Hello and welcome to this video on topic two for the Edexcel specification bonding and structure. My name is Chris Harris and I'm from AlleryTutors.com and the whole point of this video is basically just to go through topic two as a revision tool and so we're going to give you an overview of the topic to make sure that you've got everything that you need now the slides that i'm using here and you can purchase them for a very reasonable price if you just click on the link in the description box you'll be able to access them there but they they are a really good value because you can use them for your revision you can use many smartphone you can cycle through take notes from them do whatever you want with them print them them off you know so it's really good it's just not it's just a kind of to supplement the notes that you've already got in revision material that you've got okay so like I say this is specifically geared toward an Excel specification and it meets these specification points here that you can see on here so basically it matches topic 2 for edXL okay so let's make a start and look at ionic bonding okay so ionic bonding basically it's where you've got your obviously ions and these are opposite charge lines you've got positives and negatives and the held together by electrostatic attractions. Okay so these are obviously the positive ions. Now ions are formed when an element loses an electron and another element gains an electron to get a full shell of electrons. Now in this case you can see that sodium's got one electron and chlorine's got the seven and the sodium will give up one electron and underneath you've got a full shell underneath and chlorine will then have a full shell as well and then obviously they'll form ions and then the opposites will attract.
So like this. There you go. And we put square brackets around them to show that they're ions. So positive around the sodium, negative around the chlorine, and basically they attract.
It's an electrostatic attraction. Now we can obviously use the periodic table to work out the types of ions. So group 1s form 1 plus ions. Group 2s form 2 pluses. 3s form 3 pluses.
Now when you get to group 5, there are 5 electrons in or out a shell. It's easier for group 5s to... gain three electrons and to lose five okay it's energetically lower to gain the three so that's what it does and that's why they form three minus ions group sixes form two minuses and group sevens form one minuses you also need to know your molecular high molecular ions as well so we've got hydroxides which are oh minus nitrates are no3 minus ammonium is nh4 plus Sulfates are SO2-and Carbonates are CO3-.
So make sure you know the molecular ions because obviously you're going to use them quite a bit in your compounds. Okay, and we can obviously work out the formula of ionic compounds and we can use something called the swap and drop method. So the first thing we need to do is write down our two ions. That's Calcium and 2+, and NO3-so I'm just picking these as an example. We then swap the charges over.
So you can see we've got a minus on the Calcium now and 2+, on the Nitrate. Then we drop the charges, so we get rid of the minus and the plus, and drop the numbers so they're lower. Now you can see I've put brackets around the NO3 now, because we want both two lots of NO3 and put two outside, and then the calcium.
And then we just merge them to form the formula, and we simplify if we need to, but this is already the simplest form. So it's CaNO3 2 calcium nitrate. So I've got another example, Ca2 plus and O2 minus.
Again, same thing, swap the charges over, so put 2 minus and 2 plus, drop the charges, Ca2 and O2, combine and simplify. So Ca2O2 can be simplified to CaO, so we call this calcium oxide. So you can see it's a very neat method of just working out formulas of ionic compounds. Ionic compounds such as sodium chloride, for example, they have giant ionic structures. These are big, big structures.
Here's an example of a sodium chloride structure. And they are a regular structure, so we've got these nice neat rows of chloride ions and sodium ions. It's a cubic shape, and we've got this giant repeating pattern as well that just keeps on repeating and repeating and repeating.
So they're big, giant structures. Most ionic compounds, they dissolve in water. This is because, obviously, water molecules are polar.
They attract the positive ions and the negative ions, and we'll see that later on as well. And basically, because they can attract... The different parts of the ionic compound, they are soluble. They do conduct electricity when they're molten as well, because the ions are free to move around. And it's the same in solution as well.
When we dissolve it in solution, the ions are free to move around. So therefore, it can conduct electricity. And they have really high melting points. There's lots of strong electrostatic forces between oppositely charged ions. So lots of energy is needed to overcome these forces.
Got to make sure you're including all these words. they're all underlined there make sure that you're describing why sodium chloride has a high melting point Okay, they're also quite brittle as well and basically when we hit them with a hammer and we strike them the layers will slide over each other. But what then happens is your positives will then be right next door to a positive and they'll repel and so they'll break apart really easy.
And this explains why ionic compounds like sodium chloride are brittle because of this repulsion when you knock them and you get the same ions close by. So two positives together and two negatives together. Okay, so the size of the charge on the ion affects the strength of the ionic bond. Okay, so the bigger the charge on the ion, the stronger the electrostatic attraction between the ions.
And so more energy is required to overcome these forces, and so they have high melting and boiling points. Okay, so if we look at potassium chloride as an example, so it's made up of K plus and Cl minus ions. It has a melting point of 770 degrees Celsius. But if we look at calcium oxide, which is made of calcium 2 plus, O2 minus ions it has a melting point of 2572. Now if you look at the ion difference there calcium is obviously a bigger 2 plus and oxygen is an O2 minus.
