Bonding Revision (AQA)

[Music] hello and welcome to this revision video this is for AQA and in this video we&#39;re going to look at bonding our whole point in this is just to go through and give you an overview of of the the topic of bonding for AQA and these videos are very specific to AQA as well so if you&#39;re studying area that&#39;d be great also just before I start these powerpoints that I&#39;m going to be using on here you can purchase them if you want to use them for a revision or if you want to print them out or look at them in your own time etc if you look in the comments box of the description box of this video you&#39;ll see the link there where you can where you can get ahold of them okay so I can say these videos are dedicated to AQA and they are tailored to the specification and respect points are all on there and so ok so let&#39;s have a look at the first this is ionic bonding okay so only bonding or basically where you&#39;ve got charged ions and they&#39;re held together by these very strong electrostatic attractions these are oppositely charged ions and basically that&#39;s all an ionic bond is and it works in a really simple way really and so to get a full shell of electrons in this case we&#39;ve got sodium and chlorine as you can see here it is your surgeon coins get a full shell of electrons thus owed you will have to give up an electron to chlorine so it gets a full shell sodium has an empty shell or I&#39;ll have a full one underneath this shell here and the becomes stable by it being attracted to each other by forming an oppositely by forming an electrostatic attraction between the oppositely charged ions so let&#39;s have a look there&#39;s the electron and the you can see it jumps onto the coin and we form these ions I notice we draw box as well and square boxes and we put a plus and the negative charge around them to show that they are oppositely charged okay and these ions obviously depend on the group or depend on what type of iron that you form so if you look in Group 1 group one&#39;s all form plus one is group to form 2 plus 3s on three pluses Falls very very rarely really form ionic bonds they&#39;re covalent with each other so we don&#39;t bother with them group 5 right these form 3 minus ions remember it&#39;s easy for these to gain the 3 electrons then it is to get rid of 5 sets 1 upon 3 minuses group 6 form two minuses with 7 4 1 minuses okay you need to know if you these molecular ions as well so these are a little bit different so hydroxide is o h minus nitrates is no.3 minus and ammonium is NH 4 plus sulfate is so4 2 minus and carbonate is co 3 2 - make sure you know these you&#39;ve got to these are so so important because you&#39;ll see them a lot in chemistry obviously and so admit make sure you know the charges of them is one positive on there which is ammonium for some a negative ok right so let&#39;s have a look a little bit more at some ionic bonding I&#39;m not sure what we can do is you can work out the formula of an ionic compound and we can use this method called the swap and drop method I&#39;m going to show you two examples here so the first thing you have to do is write down the two ions in this case we&#39;re going to look at calcium ion and the nitrate ion as you can see there what we then do swap the charges over so the two now moves over to the nitrate and the - moves over to the calcium then drop the charges that includes the plus and minus bit and we drop them so that they are in lower case so the subscripts so that&#39;s what happens there notice the minus is obviously gone from calcium and then we just simplified the whole number in this case this is ca no.3 - so it&#39;s calcium nitrate and that&#39;s basically how we work them out so that are slightly different want to stick with calcium and this time we&#39;re gonna use oxygen again swap the charges over to - + - + drop the charges CA 2 + o2 and then simplify so we&#39;re left with CA 202 but we can simplify that to just see 800 calcium oxide so you can do this at any type of ion it&#39;s dead easy to see okay really quick and easy okay I&#39;m only compounds for example sodium car they have giant ionic structures so they form things like this here here&#39;s an example sodium Clyde the yellow circles or spheres are so divine chloride ions you can see they pack together in a regular structure very cubic shape giant repeating pattern you&#39;ve got to know and the description of the structure of these things is pretty important these things dissolved in water pretty well and they water&#39;s polar it can be attract the positive and negative ions obviously the Delta negative oxygen and water is a trap to the positive sodium ion and the Delta positive hydrogen&#39;s are attracted to the chloride ions and so water can break these structures up they can also conduct electricity when they&#39;re molten or dissolved in solution because the ions are free to move around when they&#39;re solid like this they can&#39;t because the ions are not free to move around as you can see they&#39;re fixed to the cubic shape and they have very high melting points because we have lots of strong electrostatic attractions and these basically require lots of energy to overcome them so and this is why sodium Clyde and other ionic compounds have such high melting points very important to include things like strong electrostatic forces and lots of energy needed to overcome them so make sure you get all these keywords and your answers when you write in them ok Corinna balling is another type of bonding right this is a little bit more trickier actually even though you might think it might look simple and basically career pollens are sharing about two electrons okay very important and we basically do this very similar to ionic compounds to get a full shell of electrons except this is sharing electrons we&#39;re not giving up or receiving electrons these are the sharing of them and there is still an electrostatic attraction there&#39;s that word again this time it&#39;s between the shared electrons and the positive nucleus in the middle that&#39;s where the attraction is between these so that&#39;s what&#39;s going to hold these things together we can get single double and triple bonds and basically you can see there this is a single that&#39;s a double and that&#39;s triple so we&#39;re sharing three electrons each and and the covalent bonds can also be represented by lines and you might have seen these in displayed formula okay dative covalent bonds on the other hand are a little bit different so and these are also known as coordinate bonds and this is where an ass donates two electrons or pair of electrons to another atom in this case we&#39;re going to look at ammonia with a lone pair of electrons and the hydrogen ion H+ with no electrons in this can&#39;t form a covalent bond because now so you having the electrons to share so if it wants to bond with another atom both the electrons are bit you can have to come from the other atom which is a bit unfortunate so these are called date of credit and coordinate bonds and if you see on here there&#39;s the hydrogen look it&#39;s now sharing this electron sharing the electrons with nitrogen but both of them have come from the nitrogen to the H+ ions make sure you same where they come from where they go to so there it is there and it can be represented with an arrow and the owl shows the direction of where the electrons have been donated from and where they&#39;re going to so make sure you get the arrow the right way around okay and the other type giant covalent so they were all simple molecular very simple small ones giant covalent are a lot bigger and these include things like graphite and diamond so that&#39;s look at graphite first graphite is basically made up of hexagons and each carbon is bonded three times and the fourth electrons delocalized you can see here look if you look at a carbon one two three times delocalized electrons are being delocalized so lots of strong covalent bonds means it&#39;s got a very high melting point members really loads of energy to break these bonds so else can