now our group four elements right just what does the US say about group four um somebody yeah we're going to go through parles yeah so group four where is it in syllables here's it okay we're now going to have to really explain the variations and physical properties of the elements in terms of structure and bonding we're going to have to look at the bonding of our group four tetrachlorides as well as reactions of tetrachlorides with water you must all Al discuss generally the trends in bonding acid Bas acid based character right thermal stability of our oxides and um or oxides of our oxidation state 2 and four plus right then discuss the relative stabilities yeah then we're going to discuss the relative stabilities of the oxides and equal cations of elements in the their higher and lower oxidation States and then we must discuss the use of ceramics right based on Silicon for oxide so the fact that silicon for oxide is used in creation of furnace Linings forine um Ceramics right and glass right so the chemistry of the elements that's really what it what we must look at so our group four elements are are placed within our P block our group 14 elements and our group 14 elements will consist of our carbon silicon germanium tin and lead right and here is generally the trend right um of these elements right so we must note that carbon is a nonmetal we know that carbon is a nonmetal and carbon is so interesting in which its hybridization and Tetra veency allows for a specific of chemistry which is organic chemistry go ahead somebody you ask a question I don't here questions coming from okay so carbon is so interesting in which would have an entire field of chemistry based on only one element right but when we look at this now silicon right has a similar characteristics to carbon as well because it's right there right beneath right carbon in the group itself right but carbon is generally the only nonmetal there and silicon to gerum right 10 they are metalloids right as we look at the band theorem if anybody the solid state physics you know physics unit to right and then we'll look at tin and copper who well tin and lead rather which are just complete Metals right so we have the outer shell configuration for these elements and we know that the noble gas that come before these elements would be what which noble gas configuration comes before this which noble gas comes for carbon repeat helium lovely right and then for the Silicon for germanium it would be argon argon lovely argon and then Krypton going down all right so notice you know remember in chemistry in fourth form chemistry where chemistry teacher tell you that groups are defined right by the fact that the elements in the groups have the same number of valence electrons look at that there same electronic configuration generally going down the group so this is why we categorize them as a group right and this this electronic configuration lends to the reactivity of these compounds therefore or of their compounds rather right therefore right that's why they would have General Trends and similar um Trends overall and they're plac within the same group right so we can note that the atomic radius tends to increase as we go down the periodic table can anybody explain why okay we're going to go go into why that right so just think about why that is right we also have the first ionization energy increases as we go up the group right the structure becomes more metallic metallic characteristics increase as we go down the group the density fluctuates right but it generally increases going down the group action right so carbon will exist there in it's many type of allr carbons of carbon has over 50 going upwards allotropes but we only look at generally three but carbon has many allotropes such as B Min of ferine right and different types of um um materials right um such as graphine graphite diamond we would have looked at those in CC right um melting point will increase right what's Happening Here the melting point I have a feeling that this melting point is in the wrong direction all right we're going to fix that as we go through this um PDF but the conductivity tends to increase as we go down the group as we move from nonmetal to metaloid to metal all right so for the atomic radius right we know that generally the atomic radius are Atomic radi increase going down groups right because we jump periods right as one as we descend the group full orbitals are added to successive elements and we should understand that we should know that we should internalize that right so if we remember how electrons are added to electron energy levels right electronic energy levels right are in orbitals once we fill the p orbital we go on to the S to another s orbital right right and then another p orbital then a d orbital so what we need to note is that as we fill a orbital right we start to fill right another orbital that is way um that that's distance is way further from the nucleus right so what happening here is that as we go go down let's say that this is carbon right once it fills this orbital exemplified by the circle here in the corner of the screen once it fills that orbital right then you're going to gain another orbital so this could be let's say silicon right and then this would be germanium and then this would be thin so as we increase um as we descend the group or increasing the period and we know that the period would have different numbers of orbitals and they will in that will increase gener as we go down the group so atomic radius increases as we go down the group specifically right so out electrons are placed further and further away from the nucleus causing an increase in the size of the atom right and we should understand that now in relation to ionization energy now it will decrease going down the group why is that electrons are found further and further away from the nucleus the further an electron is away from the nucleus we can note the fact factors that affect the iation energy from there can anybody tell me the factors that affect iation energy I already stated distance shielding effect size of the nuclear charge okay the size of nuclear charge okay of the size of the nuclear charge H shielding the screening effect okay shielding screening same thing different shape the shielding effects anything else well I think this is all that is at the the