There's a bigger charge difference there so therefore they have a much much higher melting point compared to something like potassium chloride. Okay so the size of the ion which is the ionic radii that affects the strength of the ionic bond too. So the smaller the ion the stronger the electrostatic attraction between the ions. Smaller ions these can pack together much more closely more energy is needed to overcome these forces much much stronger forces and therefore ions which are smaller have higher melting and boiling points sodium chloride for example is made up of Na plus and Cl-It has a melting point of 801 however if we look at potassium chloride now potassium is a bigger ion It has a melting point that is 770 degrees Celsius, so it's lower because the potassium is a bigger ion So therefore, the attraction isn't as great.
And generally, the smaller the ion and the higher the charge, the stronger the electrostatic attraction. So we have high melting points. The word we use to describe this is charge density.
We say they have a high charge density. So a big charge and a small ion has got a bigger charge density than a small charge on a big ion. So remember that word, charge density.
Okay, right. So let's look at some trends. So the ionic radius increases as we go down a group. Okay, so as we go down the group, the number of electron shells increases, so the ionic radius increases too. So that's quite logical really.
So you might have learned that from topic one. So the ionic radius in isoelectric ions decreases as the atomic number increases. Okay, right, so we need to have a look at some of these. So isoelectric ions, these are just different atoms.
have the same number of electrons okay so obviously they would form ions so let's have a look at these so we've got here's your ions n3 minus o2 minus F minus na plus mg2 plus Al three plus so these are all isoelectric because they have the same number of electrons but look at the number of protons here okay it's increasing and then look at the ionic radius as well the ionic radius gets smaller okay and so we need to know why And the reason why is because we've got these increased attractive force because we have more protons. Look at the number of protons. It's increasing. This is pulling these atoms in a little bit more and they have the same number of electrons. So the outer electrons pulled in a little bit more.
Remember, the outer shell is the same for all of these ions in this table here because they're all isoelectric. So you're looking at the number of protons here. Okay. So we can look for evidence for charged particles and we can use this kind of nifty little bit of kit here.
We use electrolysis of copper 2 chromate, which is chromate 6, on wet filter paper and it shows evidence for charged particles. So you can see on the diagram here we've got wet filter paper. It's just mounted on a glass slide like a microscope slide. We've connected the electrodes positive and negative to either side.
Okay, so when we put a drop of green copper 2 chromate 6, it's placed on wet filter paper, and electricity is passed through it, it starts to separate out. So you see there's the green copper 2 chromate 6 in the middle. And then what happens when we switch the electricity on, we get the positive copper bit, the Cu2 plus ions, they move towards the negative cathode, okay, because opposites attract. And what you'll see is a blue solution move. So this is quite useful because it's coloured.
Because you can see the actual ions separating out. The yellow chromate 6 ions, these ones here, these have a negative charge, CrO42-. These move towards the positive anode. And you see the yellow solution start to move towards this side. And all this basically shows evidence for charged particles.
Because remember, ions are positive and negative charge. And so if we... pass an electricity through them they can conduct electricity and also obviously they'll separate out according to the charges on the poles okay so let's have a look at covalent bonding so covalent bonding is the sharing of outer electrons okay in order in order for atoms to obtain a full shell so instead of in like ionic bonding when we had an atom moving electrons from one place to another these ones are now sharing so there's an electrostatic attraction there's that word again between the shared electrons and the positive nucleus in the middle of the atom.
And we have single, double, and triple bonds. And basically, this is where we've got more electrons being shared. So in the single bond, we've got two electrons being shared. In the double bond, we've got four. And in the triple bond, we've got six.
And covenant bonds can be represented with lines as well, as you can see here. So this is a displayed formula. And we can also have another type of bond, and we call this a dative.
covalent bond or coordinate bonds and this is where one atom donates both electrons to another atom or an ion so i've got this example here of ammonia this is nh3 with a lone pair of electrons and hydrogen h plus ion um doesn't actually have any electrons so normally this can't form any current bond because it doesn't have any electrons to share so if it's going to bond with something um the electrons have to both electrons have to come from another atom and this is what we call a dative covalent or coordinate bond So let's have a look. There you go. H plus sits on top of there. And we form this date of covalent bond.
Both electrons are being donated from the nitrogen to the H plus ion. And we can symbolize it using an arrow like down here. So this represents a coordinate bond. Remember, the arrow shows the direction of where the electrons are moving from and where they're going to.
So they're going from the nitrogen to the hydrogen. And overall, we've got that positive charge. Also, um... Carbon monoxide, it has a double covalent bond and a coordinate bond.