take quite a bit of energy but the layers graphite is very unique in its structures made up of layers there&#39;s weak forces between these layers and they can slide over each other relatively easily and then do localized electrons allow you to conduct electricity they can carry a charge very important that you say it carries a charge they can conduct electricity which is unusual for and obviously nonmetals they carbon the layers are really far apart in comparison to a covalent bond length and so this means graphite is pretty low density compared to other giant covalent structures okay so it&#39;s insoluble doesn&#39;t dissolve and the bonds are far too strong and you can&#39;t break them when you velcome to water so you can you can basically put a pencil in water they will dissolve thankfully just in case you want to do that home okay diamonds diamonds another one diamonds a bit different though it has each carbon has bonded four times instead of graphite where it&#39;s three times and now because this doesn&#39;t have and and because this doesn&#39;t have the gap like graphite does and it means that they are they can conduct heat pretty well because they&#39;re really tightly packed together and diamonds on see and are form from volcanoes so obviously they will be able to conduct and absorb some of this heat and so unlike graphite diamonds can be cut to make gemstones so you can cut them and graduate carcasses just prior to other breaks in their into layers and it&#39;s got a really high melting point bit like graphite and loads of strong covalent bonds you need a lot of energy to overcome them it doesn&#39;t conduct electricity very well though and in fact it doesn&#39;t conduct electricity and at all really didn&#39;t have any free electrons and like graphite does so and it&#39;s non conductive electricity and again just like graphite it&#39;s insoluble the bonds are far too strong for water to break them apart okay so make sure you know these examples of giant covalence okay shapes of molecules this could be a bit tricky this and there&#39;s a very kind of certain rule in which you can apply you might have seen other rules as well but I think this one&#39;s relatively straightforward to use I hope you agree so we&#39;re going to use the number of bond pairs and lone pairs of electrons to work out the shape of a molecule so let&#39;s have a look at this molecule here they&#39;re very specific shape and this is because the bonds repel each other equally and they trying get is far away from each other as possible we&#39;ve got electrons in these bonds they&#39;re like charges so they will repel each it so this means they have a very specific shape so if we have a lone pair at next to a bond pair though these repel more than two bond pairs together and two lone pairs repel even further so that&#39;s one look at an example here&#39;s one here this is one with one lone pair now if you look at the angle of this one with no lone pairs of electrons and this one with a lone pair look the angle has decreased this is pushing these bonds closer together because the low become even further and if we have two lone pairs like this one the bond angle shrinks even further and so you can see here so what they&#39;re doing is they change the shape of the bond angles and lone pairs what they do is they push them closer together and January is a rule not all the climbed over for things like tetrahedral structures which we&#39;ll look at in a minute every time you have a long pair you reduce the remaining bond angle by two and a half degrees so and we have got to be careful for some other molecules where this doesn&#39;t happen because they cancel out but I&#39;ll point it out later on okay so let&#39;s have a look at some of these so we&#39;re going to use the bond pair and lone pair electrons to work out the shape of a molecule so the first thing you should do is really drop your dot cross compound a dot cross model here basically just helps you to work out how many bond pairs and how many lone pairs you&#39;ve got if you&#39;ve got an ionic molecule which they sometimes give you all you do is and you add electrons to the central atom for a negative I&#39;m so and this is if your molecule has negative ID and remove them if it&#39;s a positive ID so for example ammonium nh4 plus the nitrogen would have four electrons normally nitrogen has five but we take one away to take through calculus so it has four and this is only of course this is not what happens in reality we are only using this method to work out the bond angle and the shape okay this is not what happens in reality it&#39;s just a method of working out so and so yes so let&#39;s have a look at this one so the total in this case if we add them all up tells you the shape in this case it is a tetrahedral got four bond pairs there are no lone pairs left on the central carbon so it&#39;s a total of four and so if you and have lone pairs you need to replace the bonds for lone pairs and change the shape and the bond angle which we&#39;ll look at in the middle a got special needs so something like this yeah okay and you can see water obviously this has to bond pairs two lone pairs and but crucially it&#39;s based on a tetrahedral that&#39;s the very important thing is interceptor you draw it&#39;s based on that and all we do is reduce the bond angle by two lots of two and a half words five degrees okay so let&#39;s look at the shapes with no lone pairs first okay so you&#39;ve got to know the names of these shapes and you&#39;ve got to be able to know the angles of them and recognize them etc so we use the bond pairs and lone pairs of electrons to work at the shape of a molecule so let&#39;s look at the first one this one&#39;s got to bond pairs no lone pairs here&#39;s an example and which is beryllium dichloride the name of the shape is linear you can see it&#39;s 180 degrees here and we have two atoms that are obviously and joins either side no lone pairs okay let&#39;s look at the other one bf3 bond pairs three no lone pairs so if you have this scenario we have three bond pairs no lone pairs we have a trigonal planar this is a flat molecule 120 degrees either side okay for bond pairs and no lone pairs these molecules are called tetrahedral it&#39;s the first time when we actually have a 3d shape and so here&#39;s an example of a tetrahedral we have one bond one atom coming towards you one in your plane of vision that one&#39;s in your plane of vision and that one&#39;s going away from me the bond angle is 109 point five four all tetrahedral okay if you have five bond pairs and again all of these got no lone pairs the shape we are looking at is trigonal bi-pyramidal so you can see obviously this is an example is PCL 5 but here&#39;s an example here so you see we&#39;ve got three here there&#39;s the threes you&#39;ve got one two three okay in the plane and then we have to top and bottom as an axis so this bit is planar and it&#39;s triangle shapes and call it trigonal planar but this is trigonal bi-pyramidal and so it&#39;s basically if you look at the trigonal planar one here that&#39;s basically that but just tipped up on its side you can see but it&#39;s by pyramidal because actually if we draw a line going from there to there there to there and there to there we form a pyramid on the top we form a pyramid on the bottom results as well quite a trigonal bi-pyramidal bond angles 120 because that&#39;s basically just that tips on the side but the bond angle between the top the polar ones here and the trigonal plane a bit in the middle is 90 degree so this one&#39;s got two angles so make sure you know the difference okay final one is one with six bond pairs and no lone pairs and this is an example of an octahedral structure it&#39;s very similar to trigonal bi-pyramidal except in the middle as you can see here we have a square shape you can see the square shape here in the middle two axes top and bottom the bond angle here is nice degrees we call octahedral okay so let&#39;s look at them now with em with lone pairs so there&#39;s not as many of these but we&#39;re based on and remember on the original