orbital well the orbital penet okay orbital sub penetration orbital penetration repeat something about the atom repeat that the size of the atom the size of the atom okay the size of the atom will speak to the distance right there from the electrons are from nucleus but we could just add that the size of the atom okay all right so we would know that these things affect them right so what's happening you know is as we in as we go down the group we're increasing the period right so what we're noticing is that nuclear charge does affect right the ionization energy yes right but when nuclear charge how nuclear charge actually has a better effect right an energy is across a period because as we go across a period we know that the size of nuclear charge will increase right and the electrons right themselves will be added to the same energy level right to the same energy level therefore the size of nuclear charge would be more significant than the distance of the electron because once we go across period we are increasing in nuclear charge but the distance of electrons are more or less the same they it doesn't change right so as we go across a period the anation energy will increase but when we incre when we go to the next period now we're going to have a severe decrease in ionization energy why because the ionic radius will increase as we change periods as we go down the group we increase the period we increase the distance right of the electrons of the outer electrons right from the nucleus right plus those other electrons that are between the outer electrons and the nucleus will contribute to shielding right so we know that if we have a nucleus here right and we have an electron here electron here right just looking at electrons like this and we have electron here these electrons will actually help to Shield the nucleus so that's what we really speak of as shielding right so there's going to be some negative charge that exist between the nucleus the posi charge nucleus and the negatively charge electron all right hopefully we know that and remember that all right so that's generally what's happening there we can see that there's some shielding occurring there so once we go down the group we increase the shielding and what we're doing is increasing the distance of the veence electrons from the nucleus and that will decrease energy because the further away an electron is from the nucleus right the the smaller the amount of energy required to remove it what is ionization energy somebody Define the term define it please sir first ionization energy is the amount of energy you need to remove one mole of electrons from gas ion to form one Mo of gas C all right so remove the energy required to remove one Mo of electrons from one Mo of gas atoms to produce one of gas catons right and then the first first and second that would just speak to the order of electrons that you remove did you remove the first electron alone and after you remove the first electron you can remove the second electron and those type of stuff right so we know that so as we increase the atomic radius it would be successively easier for the ionization for the ionization energy to be reached right so ionization energy will decrease as we go down the GL so there's a large difference in ionization energy between carbon and silicon after the different after that the difference is relatively small right right so this is a consequence of the D Block element so once we actually transition into d block elements right the jumping ionization energy is actually much where it's much larger significant the D orbitals do not screen the nucleus as effectively so that's where we get into orbital penetration the S orbital penetrates um through shielding effects more so than the p orbital and the p orbital penetrates through shelding effects more so than the D orbital so what's happening when you reach to D orbital electrons they do not feel the negative charge a nuclear not not negative charge a nuclear charge they do not feel the positive charge of the nucleus right as greatly as S and P electrons do so once we get to the D orbital now the D Block elements now their ionization energy will be much smaller because their veence electrons are on D orbitals and the orbitals do not feel the effective nuclear charge as best right as the S and P orbitals we must note these differences once you understand the concept you can apply to any element on the periodic table any questions or concerns before move on structure and bonding moving on for structure bonding now however give me a second let me check something for structure and bonding right we would note that carbon is the only nonmetal in this group go ahead go ahead christe go ahead sir somebody ask if it could explain the last part again yeah okay so we have something called or penetration right I do not want to overload on too much information right that you guys will not need right but what orbital penetration speaks to is how much nuclear charge right do the electrons in this orbital feel how effective is nuclear charge at at penetrating these orbitals meaning when I'm in a s orbital I feel the strength of the nucleus much more than if I was in a p orbital so p orbital electron Fe the strength of nucleus way less than S orbital right so remember you know the S orbitals are something like this around the nucleos sperical right the P orbitals are going to have our dumbbell shape right but when we look at D orbitals D orbitals exist outside of this right the orbitals exist outside of this in much more complex ways right so the orbitals are actually found further away from the nucleus in some in some cases right but let me not go too much into that and just state that s orbital electrons feel positive charge much greater than the p orbital electrons which FS that charge much greater than the D orbital electrons make sense and that is explaining why when we jump why there's a huge difference in ionization energy from carbon to Silicon lovely now for structure bonding we did say that carbon is the only nonmetal in the group as we go down we enter into the well we note the fact that silicon and germanium are actually metalloids right and then the tin and lead are actually Metals so when we look at this now the physical properties differ because the structure of the elements change right so we know the