So it's got both. So just make sure that when you're drawing things specifically for carbon monoxide, that you do have that double bond and a coordinate bond, and you've got both types in that molecule. So as long as you're aware of that, that's the main thing. So let's have a look at bond enthalpy.
So this is the strength of a bond. So bond enthalpy is linked to the length of a bond. okay so how long it is the shorter the bond the higher the bond enthalpy okay so in covenant molecules there are forces of attraction you can see them just forming there okay these are between the positive nuclei and the negative electrons being shared so the green lines here represent the attractive forces which are there so this is the attractive force between the nucleus and the shared electrons so but there are repulsive forces between the two positive nuclei so there that's the red line in between and Between the electrons in the atoms, so we've got electrons here that are not involved in bonding, but they're repelling the electrons that are involved in bonding, so we do have that repulsion. Okay, so there's a balance between these two forces, and the result is something what we call a bond length. Okay, so we've got this balance between attractive forces and repulsive forces.
So the greater the electron density, okay, between the atoms, the stronger the attractive force. So this means that the atoms are pulled further towards each other. And this leads to a shorter bond and a higher bond enthalpy.
Okay, so the shorter the bond, the higher the bond enthalpy. So we've got a big electron density here, loads of electrons. We've got this really strong attractive force, pulls these atoms closer together and strengthens the bond.
So if you have a look at a C-C single bond, this has a pretty low electron density in comparison to a double and even a triple, which has a much higher electron density. And so the... bond length is going to be much shorter for a triple bond than compared to a single bond and therefore the bond enthalpy to break a triple bond requires a lot more energy than it does for a single bond. Okay, so let's look at some shapes of molecules because obviously we're on the topic of covalent bonding here. So we need to know some rules for determining the shapes of molecules.
So we're going to use the number of bond pairs and lone pairs of electrons to work out the shape of the molecule. So molecules have a specific shape with specific angles. And the reason is, is because the bonds repel each other equally.
Okay, so the bonds contain electrons, so they will want to be as far away as possible. okay because they continue these electrons which are repelling each other so here's an example here this is a example of a tetrahedral shape and it's pushing each other apart equally okay so the lone pair next to the bond pairs repel more than the two bond pairs together and two lone pairs repel even further so you can see here there we've got a lone pair this is repelling these bonds much much more squashing these down and so the bond angle reduces in this one we've got 104.5 because it's reduced even further because we've got two lone pairs now really squashing these bonds close together and so lone pairs what they do is they change the shape and the bond angles so lone pairs push bonding pairs closer together as we say generally for every lone pair you reduce the remaining bond angles by two and a half degrees and you can see that's what's happened here but sometimes we don't do that there is some exceptions which we'll show you later on you just have to remember the bond angles for them okay so The rules. So use the number of bond pairs and lone pairs of electrons to work out the shape of a molecule. So what we have to do is draw a dot cross, and this is to work out how many bond pairs and lone pairs we have. So you can see here that we've got four bond pairs here and no lone pairs.
Okay, you've got to be careful with ions as well. So with ions, all we do is we add electrons to the central atom for negative ions and remove them for positive ions. So for example, NH4 plus nitrogen would have four. electrons are normally night in group 5 so the five electrons because it's got this positive charge we just take one electron away we're left with four and all involved in bonding and so would be tetrahedral in this case now this is just a model okay this isn't actually what happened so when we have an eye on we don't actually take an electron away from central atom it's just a method which we work out the the shape a molecule so this is just a model like I say just a method of working out okay so like I say going back to this one Four bond pairs here, no lone pairs, the total is four.
So this tells you it's a tetrahedral. So if you have lone pairs that you need to replace bonds for, then the lone pairs will actually change the angle and the shape of the molecule, but we'll look at them later when we come on to that. So look at this one here, this is water.
And you can see water, we've got two bond pairs here and two lone pairs. Now if we do the same sum as what we did here before, we'd have two bond pairs, two lone pairs, the total is four. This molecule is based on a tetrahedral, but it has two lone pairs.
So we've got this four ton as it's based on that. So what we do is we reduce the bond angle by two lots of two and a half, which is five degrees. Okay, so what we're going to do is we're going to look at some of these.
We're going to look at the ones with no lone pairs first, then we're going to look at the ones with lone pairs. Now for these, you're going to have to remember these names, and you're going to have to remember the bond angles involved as well. Okay. So we're going to use the number of bond pairs and lone pairs of electrons to work out the shape of the molecule.
So once you've got that, you then use this table as a reference, but you will have to remember them, unfortunately, and basically work out the shape of them. So if we've got a situation where we've got two bond pairs and no lone pairs, we form a linear shape. And you can see here's the diagram here, the picture being formed.