shape so we&#39;re going to use the bond pairs and lone pairs again to work this out so here&#39;s an example this one&#39;s got three bond pairs one lone pair so this is the example of pyramidal an example is ammonia there we go looks a little bit like tetrahedral so it got a lone pair there if we go for the two bond pairs and two lone pairs we get a bent molecule or nonlinear it&#39;s also known as so we&#39;ve got the lone pairs there on the side again look we&#39;ve shrunk the bond angle from 107 to 104 point five so that&#39;s been shrunk by two and a half degrees okay three bond pairs and two lone pairs an example of clf3 trigonal planar now this one 120 again we&#39;ve taken the bonds from the top and the bottom and we get this flat trigonal planar structure two lone pairs notice the bond angle here is still the same as what it was before these lone pairs are repelling each other equally this one repels these bonds down this one repels and back up again so they cancel out and this is exactly the same four and this one with four bond pairs and two lone pairs based on octahedral and you can see here the bond angles 90 degrees lone pairs are pushing equally and so that&#39;s quite important so the bond angle actually remains unchanged the two lone pairs repel equally from opposite sides make sure you know the names of these shapes okay electronegativity so we need to remember this definition electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond and basically the further up and right you go in this and in this here this periodic table the more electronegative the element is excluding the noble gases this is one would black them out so fluorine is the most electronegative element and we can use the another powering scale which helps us to quantify how electronegative something is there&#39;s Florina fluids the most electronegative for B the biggest number here and but as we go further away the electronegativity drops basically the bigger the difference in electronegativity the more polar the bond is so just watch out for this scale in your exam okay so let&#39;s have a look at polar bonds so covalent bonds can become polar the atoms attached to it of a difference in electronegativity so the bigger the difference in electronegativity the more polar will be so looking at this example here we&#39;ve got h and cl bonded together the chlorines the most electronegative is pulling the electrons towards the chlorine more than the hydrogen we give it this little Delta negative symbol to show to show that it&#39;s pulling the electrons towards it hydrogens got the dells positive okay here&#39;s another molecule this atom is basically of this molecule sorry is a molecule of chlorine they&#39;re both just as electronegative as each other so for that reason this is not polar because the electrons sit bang in the middle of the bond and so these things are not polar hydrocarbons are the same they&#39;re classed as non polar as well so like alkanes etc and alkenes so these things and the electrons are just shared equally within that bond sometimes uneven distributions well they do actually uneven distributions of charge leads to polar molecule so water is a classic sign of that see the electrons are now distributed unevenly they&#39;re not spread across the charge and spread equally across the molecule so water is an example of an uneven distribution however if you look at ones where we&#39;ve got a bit of symmetry like here this is a completely symmetrical molecule and so this one are actually nonpolar even though they&#39;ve got polar bonds the electrons have been pushed either side equally so we call this a symmetrical molecule and so these actually have no overall priority she could be really careful for that and what&#39;s out of that okay intermolecular forces these are really really important okay the first one our van der Val&#39;s also known as induced dipole dipole these exist between atoms and molecules but not total bonds here what weak forces between molecules and you need to know these fit into like a little kind of league table of other intermolecular forces and Vander vials are the weakest out of Bolin and they exist in any molecule that basically has electrons in them and so basically evander valve is where a dipole can be created when they move near another atom or molecule so let&#39;s have a look at this example here so you can see that we&#39;ve got an atom here and the electrons in this molecule are not normally unevenly distributed but when they move near another molecule they are you see the electrons have been moved over to one side of the molecule we get this Delta negative and Delta positive on there this dipole is only temporary it&#39;s only there when it&#39;s next to another molecule that&#39;s nearby and obviously that molecule will have electrons as well and basically when they move away the interaction is destroyed so you can see here when this molecule is near another molecule we have a delta negative and a delta positive as soon as this molecule leaves the electrons will shuffle back on to the other side and even itself back out again but for the time that they&#39;re close by there is this weak interaction Delta negative Delta positive that&#39;s really called an induced dipole dipole it&#39;s only brought about when it&#39;s near end of the molecule okay we get this force of attraction um okay so let&#39;s see if we can apply this to something specifically so for example iodine iodine is a classic example van der Waals forces can hold some of these molecules and crystal structures so iodine have this type of crystal structure it&#39;s pretty pretty and it&#39;s a it&#39;s solid at room temperature I D and the key thing is we have weak van der Waals forces holds the iodine molecules together remember it&#39;s I 2 but we have strong coven of bonds that hold the two iodine atoms together don&#39;t get these two confused okay the bigger the molecule the atom the more Van der Waals forces you have so that&#39;s what you&#39;re looking out for because they have larger electron clouds so when we boil the liquid and it could be water for example we&#39;re breaking the weak van der Waals forces and not the covalent bonds actually water has hydrogen bonding as well so I tried break them but when we boil a liquid and that&#39;s what happens so for example bromine we&#39;re just breaking these weak forces not the bonds and never mention the bonds when we talk about these small molecules we&#39;ve gotta have enough energy to overcome these forces and hydrocarbons remember we said we don&#39;t have this polarity so they have bound the valves and basically the longest straight chain hydrocarbons have more van der valve than air forces and energy and then you&#39;re kind of non straight or branched chain and shorter ones so basically these need more energies needed to overcome these forces and this increases the boiling point of these molecules and like I say the branched ones these are hydrocarbons with bits sticking out of them they can&#39;t pack together as much so this means that the van der Waals force is a lot weaker and the and the chain if the chain is shorter as well this lowers the boiling point as well so short the hydrocarbons all hydrocarbons a lot of branching have lower boiling points okay well it&#39;s look at the next in specular force these are called dipole-dipole they&#39;re a little bit stronger than around the valve these exist when you have a molecule as a permanent dipole so basically the dipole exists irrespective whether they&#39;re near another molecule or not so things like where you have a polarity so permanent dipoles you&#39;re looking for polarity HCl is an example of a permanent dipole and you have this weak electrostatic force between a delta negative chlorine the electronegative element the Delta positive hydrogen atom in this case and so the Delta negative is attracted to the Delta positive as we say so in Lake van der Waals forces dipole-dipole interactions involves molecules for the permanent dipole they are stronger than van der Waals it is important to note though that even though