giant molecular structures right two giant metallic structures all right so we would note here carbon carbon can form giant molecular structures like diamond and graphite same thing with silicon when we create silicon semiconductors and stuff like that they're molecular structures giant molecular structures we can use silicon and germanium to create those um super structures right so these giant molecular structures right are shown by both carbon right carbon would have that um giant molecular structure generally right and then silicon and germanium will have giant molecular structures that are more metallic than that of carbon right so it's similar to how Diamond would be arranged for the Silicon um structure while the tin and lead have metallic structure so metallic lates basically islands of positive charge the cations surrounded by a sea of the localized electrons that would be the structure of tin and lead right so as the atom get larger right the attraction of the nucleus for the electrons actually increases right and it the attraction of the nucleus of electrons in Co valent bonds get weaker right therefore electrons become more delocalized so that's what's happening so as we have as we gain um as we go down the group our positive charge our nuclear charge actually increases right and causes electrons to actually be held more closely right and allow for covalent bonds to be no longer feasible so what tends to happen here right is that the electrons CS right are going to be more delocalized in the structure they are not they are not going to want to actually form calent bonds right so since they deize they tend to exist around the nuclei in a metallic structure rather than form neatly organized bonds like the giant molecular structures of germanium silicon and carbon Bond right so as we go down the group we increase in our metallic characteristics in which our electrons will become more delocalized because of the large strength of the nuclear charge right so we have that there so the co valent characteristic um quickly turns to metallic characteristics within this group as we descend the group all right so we're looking at this fruit now we have the melting points so silicon and gerum have giant coent structure similar to Diamond right so they have a very high melting point right so because of a lot of energies because a lot of energies need to break the calent bond so that is why we have a high melting point right so if we look at a melting point of diamond that is 3,730 De very high melting point when you look at silicon now right is a melting point of around 1,00 410 then geranium of a mum point of our own um 9 37° right so the melting point is actually decreasing as you go down the group there but with tin and lead respectively right they have lower melting points compared to most other metals so tin and lead tends to be tend to have lower melting points compared to rest of the metals right the strength of the metallic bonding decreases as the size of the ions increase so as we go from left to right across periodic table right our Metals would actually start losing their metallic characteristics as we go across all right so tin and lead are have fairly large ions right so they have relatively low melting points right so we look at the melting point of tin which is just 232° C and Lead which is around 237° C right really really small I want us to understand what's happening here now do we want to have a look at a period table to see what I'm talking about here with a peri table help okay let me stop let me stop sharing and let's get that per table up so let's share all right so I want us to understand all it is is periodic trends that's all we're looking at periodic trends right looking at the periodic table we should be able to see these different Trends and understand what's happening all right so just give me a second I'll be able to reare soon the presentation ends all right okay so let's let me share my entire screen all right so we can have a look at this now okay hold on me one moment all right so let me just bring up P table I should be able to look at that all right so what we're talking about is a metallic characteristic of our lead and our tin right within our group four so lead and Tin right they would have lower melting points right and less metallic characteristic as the rest of our Metals right and I want to look at the oh it's loading it's loading it's not there for me either it's still loading right that's because I'm recording are you presenting something right now yes I am okay you see that there all right so with this now right we can see that we're talking about here our group 14 right so this would be our carbon our silicon our germanium tin and our lead we don't go to fium or anything like that right at this level right we don't we do not speak about the rest of these elements over here all right so we're looking at Carbon we know that carbon here is a nonmetal we look at silicon and ger ium to be metaloid so looking at our metalloids we know that so our nonmetals over here carbon in group 14 is nonmetal and our metalloids would be silicon and germanium and then our post transition metals down here would be our tin and our lead all right but what we're saying is that as we go across the periodic table our metallic characteristic decreases so notice know right if you look at period two right right we have the metal up here right these two are metals right lithium and are metals Boron is a metaloid and then everything else is a nonmetal right so notice how metallic characteristic decrease as we go across the periodic table notice that on one side it's metallic on the other side it's non- metallic but remember that we stated before we said that that metallic characteristic increases as we go down the periodic table therefore if we look at period like period 3 and period four right mear characteristics continues into Group 13 and group 14 as we go across therefore if we look at something like a group 16 right even group 16 we have a metal present right so it's just that as we go across Peri table metallic characteristics decreases and as we go down it increases right so we have that there so what we're doing is focusing on carbon silicon germanium tin and lead here right so we have that there so since and it's based based off of the increase in atomic size we