The linear has 180 degrees. If we've got three bond pairs and no lone pairs, we form a trigonal planar structure. Now a trigonal planar has this shape and the bond angle is 120 degrees. It's a flat molecule. If we have four bond pairs and no lone pairs, we form a tetrahedral shape.
Now the tetrahedral shape obviously has four different bits coming on. You can see it being constructed there. The bond angle is 109.5 for all tetrahedral shapes. If we've got a molecule with five bond pairs and no lone pairs, then we form a trigonal bipyramidal shape. and you can see on here that we have the trigonal bipyramidal shape being formed and the bond angles has two 120 for that bit 90 degrees for the two axes now these ones are three dimensional and the tetrahedral and trigonal bipyramidal you've got the wedge showing the um atom coming towards you dotted lines showing them going away and then we've got the solid line which shows us in the plane of vision you can see here we've got 120 this is the trigonal bit in the middle basically it's just this turned on its side And then we've got the two poles top and bottom.
Okay, if we've got a situation where we have six bond pairs and no lone pairs, then the shape would be octahedral. And you can see it's being constructed right there now. And the octahedral shape has a square bit in the middle. So it's like a flat square.
You can see here, there it is. And we have 90 degrees between each of these. So it's all 90 degrees and it's called octahedral. Okay, so we've got this shape here. Two coming towards us, two going away, and two in the pole there.
It's called octahedral, by the way, because even though it's got six bonds, if you draw a line from each of these peripheral atoms here on the side up to the axis at the top, you'd have a four-sided pyramid on the top. And if you do the same on the bottom, it'll be four-sided as well. Obviously, the number of sides in total would be eight. That's why we call it octahedral. Okay.
So let's look at the ones with no lone pairs. These get a little bit tricky here. So you need to really remember these. So for three bond pairs and one lone pair, we form a trigonal pyramidal shape.
There it is there, look. So this is an example of ammonia. Look at the bond angle, 107 degrees, because we've reduced that by 2.5 degrees for the lone pair that's on there.
If we have a situation where we've got two bond pairs and two lone pairs, we get a bent molecule. And again... two lone pairs here they repel the bond angle and we get 104.5 degrees remember we're reducing it for every lone pair that we add we reduce it by two and a half degrees so we get 104.5 gets a little bit confusing at this bit okay this is where that that kind of electron rule breaks down that we mentioned so bond pairs three lone pairs two so what we get is a distorted t-shaped and we get something like this here again it would have been 90 degrees at this point here but reduce again we reduce the bond angle by two and a half degrees but just look how we've drawn it okay it's all in the plane but we've just distorted this bit here it looks like a letter t but you reduce it by two and a half degrees from 90 remember it was 90 in that one so it's two and a half degrees so that one's gonna be 87.5 degrees um if you've got four bond pairs and one lone pair And we actually get a seesaw as well.
So let's have a look. There it is Okay. Now again, if you look at the bond angles here, okay, this is where it kind of breaks down You've got to be really on the ball here to know these bond angles for these two one lone pair and four bond pairs And so again you draw your dot cross and if you've got that scenario, this is the type you're looking for So this is called the seesaw mainly because if I flip that on its side, it looks a little bit like a seesaw Okay, five bond pairs and one lone pair. Right, what we get is we get a square pyramidal shape. So, for example, IF5 is an example.
Here it is here. Now, you can see we've got our lone pair on here. And what this is going to do, it's going to push these four up.
And it's going to just close this angle down a little bit. So, it's going to be 81.9 degrees for this angle. 90 degrees remains for the square bit in the middle.
And for four bond pairs and two lone pairs, we get this square planar shape. Now you can see here that we've got two lone pairs of electrons here. Now this is the one where the bond actually remains unchanged because we've got the lone pair here pushing these ones down and this lone pair pushing them ones back up again. So actually it has no effect on the bond angle. So this one remains the same at 90 degrees because they repel equally from opposite sides.
Okay, so let's look at some... As soon as we're still on the covalent side, let's look at some giant covalents. So examples of giant covalent structures include graphite and diamond. Okay, so let's have a look at graphite.
Graphite is a big structure and it contains lots of carbons. Each carbon is bonded three times. And the fourth electron that is involved in carbon is actually delocalized and it forms these layers.
Now we've got lots of strong covalent bonds between the carbon atoms. It means graphite has a really high melting point. However, what we have, like I say, we've got these layers and these slide over each other really easily. You have these weak forces between the layers and that allows it to slide, which makes it ideal for using as a pencil. We have these delocalised electrons as well.
These are between the layers. This allows graphite to unusually, as a non-metal, conduct electricity because the electrons can carry a charge. OK, the layers are really far apart from each other as well compared to...
the covalent bond length and this means that graphite has got a low density. So that's pretty useful because it makes it lightweight so you can carry these things around, like I say in pencils. Graphite's insoluble, it doesn't dissolve, the covalent bonds are far too strong to break for water. So thankfully you can put a pencil in water and it won't actually dissolve.
