals are generally malleable and ductile.

So you can hit them, you can bend them, you can draw them out to strings and so on. They don't tend to crack and break like ionic substances do, like rocks. And the reason for this is because there's no set arrangement of atoms within this, within the sea of delocalized electrons. It doesn't really matter if the atoms, you know, if one metal, ion is here or a bit over that way or a bit over this way, as long as it's within that sea. Whereas in an ionic lattice, it does matter. If you moved one ion, just one spot over, it will be repelled by all its neighbours instead of attracted by them and so that the lattice

becomes unstable. It doesn't matter so much with metallic bonding. So it's much easier to bend and move around the metal than it is in an ionic lattice. And once again, there's no distinct molecule that's formed in metallic bonding. It's all sort of one big lattice. Okay, so now that we've gone through the three main types of bonding, I just want to talk about the distinction between polar and non-polar bonds, or polar and non-polar molecules. And this is important because it plays a big role in intermolecular bonds, which, as I said, I'll talk about in a future podcast. So if two atoms in a covalent bond are the same, then they have a sort of a reciprocal attraction to each other. You know, they're exactly the same. An example of this is oxygen. The air, you may know that the air is

like 20% oxygen. It's not actually oxygen atoms though, it's oxygen molecules, which is O2, two oxygen atoms bound together. These are obviously, you know, the two oxenogens are the same, so there's no asymmetry there, there's no difference, and so that is what we call a non-polar molecule. However, if you have, for example, a water molecule, which is a hydrogen, sorry, two hydrogens and one oxygen, there is an asymmetry there. There are different types of atoms. In this case, the oxygen atom has a much higher electronegativity than the hydrogen atoms. And so the oxygen will

tend to, remember I said that the electrons sort of form an orbital around both hydrogen and oxygen atoms. But because of the higher electronegativity, the higher pulling power of the oxygen atom, the electrons will tend to spend more time around the oxygen atom

than the hydrogen atoms. You can think of it as the probability cloud of the atomic orbital of the electron is much more is much denser around the oxygen than it is around the hydrogen. So it's much more likely that you'll find the electron near the oxygen than near the hydrogen. And that produces what is called a partial electric charge or a dipole, a partial charge of the molecule.

It's not, the hydrogen is not an ion. It hasn't been, it hasn't completely lost its electron, but it sort of partially lost it. And so it's got a partial positive charge because it doesn't have the electron as often as it should, in a sense. whereas the oxygen has a partial negative charge because it has, it's got more of the electron than it should. In fact, it's got more of both of the electrons from each of the hydrogen atoms than it should in our example of the water molecule. And so it's got a partial negative charge. And so this is what is called a polar bond. A polar bond just exists whenever there is a covalent bond between two or more elements with different electronegativities. Now when you have, and that occurs,

for example, in a water molecule, but does not occur in an oxygen molecule because they're both the same. both oxygen. Now, that's the difference between polar and non-polar bonds, but there is also a related but slightly different distinction between polar and non-polar molecules. Now, the point to understand is that all non-polar bonds are also non-polar molecules. If you don't have any partial charges within the bond, you can't have them within the molecule. But it is possible to have non-polar molecules even though you have polar bonds. And a good example of that is carbon dioxide. Now, carbon dioxide is a molecule that has one carbon atom and two oxygen atoms. However, because of the way the electron orbitals arrange themselves, the two oxygen atoms sit on opposite

ends of the carbon atom. And so although the oxygen atoms have a high electronegativity than the carbon, and so each of the two carbon oxygen bonds are polar, the polarities exactly cancel each other out because one electron is pulling, it has sort of a partial charge towards one side, of the carbon, but then the other oxygen on the other side, we'll have a partial charge sort of in the other direction, and they exactly cancel each other out. And so the molecule as itself has no polarity to it. There's no partial charges, there's no asymmetry, it's symmetrical, and so it's a non-polar molecule, even though it's made up of polar bonds. That does not happen all the time. There are many molecules where the polarities of the bonds don't cancel

each other out, and water, as I mentioned, is a classic example of that. The, um, the, the non-polar bonds don't exactly cancel each other out because then it's not exact, the molecule is not exactly symmetrical, and so you get a non-polar molecule. The molecule itself has a sort of a slightly positively charged end and a slightly negatively charged end, and so it's called polar, it has a polarity to it. And this brings me to probably the last thing that I'll talk about in this podcast, which is the shape of molecules. You may be wondering, why is it that carbon dioxide is non-polar while water is polar?