should know let me go back to the credit table okay so we should know right that as we go across the periodic table atomic size increases right so the atomic radius would increase so if we look at something like a lithium lithium's radius is 167 pom burum okay I don't think I should have used this because I know that there's a difference in Trend as we go across period too all right but if you look at something like carbon right right we're looking at the at the size of the atoms right so if we look at the radius we can see that the radius is actually decreasing as we go across right radius in um decreases as we go across the periodic table right so large metal ions right um metal atoms like um calcium um well calcium is here right which is 194 pom right if we look at that our rest of our transition metals tend to have smaller sizes so it's important to note that the small all the ion things they get right they have relatively large sizes compar of the period but compared to elements um well other metals rather right they would actually be really really small right so what we're seeing is that tin and lead right have relatively large sizes right compared to the rest of their group so compared to carbon Etc right but as we compare them to Metals their sizes are very small so right and they have very large ions fairly large ions compar of the group right so when we look at this now the metallic the metallic um characteristics decrease and the melting points tend to be small right because some people are looking for the melting points right of tin right right and lead to be really high but they're actually really small in this group right tin can also exist as gray and white tin so tin also has alres right looking at these bonds now the decrease in the melting points reflects the weakness in the calent and metallic bonds right so if we look at this now the bond strength is equal to bond length right is proportional to bonding right so well inversely proportional to bonding so if we look at that the bond strength between right carbon to carbon silicon to Silicon tin to Tin is actually decreasing as we go down the right so carbon carbon carbon to carbon single bond has the greatest bond energy compared to the rest of them right so that's is um also something that we can look at to really see that our melting points actually decreasing as we go down okay and if you can note that group four elements can actually bond to each other they possess cation right patination is is the ability for the element to bond to itself okay to form intricate structures all right one thing that I normally see on passer question is the conductivity right of our substances within group four right we know that carbon does not conduct electricity unless it's in its allotropic form of graphite right and carbon when Bond into four other carbon atoms tend to have a tetrahedral Arrangement right so there are no no free electrons present Within These carbon compounds conduct electricity unless it is in graphite in which we have coent bonding right to three other carbon atoms and the fourth carbon atom is actually delocalized in the electron cloud so that's how we have the conductivity of graphite the fourth carbon electron is delocalized all right if we look at silicon and germanium they conduct electricity to a very small extent they are semiconductors right so our metalloids are all semiconductors they do not necessarily have delocalized electrons per se but some of their electrons can move in and out of positions right through their n and p Junctions physics students should understand this you have n and p Junctions with within semiconductors that allows positive and negative charge to change right and due to this movement of positive and energy positive and negative charge within NP Junctions right we can actually have electricity flowing so people doing physics unit 2 You' understand this to the T right I think some of some of the understanding is from physics physics C physics as well right so semiconductors right don't necessarily have deiz electrons right but they're they're structured in such a way in which electrons can actually move or charge can actually move right within their latices right so this so so this is it their conductivity actually increases with temperature as opposed to elements that um like Metals right that actually increase conduct that conductivity decreases with temperature right so with semiconductor when we increase the temperature we increase the conductivity right with General Metals that are just general conductors conductivity actually decrease when we increase the temperature all right so when you go into stuff like super conductors and stuff like that super conductor can conduct a lot of electricity at low temperatures you can look at all of that when you start to talk about electricity and magnetism in physics right so this is why the structures are called metalloids because they're not Metals per se they don't have the localized electrons per se right but right they have NP Junctions within their laes so there are points in their laes in electrons can jump out then jump back in so it's not going to be free and delocalized but there's some amount of movement within their laes and that's why they have such metallic characteristics right with tin and lead now we know these are complete electrical conductors because they are typical metals that have the the ca of the localized electrons all right that make sense so far guys does conductivity make sense cuz normal conductivity when they're asking questions go ahead or you just raising your hand to show that you okay make okay M all right that's fine hopefully it makes sense do you guys understand what's Happening Here the fact that conductivity increases going down the group okay lovely with our group four Tetra CS now we must understand how our tetrachlorides react so group four elements form tetrachlorides tetrachlorides are all simple coent molecules we have to look at this so this is shape of silicon T chloride here which is tetrahedral in nature all tetrachlorides are volatile liquids that we must know they have low melting points and boiling points and we must know that because the the elements the car the the chloride Bond right becomes longer as we go down the group and that is because the atomic radius of our elements get larger