So that's pretty useful, just in case you wanted to do that. Right, let's look at diamonds. Diamonds are another giant covalent, slightly different structure though. This time each carbon atom is bonded four times, unlike graphite which is just the three. It's tightly packed, rigid arrangement allows the heat to conduct well in diamonds, so that's pretty useful.
But unlike graphite, diamonds can be cut to make gemstones, so you can make it into jewellery. At really high melting point just like graphite loads of strong covalent bonds really hard to To break these because you need so much energy to overcome these strong covalent bonds Diamond doesn't conduct electricity well and it doesn't have any delocalized electrons, but it's an insulator really It doesn't have any delocalized electrons so unlike graphite and again just like graphite Diamond is insoluble Covenant bonds are far too strong to break, so you can put a diamond in water. So if you've got a diamond ring and you're washing up, luckily it won't dissolve in the water, thankfully.
Otherwise it's going to be a very expensive material to pay for. Silicon dioxide is another example. It actually has the same structure as diamond, same arrangement of atoms. And obviously the properties match as well. So as long as you're aware that the structure for silicon dioxide is the same as for diamond, that's the main thing.
Okay, so another example of a giant covalent structure is graphene. Okay, now graphene is a bit unusual, very useful property actually. It's one layer of graphite. So we just take one layer from that, it's one atom thick, it's made of loads of hexagonal carbon rings. As graphene is really, because it's really thin, it's only one cell thick, it's lightweight and it's transparent.
So it makes it ideal for electronic uses. And this is because it has delocalized free moving electrons between this sheet here. Excellent conduct of electricity and they can carry a charge.
So the same delocalised electrons strengthen the covalent bonds. And this gives graphene a really high strength property. So you can twist it and bend it and it's going to be really difficult to break this structure.
And the uses of it. Well, aircraft shells is obviously quite useful because of its strength and it's lightweight. Supercomputers because it conducts electricity. So high speed computing. Very little very low resistance electrical resistance and this stuff compared to silicon and it's using smartphone screens for that reason So you'll find May be in the next 10 20 years.
You're going to get these materials very commonly aware very commonly around You'll probably get these really thin Like sheet like materials and you already seen them in some applications now, for example Google specs and you obviously get these really thin screens that are mounted onto lenses. So you'll see a lot of this as well and particularly the rise of graphene. So what's this space I suppose? Depends when you're watching this I suppose.
Okay metallic bonding is another type of bonding. So these obviously occur in metals only. So they have giant metallic lattice structures.
What they have is positive metal ions. Okay so these are in the middle. These are formed from those to the metal atoms. They donate electrons into a sea of delocalized electrons that you can see here so there's a few of them floating around there is an electrostatic attraction between these positive metal ions and the negative delocalized electrons so there they are there there's your delocalized electrons and they're attracted to the positive metal ions and the more electrons an atom can donate to the delocalized system the higher the melting point so for example magnesium has a higher melting point than sodium because magnesium is in group two can donate to electrons to the delocalized system where sodium can only donate one per atom.
So magnesium has a higher melting point. They're really good thermal conductors because of the delocalized electrons they have. They can transfer this kinetic energy. Remember when we heat these up, we're giving the electrons energy.
They start to vibrate and they knock into neighboring atoms and they pass on this kinetic energy. So that makes them good conductors of heat. Obviously they're good electrical conductors as well because they've got these free moving electrons. They can carry this current and they can move it through the electron or the metallic structure. So obviously no surprises, metals are good conductors of electricity.
They have high melting points as well because they've got these lots of strong electrostatic attractions between the delocalised electrons and the positive metal ions. And they are insoluble. because of the strong electrostatic attraction between the positive metal ions and the delocalized electrons. Now you can obviously see a theme here. There's a lot going on about the strong electrostatic attractions.
And metals are malleable and they're ductile as well, as the ion layers can slide over each other when we hit them with a hammer. And they can still retain this attraction between the delocalized electrons and the positive metal ions. So because this can still be retained despite... distorting the kind of ion structure. This gives metals this property where you can hit them and hammer them into shape and draw them into wires as well.
Okay, electronegativity. So electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond. A really, really important definition that you must, must remember. So the further up and right you go in the periodic table, excluding the noble gases, the more electronegative the element is. So fluorine is the most electronegative element in this series.
So the Pauling scale helps us to quantify. How electronegative an element is, and you can see here that fluorine is the most electronegative. It has a value of 4, whereas on this scale here, hydrogen is the least electronegative with a value of 2.2. And basically, the bigger the difference in electronegativity, the more ionic a compound will be.