this does have this molecule in particular has a permanent dipole it also has banned the valves as well the strongest in specular force is the permanent dipole dipole and we can kind of prove this as well we can prove that polar molecules like water exist and if we take a charged rod and we put it near a steady stream of a polar liquid you can do this in a burette what is an example of polar liquid and then what happened is you should see the liquid bed towards the rod and because obviously the charge difference here and basically the molecule will align the line itself to the oppositely charged rod so you can see here this is the positively charged rod in this example so the Delta negative oxygen will be attracted towards a positively charged rod so it&#39;s a very important little test for polar molecules okay and the lust intermolecular forces hydrogen bonding this is the strongest of the lot the hydrogen bonding is quite unique is that it&#39;s basically a type of dipole-dipole force and you basically get hydrogen bonding when you use very electronegative element so your most electronegative elements are nitrogen oxygen and fluorine it&#39;s these three and hydrogen so any molecules that contain nitrogen oxygen or fluorine and a hydrogen will get involved with hydrogen bonding and it&#39;s basically the lone pair of electrons on these atoms that get involved that plays a very important role so here&#39;s an example of showing hydrogen bonding we&#39;ve got all our partial charges on this or Delta negatives Delta positives there&#39;s the lone pair on the oxygen the interaction is between the lone pair on the electronegative element and the hydrogen and this is a hydrogen bond very common form to ask this and you gum so make sure you know how to do it and again just like with the previous one molecules that have hydrogen bonding also have van der valves and they have dipole-dipole they have all three it&#39;s just the strongest one is hydrogen bonding so make sure you&#39;re very vigilant exam and you watch out for them um so let&#39;s have a look a little bit further ice obviously is just frozen water and it forms this very regular structure held together by hydrogen bonds but ice is strange because when you cool normally many cool things down they get smaller they contract but ice actually expands it gets bigger and the reason why is because we have hydrogen bonding that&#39;s pushing the molecules further apart and this makes ice less dense so if an ice cube in water and it floats and let&#39;s look at the boiling points and properties of these things as well so if you look at the h-f h-f has a higher boiling point than hate see L remember HF is one of the molecules that can hydrogen bond between other molecules because hydrogen bonding is stronger it needs more energy to overcome these forces looking at the other ones things like HCl HCl can&#39;t hydrogen bond it drops massively less energies needed to break these however it starts to pick up again when we go for things like HBR and H I and this is because and yes they have dipole-dipole forces all of these do but the biggest effect have the biggest effect that&#39;s having effect on the boiling point is basically the band of ours is increasing the out and getting bigger the molecules getting bigger there&#39;s more electrons Barneveld forces increases so therefore the boiling point increases from heat CL thi okay metallic bonding so metallic bonding and these are giant metallic lattice structures so these are quite large structures they have a very unique setup you have positive net lines which is the green circles there they&#39;ve formed when metals donate electrons into the sea of delocalized electrons and there&#39;s basic electrostatic attraction between the positive metal lines and then delocalized electrons and basically the more electrons and assam can donate the higher the melting point and so for example magnesium has a high melting point of sodium it can donate two electrons into the delocalized cloud where sodium can only donate one so therefore the forces of attraction between the positive and negative charges are a lot weaker metals are good thermal conductors as you&#39;d expect and they have delocalized electrons so they can transfer this kinetic energy very easily remember when you heat things the the particles move and electrons move it&#39;s the electrons that allow the conduction of heat there are also good electrical conductors because they can carry a charge or current they can move that through the lattice there are free electrons delocalized electrons okay they have high melting points again strong electrostatic attractions between the delocalized electrons and the positive metal line and solid metals are insoluble then the metallic bond is far too strong to break and don&#39;t bleed and otherwise you&#39;ll get metal structures dissolving when it rains so that&#39;s pretty useful okay so quickly this particle model F which straightforward this bit really and solids Titan packed regular arrangement really high density you can see the arrangement of these here they vibrate on the spot and they can&#39;t be compressed liquids they are tightly packed again they don&#39;t have many gaps between them random arrangement high density they move around freely and they slide over each other and they really difficult to compress because they&#39;re tightly packed and the particles have a little bit more energy in them in liquid than you do in solids and that&#39;s why they&#39;re moving randomly gases well they&#39;re very spaced out random arrangement and very low density particles able to move around freely and they have a lot more energy than liquids and and solids obviously because they&#39;re moving out freely so don&#39;t really comment on them that starts the bathing and finally just summarizing the bond types and make sure you can summarize these all really it&#39;s a good way of just making sure that you know all these things giant covalent graphite diamond silicon dioxide for example normal temperature their solids they don&#39;t conduct apart from graphite that don&#39;t conduct electricity as liquids very difficult they normally sublime anyway they don&#39;t really melt soluble in water high melting points give you lots of energy to break them simple moleculars these are covalent as well liquids or gases normally won&#39;t conduct electricity and they might be soluble depending on the polarity the low breaking low melting and boiling points and because you&#39;re breaking weak forces giant ionic much much higher melting points and generally soluble more so because the polar they will conduct if they&#39;re dissolved in water or molten and normally they&#39;re solids very high melting points strong electrostatic forces and metallic it&#39;s pretty much the same except they will conduct solid and liquid conduct electricity and which is unusual and like the rest of them and a usual temperature normal state is solid again high melting pipes because they have strong electrostatic forces and so the polarity you&#39;ve got to be watch out for this polarity zapala molecules they do dissolve well in polar solvents like water and but nonpolar molecules don&#39;t hydrocarbons you need nonpolar solvents really for them to dissolve so there we go that is the end of the on the video I thought I hope that was a useful summary of bonding is quite a bit their bond angles are pretty important and just again if you want a copy of this of this PowerPoint be good for revision and just look in the the description box for this video and click on the link there and you get a hold it in there that&#39;s it bye bye