Because both of them are just one central atom with, two smaller ones on either side. In one case, it's one oxygen two hydrogen, the other case it's one carbon, two oxygen. Wouldn't it be the case that the two hydrogen atoms would cancel out in their effect on the oxygen atom, and therefore that would be non-polar as well? Well, that's not the case.

The reason is because the water molecule is bent. In other words, the hydrogen atoms are not on directly opposite sides of the oxygen, they're kind of a bit closer together. They're kind of bent towards each other, closer than you would expect. And so the water molecule is not perfectly symmetrical, and that's why it's not non-polar, it's polar.

Why is it bent towards each other? Why is it not symmetrical? Wouldn't symmetry be a lower energy state? Usually atoms or molecules will form symmetrical arrangements because the electrons around the atom or the molecule repel each other, and so they tend to move as far away from each other as possible. So, for example, if you have six electrons around an atom, one will be up, one will be down, one will be left, one will be right,

one will be front, one will be back. And that's perfectly symmetrical, and they'll be as far away from each other as they can. And similarly, if you have four electrons around an atom, the arrangement's a bit more complicated, but they all form a symmetrical arrangement called a pyramidal shape with about 109 degrees or something between each of the bond direction. However, in the case of water, there are four different orbitals around the oxygen atom, four orbitals times two electrons in each makes eight,

which is the number of valence electrons it should have with a full outer shell. However, two of these orbitals are just a non-bonding pairs. They're just two electrons together. So that's a negative charge of two. However, the other two are also two electrons, but they have a proton with them. And that's the proton in the nucleus of the hydrogen atom. So there's an asymmetry here. We've got two orbitals with just electrons and two orbitals with electrons plus the extra hydrogen atom.

And so you can see there's a difference in charge there. Because of that difference in charge, the non-bonding pairs of electrons are a bit more negatively charged. repel the bonding pairs of electrons a bit more strongly, so the bonding pairs of electrons tend to bend a bit towards each other, and that in turn makes the molecule as itself asymmetrical, thereby producing the polarity of the water molecule. It's very important that water does have the shape, because if it didn't, if it did not have this polarity, it would not be able to form what are called hydrogen bonds. That's a form of intermolecular bonding. And I'll explain that in a future podcast. The point I want to emphasizes the shape of the water molecule is very important because without that shape it would not be

able to form these intermolecular bonds. And without those intermolecular bonds, it would not be liquid at room temperature or at the ordinary temperatures on the earth. And if that was not the case, Earth would almost certainly not have been able to support life, because all forms of life that we know of rely on liquid water. So, the shapes of molecules do matter. Another example of the shapes of molecules being highly important is proteins and other organic macromolecules within living organisms, the particular shape of these molecules often determines their precise function within cells, acting as a hormone or an ionic channel letting in different things into the cell or as part of our immune system, recognising particular pathogens or whatever exactly it's doing.

The particular shape determines function, and the particular shape of molecules in turn is often highly impacted by the forms of bonding that make up that molecule. that is the ionic, covalent, metallic, and the polarities of the different bonds within that molecule. And so that's why it's very important to understand chemical bonds, and so that's why it's very important to understand chemical bonds and how they form, why they form, and the different types of bonds. And so that's about all I have for this podcast.

If you enjoyed this episode, please share this podcast with anyone you think might be interested. I'd also love some extra reviews on iTunes. I've only got one so far, and more reviews is always a good thing. If you want to contact me, my email address is Fods12 at gmail.com. I'd love to hear any feedback you have about the podcast or suggestions for changes or future episodes or anything. That's all for this week, and I'll talk to you next time.