And if there's no difference, then it'll be purely covalent. So, for example, Cl2, two chlorines, because they've both got an electronegativity value of 3, the difference is zero so that would be an example of a purely covalent compound whereas something like hf because you've got this big difference here 2.2 and 4 this will have more ionic character um than just purely covalent so even though it would be um a covalent bond there would be some ionic character there because of the difference in electronegativity so yes you can have covalent and ionic character in the same thing effectively that's what i'm trying to say uh right okay polar bonds So covalent bonds can become polar. And it's very important this is only in covalent bonds. If there is a difference in electronegativity.
So the bigger the difference in electronegativity, the more polar the bond will be. So let's have a look at that example that we looked at where we used HF before, but this is HCl. Chlorine's more electronegative.
So what that's going to do is pull these electrons towards itself in this covalent bond. So the electrons are going to be closer towards the chlorine than the hydrogen. We have this delta negative on the top.
Seeing that this is obviously electronegative, delta positive on the hydrogen. So this is basically to show the polarity. And you can see in this example, this is Cl2, the atoms are bonded with the same or similar electronegativity value.
So these both have a value of 3, a powering value of 3. So these are not polar. The shared electrons you can see sit bang in the middle. And hydrocarbons are also classed as non-polar as well.
So even though there might be a slight difference in electronegativity between carbon and hydrogen, they are classed as non-polar. Look out for molecules like this. You might have uneven distribution of electrons. So these are polar.
So water is a classic example of that. Carbon dioxide is a molecule. It appears to be polar. It looks as though it's polar.
But because it's a symmetrical molecule in terms of the distribution of electrons, this isn't actually polar. This is a non-polar molecule because the overall polarity is zero because we have this symmetry. So carbon dioxide is an example of that.
Okay, so intermolecular forces. So we need to know about the three types. I'm going to look at London first. So London, also known as instantaneous dipole-induced dipole, you can see why we just call it London, these forces exist between atoms and molecules. Now the London forces are the weakest ones.
Okay. Now any molecule with electrons can form a dipole when they move near another atom or molecule. So here we're going to use Cl2 as an example. So this occurs as electrons in a molecule or an atom they can move from one end to the other and you can see here that we've got a distortion of electrons. So we've got more electrons at this end of the molecule than we do at this end.
And because we've got more on this side we have this delta negative because more electrons are on that side than on this side. Now this happens when this molecule goes close to another one it might be another chlorine molecule for example and obviously the electron clouds in both molecules will repel each other and so we form this temporary dipole here this instantaneous dipole induced dipole okay so they basically do this but when the molecules move away that interaction is destroyed so you can see here that we've got two I've got few chlorine atoms here now obviously the electrons in here will repel the electrons in the neighboring molecule So that will push them to that end of the molecule. And what we have is this delta negative and positive attracting to each other. This weak attraction is called a London force.
And so this only exists when these two molecules are nearby. When this molecule leaves or moves somewhere else, then the electrons in here will move back to its original position and the whole thing will start again. Okay, an example of a London force is in iodine. Now... these can hold these crystal structures together remember iodine is gray solid it comes in crystals so iodine looks like this you can see this nice regularly arranged structure of i2 molecules Now what we have between these molecules are weak London forces and these hold the I2 molecules together.
However, what we've got between the atoms of iodine, we've got strong covalent bonds and these hold the two iodine atoms together. Now you've got to know the difference between them. A force happens between molecules, hence the word intermolecular, and a bond happens between atoms.
And the bond is much stronger than a force. So when we actually heat this up, iodine actually sublimes. Then what we're doing is we're weakening the London forces, not breaking the covalent bond. So the bigger the molecule or the atom, then the more London forces we have because we've got larger electron clouds.
Remember, it's the electron clouds that get involved with London forces. So when we boil liquid, it could be a liquid such as a fuel, for example, although that would ignite. maybe that's not a good example um but yeah what we're doing is we're breaking the weak london forces not the current bonds okay it's really important when you're talking about boiling points of little simple molecular molecules like this we're talking about forces not bonds okay we've got to have enough energy to overcome these forces and that basically uh determines their boiling points or melting points if it's a solid so longer straight chain hydrocarbons these have more london forces so more energy is needed to overcome these forces This means that their boiling point increases.
So if we start off with something like, say, like butane and compare that with something like decane, which has got 10 carbons in it, that's going to have a higher boiling point than butane because of the fact that it's got more London forces. Branched hydrocarbons, though, we've got these branching here, they can't pack closely together. And so this weakens the London force interaction between these molecules and it lowers their boiling points.
So hydrocarbons with branching have lower boiling points than their counterparts. with the same molecular mass. Okay, let's look at the next strongest intermolecular force. This is called a permanent dipole-dipole. Okay, these exist.