[Music] hello and welcome to this revision video this is for AQA and in this video we&amp;#39;re going to look at bonding our whole point in this is just to go through and give you an overview of of the the topic of bonding for AQA and these videos are very specific to AQA as well so if you&amp;#39;re studying area that&amp;#39;d be great also just before I start these powerpoints that I&amp;#39;m going to be using on here you can purchase them if you want to use them for a revision or if you want to print them out or look at them in your own time etc if you look in the comments box of the description box of this video you&amp;#39;ll see the link there where you can where you can get ahold of them okay so I can say these videos are dedicated to AQA and they are tailored to the specification and respect points are all on there and so ok so let&amp;#39;s have a look at the first this is ionic bonding okay so only bonding or basically where you&amp;#39;ve got charged ions and they&amp;#39;re held together by these very strong electrostatic attractions these are oppositely charged ions and basically that&amp;#39;s all an ionic bond is and it works in a really simple way really and so to get a full shell of electrons in this case we&amp;#39;ve got sodium and chlorine as you can see here it is your surgeon coins get a full shell of electrons thus owed you will have to give up an electron to chlorine so it gets a full shell sodium has an empty shell or I&amp;#39;ll have a full one underneath this shell here and the becomes stable by it being attracted to each other by forming an oppositely by forming an electrostatic attraction between the oppositely charged ions so let&amp;#39;s have a look there&amp;#39;s the electron and the you can see it jumps onto the coin and we form these ions I notice we draw box as well and square boxes and we put a plus and the negative charge around them to show that they are oppositely charged okay and these ions obviously depend on the group or depend on what type of iron that you form so if you look in Group 1 group one&amp;#39;s all form plus one is group to form 2 plus 3s on three pluses Falls very very rarely really form ionic bonds they&amp;#39;re covalent with each other so we don&amp;#39;t bother with them group 5 right these form 3 minus ions remember it&amp;#39;s easy for these to gain the 3 electrons then it is to get rid of 5 sets 1 upon 3 minuses group 6 form two minuses with 7 4 1 minuses okay you need to know if you these molecular ions as well so these are a little bit different so hydroxide is o h minus nitrates is no.3 minus and ammonium is NH 4 plus sulfate is so4 2 minus and carbonate is co 3 2 - make sure you know these you&amp;#39;ve got to these are so so important because you&amp;#39;ll see them a lot in chemistry obviously and so admit make sure you know the charges of them is one positive on there which is ammonium for some a negative ok right so let&amp;#39;s have a look a little bit more at some ionic bonding I&amp;#39;m not sure what we can do is you can work out the formula of an ionic compound and we can use this method called the swap and drop method I&amp;#39;m going to show you two examples here so the first thing you have to do is write down the two ions in this case we&amp;#39;re going to look at calcium ion and the nitrate ion as you can see there what we then do swap the charges over so the two now moves over to the nitrate and the - moves over to the calcium then drop the charges that includes the plus and minus bit and we drop them so that they are in lower case so the subscripts so that&amp;#39;s what happens there notice the minus is obviously gone from calcium and then we just simplified the whole number in this case this is ca no.3 - so it&amp;#39;s calcium nitrate and that&amp;#39;s basically how we work them out so that are slightly different want to stick with calcium and this time we&amp;#39;re gonna use oxygen again swap the charges over to - + - + drop the charges CA 2 + o2 and then simplify so we&amp;#39;re left with CA 202 but we can simplify that to just see 800 calcium oxide so you can do this at any type of ion it&amp;#39;s dead easy to see okay really quick and easy okay I&amp;#39;m only compounds for example sodium car they have giant ionic structures so they form things like this here here&amp;#39;s an example sodium Clyde the yellow circles or spheres are so divine chloride ions you can see they pack together in a regular structure very cubic shape giant repeating pattern you&amp;#39;ve got to know and the description of the structure of these things is pretty important these things dissolved in water pretty well and they water&amp;#39;s polar it can be attract the positive and negative ions obviously the Delta negative oxygen and water is a trap to the positive sodium ion and the Delta positive hydrogen&amp;#39;s are attracted to the chloride ions and so water can break these structures up they can also conduct electricity when they&amp;#39;re molten or dissolved in solution because the ions are free to move around when they&amp;#39;re solid like this they can&amp;#39;t because the ions are not free to move around as you can see they&amp;#39;re fixed to the cubic shape and they have very high melting points because we have lots of strong electrostatic attractions and these basically require lots of energy to overcome them so and this is why sodium Clyde and other ionic compounds have such high melting points very important to include things like strong electrostatic forces and lots of energy needed to overcome them so make sure you get all these keywords and your answers when you write in them ok Corinna balling is another type of bonding right this is a little bit more trickier actually even though you might think it might look simple and basically career pollens are sharing about two electrons okay very important and we basically do this very similar to ionic compounds to get a full shell of electrons except this is sharing electrons we&amp;#39;re not giving up or receiving electrons these are the sharing of them and there is still an electrostatic attraction there&amp;#39;s that word again this time it&amp;#39;s between the shared electrons and the positive nucleus in the middle that&amp;#39;s where the attraction is between these so that&amp;#39;s what&amp;#39;s going to hold these things together we can get single double and triple bonds and basically you can see there this is a single that&amp;#39;s a double and that&amp;#39;s triple so we&amp;#39;re sharing three electrons each and and the covalent bonds can also be represented by lines and you might have seen these in displayed formula okay dative covalent bonds on the other hand are a little bit different so and these are also known as coordinate bonds and this is where an ass donates two electrons or pair of electrons to another atom in this case we&amp;#39;re going to look at ammonia with a lone pair of electrons and the hydrogen ion H+ with no electrons in this can&amp;#39;t form a covalent bond because now so you having the electrons to share so if it wants to bond with another atom both the electrons are bit you can have to come from the other atom which is a bit unfortunate so these are called date of credit and coordinate bonds and if you see on here there&amp;#39;s the hydrogen look it&amp;#39;s now sharing this electron sharing the electrons with nitrogen but both of them have come from the nitrogen to the H+ ions make sure you same where they come from where they go to so there it is there and it can be represented with an arrow and the owl shows the direction of where the electrons have been donated from and where they&amp;#39;re going to so make sure you get the arrow the right way around okay and the other type giant covalent so they were all simple molecular very simple small ones giant covalent are a lot bigger and these include things like graphite and diamond so that&amp;#39;s look at graphite first graphite is basically made up of hexagons and each carbon is bonded three times and the fourth electrons delocalized you can see here look if you look at a carbon one two three times delocalized electrons are being delocalized so lots of strong covalent bonds means it&amp;#39;s got a very high melting point members really loads of energy to break these bonds so else can take quite a bit of energy but the layers graphite is very unique in its structures made up of layers there&amp;#39;s weak forces between these layers and they can slide over each other relatively easily and then do localized electrons allow you to conduct electricity they can