We've got these permanent dipoles here, as you can see, HCl. In other words, these polarity doesn't just exist when it's near another molecule. This one is permanent. So they exist in molecules of the polarity, for example, like HCl. What we have is a weak...
electrostatic force between the molecules a little bit like London forces again delta negative on one side of the molecule delta positive except these are permanent. So the delta negative part one molecule is attracted to the delta positive on the other and then like London forces dipole interactions involves molecules with a permanent dipole and so these forces are stronger than London forces. It is important to note though that these molecules such as HCl for example you that has these permanent dipole-dipoles, they also have London forces as well. So they have both forces.
It's just the strongest intermolecular force is a permanent dipole-dipole. And we can test polar molecules by just running them against a... Basically, we take the polar liquid, could be water, for example, in this case, and we run it through a burette.
And we get a steady stream of the liquid running through the burette. If we put a charged rod next to that steady stream of polar liquid, the liquid should bend towards the rod. So in this case, what's happening here is we've got a positive rod and the negative oxygens, delta negative oxygen on the water here, is attracted to the positive rod and so it'll bend towards it. So we've got this kind of attraction and that's a test for polar molecules or if you've got a polar bond. Okay, and the strongest type.
of intermolecular force is hydrogen bonding okay they have or we get these when we have very electronegative elements okay so we're looking at the likes of nitrogen oxygen and fluorine these are the most electronegative elements and obviously because it's hydrogen bonding we've got to have hydrogen involved in there so what you're looking for is molecules of nitrogen oxygen or fluorine and it must have hydrogen there for it to be able to hydrogen bond so an example is water water can hydrogen bond with each other you What we can do is we can show hydrogen bonding. It's basically an interaction between the lone pair of electrons on an oxygen and a hydrogen on another molecule. When you're drawing these, make sure you draw all your partial charges, your lone pairs as well on the oxygen. You must show that, the interaction and the interaction to a hydrogen.
Hydrogen bonding always occurs between nitrogen, oxygen or fluorine and hydrogen. So it's important to note as well, just like with the permanent dipole-dipole, that any molecules that have hydrogen bonding also have the two other weaker ones before it. So water, for example, has hydrogen bonding, but it also has permanent dipole-dipole and London forces.
So it has all three. It's just its strongest intermolecular force is hydrogen bonding. So just be aware of that.
Okay, so we're going to look at some examples here of hydrogen bonding, but this is in terms of ice. Now, ice is a bit peculiar because when you cool water down, obviously it turns into ice, but it actually expands. It takes up more room.
Now, normally with most materials, the vast majority of materials, when you cool them down, they contract. They get smaller. But ice doesn't.
Okay, and this is why. Ice is actually a regular structure, so it's got loads of water molecules arranged regularly. We have the hydrogen bonding, obviously, as we've seen before, between the delta, the lone pair on the oxygen, and the hydrogen.
But what this does is this actually pushes the water molecules slightly further apart than if they were liquid. And this obviously makes them less dense, and hence the reason why water expands when it's frozen. And obviously, you probably would have seen this as well with ice.
If you put ice in water, it obviously floats. So alcohols, though, are not as volatile as alkanes. with similar masses and this is due to the hydrogen bonding in alcohol so alcohols can hydrogen bond a bit like water can and but if we obviously take a similar mass alkane compared to the alcohol it's not as volatile because of these stronger forces remember alkanes only have london forces their strongest intermolecular force is much weaker and makes them a lot more volatile okay so let's look at some data here you can see the boiling points of hydrogen halides you can see here we've got obviously different hydrogen halides here this one hf it's got the highest boiling point out of a lot of them.
It has hydrogen bonding, loads of energy needed to overcome these stronger forces between the molecules. The rest of them don't have hydrogen bonding. So if we come down here, we've got a mixture between permanent dipole-dipole and London forces.
Because these are a lot weaker than hydrogen bonding, they obviously have lower boiling points. But you'll notice it does kind of gradually increase from HCl to HI. And this is because we have...
As we go along here, we've got bigger halide ion. And that bigger halide ion comes with larger London forces. So that means we've got an increased mass, molecular mass of the molecule.
And we've got a higher intermolecular force between the molecules. Okay, solubility. So remember we talked about last time about these ionic compounds being soluble. So polar substances can dissolve in polar solvents. Okay, so for a substance to dissolve, the solvent...
bonds must break. Okay, so we're looking at the solvent bonds breaking. The substance bonds must break as well. And the new bonds formed between the solvent and the substance. Okay, so we've got a lot of things here.
So basically, we're breaking the solvent, breaking the substance, and then forming new bonds between the two of them. So polar solvents, these are molecules that have a polarity. Some, like water, can hydrogen bond. That's called an aqueous solvent.