carry a charge very important that you say it carries a charge they can conduct electricity which is unusual for and obviously nonmetals they carbon the layers are really far apart in comparison to a covalent bond length and so this means graphite is pretty low density compared to other giant covalent structures okay so it&amp;#39;s insoluble doesn&amp;#39;t dissolve and the bonds are far too strong and you can&amp;#39;t break them when you velcome to water so you can you can basically put a pencil in water they will dissolve thankfully just in case you want to do that home okay diamonds diamonds another one diamonds a bit different though it has each carbon has bonded four times instead of graphite where it&amp;#39;s three times and now because this doesn&amp;#39;t have and and because this doesn&amp;#39;t have the gap like graphite does and it means that they are they can conduct heat pretty well because they&amp;#39;re really tightly packed together and diamonds on see and are form from volcanoes so obviously they will be able to conduct and absorb some of this heat and so unlike graphite diamonds can be cut to make gemstones so you can cut them and graduate carcasses just prior to other breaks in their into layers and it&amp;#39;s got a really high melting point bit like graphite and loads of strong covalent bonds you need a lot of energy to overcome them it doesn&amp;#39;t conduct electricity very well though and in fact it doesn&amp;#39;t conduct electricity and at all really didn&amp;#39;t have any free electrons and like graphite does so and it&amp;#39;s non conductive electricity and again just like graphite it&amp;#39;s insoluble the bonds are far too strong for water to break them apart okay so make sure you know these examples of giant covalence okay shapes of molecules this could be a bit tricky this and there&amp;#39;s a very kind of certain rule in which you can apply you might have seen other rules as well but I think this one&amp;#39;s relatively straightforward to use I hope you agree so we&amp;#39;re going to use the number of bond pairs and lone pairs of electrons to work out the shape of a molecule so let&amp;#39;s have a look at this molecule here they&amp;#39;re very specific shape and this is because the bonds repel each other equally and they trying get is far away from each other as possible we&amp;#39;ve got electrons in these bonds they&amp;#39;re like charges so they will repel each it so this means they have a very specific shape so if we have a lone pair at next to a bond pair though these repel more than two bond pairs together and two lone pairs repel even further so that&amp;#39;s one look at an example here&amp;#39;s one here this is one with one lone pair now if you look at the angle of this one with no lone pairs of electrons and this one with a lone pair look the angle has decreased this is pushing these bonds closer together because the low become even further and if we have two lone pairs like this one the bond angle shrinks even further and so you can see here so what they&amp;#39;re doing is they change the shape of the bond angles and lone pairs what they do is they push them closer together and January is a rule not all the climbed over for things like tetrahedral structures which we&amp;#39;ll look at in a minute every time you have a long pair you reduce the remaining bond angle by two and a half degrees so and we have got to be careful for some other molecules where this doesn&amp;#39;t happen because they cancel out but I&amp;#39;ll point it out later on okay so let&amp;#39;s have a look at some of these so we&amp;#39;re going to use the bond pair and lone pair electrons to work out the shape of a molecule so the first thing you should do is really drop your dot cross compound a dot cross model here basically just helps you to work out how many bond pairs and how many lone pairs you&amp;#39;ve got if you&amp;#39;ve got an ionic molecule which they sometimes give you all you do is and you add electrons to the central atom for a negative I&amp;#39;m so and this is if your molecule has negative ID and remove them if it&amp;#39;s a positive ID so for example ammonium nh4 plus the nitrogen would have four electrons normally nitrogen has five but we take one away to take through calculus so it has four and this is only of course this is not what happens in reality we are only using this method to work out the bond angle and the shape okay this is not what happens in reality it&amp;#39;s just a method of working out so and so yes so let&amp;#39;s have a look at this one so the total in this case if we add them all up tells you the shape in this case it is a tetrahedral got four bond pairs there are no lone pairs left on the central carbon so it&amp;#39;s a total of four and so if you and have lone pairs you need to replace the bonds for lone pairs and change the shape and the bond angle which we&amp;#39;ll look at in the middle a got special needs so something like this yeah okay and you can see water obviously this has to bond pairs two lone pairs and but crucially it&amp;#39;s based on a tetrahedral that&amp;#39;s the very important thing is interceptor you draw it&amp;#39;s based on that and all we do is reduce the bond angle by two lots of two and a half words five degrees okay so let&amp;#39;s look at the shapes with no lone pairs first okay so you&amp;#39;ve got to know the names of these shapes and you&amp;#39;ve got to be able to know the angles of them and recognize them etc so we use the bond pairs and lone pairs of electrons to work at the shape of a molecule so let&amp;#39;s look at the first one this one&amp;#39;s got to bond pairs no lone pairs here&amp;#39;s an example and which is beryllium dichloride the name of the shape is linear you can see it&amp;#39;s 180 degrees here and we have two atoms that are obviously and joins either side no lone pairs okay let&amp;#39;s look at the other one bf3 bond pairs three no lone pairs so if you have this scenario we have three bond pairs no lone pairs we have a trigonal planar this is a flat molecule 120 degrees either side okay for bond pairs and no lone pairs these molecules are called tetrahedral it&amp;#39;s the first time when we actually have a 3d shape and so here&amp;#39;s an example of a tetrahedral we have one bond one atom coming towards you one in your plane of vision that one&amp;#39;s in your plane of vision and that one&amp;#39;s going away from me the bond angle is 109 point five four all tetrahedral okay if you have five bond pairs and again all of these got no lone pairs the shape we are looking at is trigonal bi-pyramidal so you can see obviously this is an example is PCL 5 but here&amp;#39;s an example here so you see we&amp;#39;ve got three here there&amp;#39;s the threes you&amp;#39;ve got one two three okay in the plane and then we have to top and bottom as an axis so this bit is planar and it&amp;#39;s triangle shapes and call it trigonal planar but this is trigonal bi-pyramidal and so it&amp;#39;s basically if you look at the trigonal planar one here that&amp;#39;s basically that but just tipped up on its side you can see but it&amp;#39;s by pyramidal because actually if we draw a line going from there to there there to there and there to there we form a pyramid on the top we form a pyramid on the bottom results as well quite a trigonal bi-pyramidal bond angles 120 because that&amp;#39;s basically just that tips on the side but the bond angle between the top the polar ones here and the trigonal plane a bit in the middle is 90 degree so this one&amp;#39;s got two angles so make sure you know the difference okay final one is one with six bond pairs and no lone pairs and this is an example of an octahedral structure it&amp;#39;s very similar to trigonal bi-pyramidal except in the middle as you can see here we have a square shape you can see the square shape here in the middle two axes top and bottom the bond angle here is nice degrees we call octahedral okay so let&amp;#39;s look at them now with em with lone pairs so there&amp;#39;s not as many of these but we&amp;#39;re based on and remember on the original shape so we&amp;#39;re going to use the bond pairs and lone pairs again to work this out so here&amp;#39;s an example this one&amp;#39;s got three bond pairs one lone pair so this is the example of pyramidal an example is ammonia there we go looks a little bit like tetrahedral so it got a lone pair there if we go for the two bond pairs and two lone pairs we get a bent molecule or nonlinear it&amp;#39;s also known as so we&amp;#39;ve got the lone pairs there on the side again look we&amp;#39;ve shrunk the bond angle from 107 to 104 point five so that&amp;#39;s been shrunk by two and a half degrees okay three bond pairs and two lone pairs an example of clf3 trigonal planar now this one 120 again we&amp;#39;ve taken the bonds from the top and the bottom and we get this flat trigonal planar structure two lone pairs notice the bond angle here is still the same as what it was before these lone pairs are repelling each other equally this one repels these bonds down this one repels and back up again so they cancel out and this is exactly the same four and this one with four bond pairs and two lone pairs based on octahedral and you can see here the bond angles 90 degrees lone pairs are pushing equally and so that&amp;#39;s quite important so the bond angle actually remains unchanged the two lone pairs repel equally from opposite sides make sure you know the names of these shapes okay electronegativity so we need to remember this definition electronegativity is the ability for an atom to attract electrons towards itself in a covalent bond and basically the further up and right you go in this and in this here this periodic table the more electronegative the element is excluding the noble gases this is one would black them out so fluorine is the most electronegative element and we can use the another powering scale which helps us to quantify how electronegative something is there&amp;#39;s Florina fluids the most electronegative for B the biggest number here and but as we go further away the electronegativity drops basically the bigger the difference in electronegativity the more polar the bond is so just watch out for this scale in your exam okay so let&amp;#39;s have a look at polar bonds so covalent bonds can become polar the atoms attached to it of a difference in electronegativity so the bigger the difference in electronegativity the more polar will be so looking at this example here we&amp;#39;ve got h and cl bonded together the chlorines the most electronegative is pulling the electrons towards the chlorine more than the hydrogen we give it this little Delta negative symbol to show to show that it&amp;#39;s pulling the electrons towards it hydrogens got the dells positive okay here&amp;#39;s another molecule this atom is basically of this molecule sorry is a molecule of chlorine they&amp;#39;re both just as electronegative as each other so for that reason this is not polar because the electrons sit bang in the middle of the bond and so these things are not polar hydrocarbons are the same they&amp;#39;re classed as non polar as well so like alkanes etc and alkenes so these things and the electrons are just shared equally within that bond sometimes uneven distributions well they do actually uneven distributions of charge leads to polar molecule so water is a classic sign of that see the electrons are now distributed unevenly they&amp;#39;re not spread across the charge and spread equally across the molecule so water is an example of an uneven distribution however if you look at ones where we&amp;#39;ve got a bit of symmetry like here this is a completely symmetrical molecule and so this one are actually nonpolar even though they&amp;#39;ve got polar bonds the electrons have been pushed either side equally so we call this a symmetrical molecule and so these actually have no overall priority she could be really careful for that and what&amp;#39;s out of that okay intermolecular forces these are really really important okay the first one our van der Val&amp;#39;s also known as induced dipole dipole these exist between atoms and molecules but not total bonds here what weak forces between molecules and you need to know these fit into like a little kind of league table of other intermolecular forces and Vander vials are the weakest out of Bolin and they exist in any molecule that basically has electrons in them and so basically evander valve is where a dipole can be created when they move near another atom or molecule so let&amp;#39;s have a look at this example here so you can see that we&amp;#39;ve got an atom here and the electrons in this molecule are not normally unevenly distributed but when they move near another molecule they are you see the electrons have been moved over to one side of the molecule we get this Delta negative and Delta positive on there this dipole is only temporary it&amp;#39;s only there when it&amp;#39;s next to another molecule that&amp;#39;s nearby and obviously that molecule will have electrons as well and basically when they move away the interaction is destroyed so you can see here when this molecule is near another molecule we have a delta negative and a delta positive as soon as this molecule leaves the electrons will shuffle back on to the other side and even itself back out again but for the time that they&amp;#39;re close by there is this weak interaction Delta negative Delta positive that&amp;#39;s really called an induced dipole dipole it&amp;#39;s only brought about when it&amp;#39;s near end of the molecule okay we get this force of attraction um okay so let&amp;#39;s see if we can apply this to something specifically so for example iodine iodine is a classic example van der Waals forces can hold some of these molecules and crystal structures so iodine have this type of crystal structure it&amp;#39;s pretty pretty and it&amp;#39;s a it&amp;#39;s solid at room temperature I D and the key thing is we have weak van der Waals forces holds the iodine molecules together remember it&amp;#39;s I 2 but we have strong coven of bonds that hold the two iodine atoms together don&amp;#39;t get these two confused okay the bigger the molecule the atom the more Van der Waals forces you have so that&amp;#39;s what you&amp;#39;re looking out for because they have larger electron clouds so when we boil the liquid and it could be water for example we&amp;#39;re breaking the weak van der Waals forces and not the covalent bonds actually water has hydrogen bonding as well so I tried break them but when we boil a liquid and that&amp;#39;s what happens so for example bromine we&amp;#39;re just breaking these weak forces not the bonds and never mention the bonds when we talk about these small molecules we&amp;#39;ve gotta have enough energy to overcome these forces and hydrocarbons remember we said we don&amp;#39;t have this polarity so they have bound the valves and basically the longest straight chain hydrocarbons have more van der valve than air forces and energy and then you&amp;#39;re kind of non straight or branched chain and shorter ones so basically these need more energies needed to overcome these forces and this increases the boiling point of these molecules and like I say the branched ones these are hydrocarbons with bits sticking out of them they can&amp;#39;t pack together as much so this means that the van der Waals force is a lot weaker and the and the chain if the chain is shorter as well this lowers the boiling point as well so short the hydrocarbons all hydrocarbons a lot of branching have lower boiling points okay well it&amp;#39;s look at the next in specular force these are called dipole-dipole they&amp;#39;re a little bit stronger than around the valve these exist when you have a molecule as a permanent dipole so basically the dipole exists irrespective whether they&amp;#39;re near another molecule or not so things like where you have a polarity so permanent dipoles you&amp;#39;re looking for polarity HCl is an example of a permanent dipole and you have this weak electrostatic force between a delta negative chlorine the electronegative element the Delta positive hydrogen atom in this case and so the Delta negative is