And some, like propanone. can have a permanent dipole-dipole interactions and London forces. These are called non-aqueous solvents. These are like organic solvents. So let's have a look.
Most ionic compounds, okay, these can dissolve in polar solvents like water. And basically what's happening is the delta positive on the H is attracted to the negative ions in the ionic compound. And the delta negative on the oxygen is attracted to the positive ions.
And the structure starts to break down. And so this is how powerful water is, really. So you can see here there's an example of it.
So you've got the positive ion here and the waters align themselves where you've got the delta negative oxygen aligning towards the positive ion. It's pretty clever isn't it? And then you've got the delta positive hydrogens here surrounding the negative ion. So what they've done is effectively they've surrounded it and pulled it apart. So we call this hydration.
We've hydrated the salt. So for this to happen the new bonds formed must be the same strength or greater than those broken. Otherwise, there's no point in doing it.
It's got to be energetically favourable to do this. If not, the substance is very unlikely to dissolve. So aluminium oxide, for example, doesn't dissolve. The ionic bonding is too strong. Even though it's obviously an ionic compound and most ionic compounds do dissolve in water, because the ionic bonds between the aluminium and the oxygen is so strong, water can't break it down.
Okay, so let's look at some other examples here. So some non-ionic substances can dissolve too. So alcohols. These dissolve in polar solvents as they can hydrogen bond with water molecules.
So we know obviously alcohols are covalently bonded. But because of this lone pair and the oxygen and the hydrogen on water, they can hydrogen bond. So we've got that interaction. So that allows them to dissolve. But this hydrocarbon part is non-polar.
So this bit of the... alcohol this doesn't dissolve in water and basically the bigger this bit is the less soluble the alcohol is so if we've got an alcohol with 10 carbons on it it's not really going to dissolve very well because it's got such a huge chunk that can't dissolve So some polar molecules don't dissolve in water. Haloalkanes don't dissolve as their dipoles are not very strong.
So they don't really dissolve very well in water. And water forms stronger hydrogen bonds between each other than with the haloalkanes. So haloalkanes are actually insoluble.
So the water is more than happy just interacting with itself rather than bothering with the haloalkane with a weak dipole. So they can dissolve in solvents that interact via permanent dipole-dipole interactions though. So we can do that. So for example, you might have propanol.
Okay, non-polar substances can dissolve best in non-polar solvents. So these are molecules, non-polar solvents, basically are molecules that don't have a polarity. So for example, these are things like alkanes.
They're typically boron molecules, okay? So like butane. All these have is London forces. There's no polarity in there.
So alkanes dissolve best in non-polar solvents. as they can form London forces between the molecules. So nonpolar molecules, these tend to dissolve well in water as water forms stronger hydrogen bonds between each other than interacting with the nonpolar molecules. So again, the water is more interested in interacting with its own molecules than with the nonpolar one. Okay, so let's summarise some of this bonding.
Now, you can see here that we've tried to summarise all of the bonds that we've seen here. So giant covalent, macromolecular. graphite diamond silicon dioxide remember these ones graphene's another one usual state they're solid they don't conduct electricity when they're solid or liquid graphite's the only exception and graphene of course um so because they don't have they're really difficult to melt so don't conduct they're not soluble in water remember them covenant bonds are too strong um and their melting points are high lots of energy needed to break them strong corvina bonds simple molecular there's two types of corvina in here iodine ammonia water little molecules like that normally liquids are gases they don't conduct electricity or when they're solid or liquid soluble water depends on the polarity really and obviously we just looked at solubility just before their melting boiling points are low because we've got weak forces we're not breaking bonds remember when we talk about these we're just weakening forces so when you talk about melting and boiling point we're talking about weak forces giant ionic Sodium chloride, calcium oxide, magnesium bromide.
These are examples of giant ionics. Normally solid at room temperature. Don't conduct electricity when they're solid because the ions are not free to move around.
But when they're as liquid they can, or even dissolved in water they can because they've got free ions. And these allow electrical conduction. Soluble in water, yes they are. Again we've seen that just before.
And the melting and boiling points, they're high. Lots of energy needed to break them. strong electrostatic attractions between the oppositely charged ions.
Metallic type of bonding, metallic. Magnesium, sodium, copper, these are all examples of metallic bonding. Usual state, they're solids obviously because these are metals. They conduct electricity as a solid, yes, because they've got delocalised electrons and they do it as a liquid as well.
They're not soluble in water, bonds are far too strong to break and melting and boiling point is normally high because they have strong electrostatic attractions and so it's very difficult to break them. So this polarity bit here just a reminder that polar molecules dissolve well in polar solvents but like water but your non-polar don't they they dissolve in non polar solvents so like hydrocarbons for example. And that's it that's just a very quick overview of the topic 2 bonding and structure.
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