attracted to the Delta positive as we say so in Lake van der Waals forces dipole-dipole interactions involves molecules for the permanent dipole they are stronger than van der Waals it is important to note though that even though this does have this molecule in particular has a permanent dipole it also has banned the valves as well the strongest in specular force is the permanent dipole dipole and we can kind of prove this as well we can prove that polar molecules like water exist and if we take a charged rod and we put it near a steady stream of a polar liquid you can do this in a burette what is an example of polar liquid and then what happened is you should see the liquid bed towards the rod and because obviously the charge difference here and basically the molecule will align the line itself to the oppositely charged rod so you can see here this is the positively charged rod in this example so the Delta negative oxygen will be attracted towards a positively charged rod so it&amp;#39;s a very important little test for polar molecules okay and the lust intermolecular forces hydrogen bonding this is the strongest of the lot the hydrogen bonding is quite unique is that it&amp;#39;s basically a type of dipole-dipole force and you basically get hydrogen bonding when you use very electronegative element so your most electronegative elements are nitrogen oxygen and fluorine it&amp;#39;s these three and hydrogen so any molecules that contain nitrogen oxygen or fluorine and a hydrogen will get involved with hydrogen bonding and it&amp;#39;s basically the lone pair of electrons on these atoms that get involved that plays a very important role so here&amp;#39;s an example of showing hydrogen bonding we&amp;#39;ve got all our partial charges on this or Delta negatives Delta positives there&amp;#39;s the lone pair on the oxygen the interaction is between the lone pair on the electronegative element and the hydrogen and this is a hydrogen bond very common form to ask this and you gum so make sure you know how to do it and again just like with the previous one molecules that have hydrogen bonding also have van der valves and they have dipole-dipole they have all three it&amp;#39;s just the strongest one is hydrogen bonding so make sure you&amp;#39;re very vigilant exam and you watch out for them um so let&amp;#39;s have a look a little bit further ice obviously is just frozen water and it forms this very regular structure held together by hydrogen bonds but ice is strange because when you cool normally many cool things down they get smaller they contract but ice actually expands it gets bigger and the reason why is because we have hydrogen bonding that&amp;#39;s pushing the molecules further apart and this makes ice less dense so if an ice cube in water and it floats and let&amp;#39;s look at the boiling points and properties of these things as well so if you look at the h-f h-f has a higher boiling point than hate see L remember HF is one of the molecules that can hydrogen bond between other molecules because hydrogen bonding is stronger it needs more energy to overcome these forces looking at the other ones things like HCl HCl can&amp;#39;t hydrogen bond it drops massively less energies needed to break these however it starts to pick up again when we go for things like HBR and H I and this is because and yes they have dipole-dipole forces all of these do but the biggest effect have the biggest effect that&amp;#39;s having effect on the boiling point is basically the band of ours is increasing the out and getting bigger the molecules getting bigger there&amp;#39;s more electrons Barneveld forces increases so therefore the boiling point increases from heat CL thi okay metallic bonding so metallic bonding and these are giant metallic lattice structures so these are quite large structures they have a very unique setup you have positive net lines which is the green circles there they&amp;#39;ve formed when metals donate electrons into the sea of delocalized electrons and there&amp;#39;s basic electrostatic attraction between the positive metal lines and then delocalized electrons and basically the more electrons and assam can donate the higher the melting point and so for example magnesium has a high melting point of sodium it can donate two electrons into the delocalized cloud where sodium can only donate one so therefore the forces of attraction between the positive and negative charges are a lot weaker metals are good thermal conductors as you&amp;#39;d expect and they have delocalized electrons so they can transfer this kinetic energy very easily remember when you heat things the the particles move and electrons move it&amp;#39;s the electrons that allow the conduction of heat there are also good electrical conductors because they can carry a charge or current they can move that through the lattice there are free electrons delocalized electrons okay they have high melting points again strong electrostatic attractions between the delocalized electrons and the positive metal line and solid metals are insoluble then the metallic bond is far too strong to break and don&amp;#39;t bleed and otherwise you&amp;#39;ll get metal structures dissolving when it rains so that&amp;#39;s pretty useful okay so quickly this particle model F which straightforward this bit really and solids Titan packed regular arrangement really high density you can see the arrangement of these here they vibrate on the spot and they can&amp;#39;t be compressed liquids they are tightly packed again they don&amp;#39;t have many gaps between them random arrangement high density they move around freely and they slide over each other and they really difficult to compress because they&amp;#39;re tightly packed and the particles have a little bit more energy in them in liquid than you do in solids and that&amp;#39;s why they&amp;#39;re moving randomly gases well they&amp;#39;re very spaced out random arrangement and very low density particles able to move around freely and they have a lot more energy than liquids and and solids obviously because they&amp;#39;re moving out freely so don&amp;#39;t really comment on them that starts the bathing and finally just summarizing the bond types and make sure you can summarize these all really it&amp;#39;s a good way of just making sure that you know all these things giant covalent graphite diamond silicon dioxide for example normal temperature their solids they don&amp;#39;t conduct apart from graphite that don&amp;#39;t conduct electricity as liquids very difficult they normally sublime anyway they don&amp;#39;t really melt soluble in water high melting points give you lots of energy to break them simple moleculars these are covalent as well liquids or gases normally won&amp;#39;t conduct electricity and they might be soluble depending on the polarity the low breaking low melting and boiling points and because you&amp;#39;re breaking weak forces giant ionic much much higher melting points and generally soluble more so because the polar they will conduct if they&amp;#39;re dissolved in water or molten and normally they&amp;#39;re solids very high melting points strong electrostatic forces and metallic it&amp;#39;s pretty much the same except they will conduct solid and liquid conduct electricity and which is unusual and like the rest of them and a usual temperature normal state is solid again high melting pipes because they have strong electrostatic forces and so the polarity you&amp;#39;ve got to be watch out for this polarity zapala molecules they do dissolve well in polar solvents like water and but nonpolar molecules don&amp;#39;t hydrocarbons you need nonpolar solvents really for them to dissolve so there we go that is the end of the on the video I thought I hope that was a useful summary of bonding is quite a bit their bond angles are pretty important and just again if you want a copy of this of this PowerPoint be good for revision and just look in the the description box for this video and click on the link there and you get a hold it in there that&amp;#39;s it bye bye

Transcript for:Bonding Revision (AQA)

Transcript for:
Bonding Revision (AQA)