as we go down the group and that means that as we go down the group The there are weaker bonds right and ares become less stable what am I saying let me rephrase that carbon chloride right is like that car um silicon if we look at silicon let me rewrite that somebody unmuted they have a question doesn't seem like they have a question not sure if they're aware that they are all right but we're looking at this right so carbon would let's say that no we don't have to look at that let's go up let's go up see it here the bond length is increasing as we go down the group therefore the bond strength actually decreases as we go down the GL so if we look at those bonds with the chlorine now the bond length between our elements and chlorine will actually continue to increase until it's no longer feasible for that compound to form a tetrachloride right so that's what we're talking about the stability of tetrachlorides decrease Going Down group four right that's what we're talking about so we have to look at the reactions of water with our tetrachloride so the hydrolysis of the tetrachlorides of group four their activity of group four tetrachlorides with water all right so all tetrachlorides with the exception of carbon tetrachloride are readily hydrolized by water to precipitate the oxide of the plus4 oxidation state so when silicon tet reacts with water we get silicon dioxide which is which is silica the same thing that we find in sand or glass or anything like that plus our hydrochlor hydrogen chloride rather so hyen chloride would be in the form of a gas okay when we hydr germanium tetrachloride we're going to get germanium dioxide and hydrochloride hydrogen chloride rather right so basically the ease of hydis will actually increase as we go down the group and we know why because the bonds become weak as we go down that means it's going to be easier to hydrolyze as we go down the gr right carbon Tetra chloride doesn't really go through that right and that's basically chloroform carbon Tetra chloride all right this is inorganic chemistry so we can't refer to it generally through organic understanding right okay somebody's mic is on and I'm not sure if there arew that that is so all right so let's go now so silicon germanium tin and lead all have empty orbitals close enough to in energy right to the occupied p orbital these at atoms allow for incoming water molecules to donate a pair of electrons their D orbitals to form a bond right so basically water will act as a lion to these metals and metalloids right so for example our silicon chloride now in water can donate a lone here right so water molecule can donate it loone pair to the Silicon right cuz silicon has an empty D orbital or unfill D orbital rather it's not going to be empty but well it's going to have a unfill d orbital right so therefore the water can actually act as a Ligon and donate electrons right to the the Silicon itself right so that's where you can tend to create our tetrachlorides right so it can form ative equalent Bond silicon oxide silicon chlorides and those type of stuff right dative Co bonds tend to be a little bit weaker than our regular Co bonds all right hopefully everything is making sense so far right with our t chloromethane though we know it's imiss in water we should know that our organic solvents are actually non well organic solvents in this case are REM missible with water which just let me not go any further they're remiss in water right so the empty 3D orbitals in carbon are too different in energy for carbon to extend its octet right to form da Co mon water therefore it's just a regular um sharing coent bond that it has okay so this is basically all that we were stating up there right and now we can actually start looking at the acidity right really so let's look at the oxides of the group four elements so we have carbon dioxide silicon dioxide germanium dioxide tin dioxide and Lead dioxide all right and these are our plus4 oxidation States so if you look at the oxidation States they must all be plus4 we should know that right to find oxy States the boiling points they are the the boiling points down there right so the boiling points here and then lead on dioxide actually decompos on heating so it's too unstable there the structure now we know it's a simple molecular for carbon dioxide for and it's going to be a giant molecular structure for silicon dioxide we should know silicon dioxide we use it as an example for gular structure from C chemistry and then we have an intermediate between the giant molecular and ionic structures for our germanium towards our lead dioxide all right all right acid based nature now we know that carbon dioxide is acidic we know that silicon dioxide is acidic and all the rest of them are ambic so what we said that acidity decreases as we go down the group basicity increases as we go down the group and when we look at the thermal stability all of them are stable at high temperatures except for our lead um dioxide which decomposes is on the warming and that's why this boiling point doesn't really exist because when you reach a high temperature just it does it just decomposes it doesn't really boil right it just decomposes like that okay so it decomposes on warming to give us regular lead oxide at a Le plus plus one oxidation state and our oxygen gas everything is fine so far I know it is a lot of information why is it more stable in 2 plus than 4 plus state that is rather interesting those are the questions I want people to be asking so why is lead so we're asking why lead oxide why lead oxide is more stable than um well generally lead to oide anybody can tell me iner P effect inner perer effect lovely um so the inner perer effect is one of the things that contribute to stability of those um higher oxidation States right so if we look at PB well sorry sorry for saying PB um lead to oxide right versus um the lead oxide stability the in ver is one of the reasons why one is more stabiliz um stabilized than the other all right so what is the inner PA effect people may ask the iner perer effect well you guys don't go through the iner perer effect right but the iner effect really increases as going down the group right the group four right so we see that the carbon plus4 state is actually not stable because of the very poor shielding of f orbitals right so that is the F orbital we're talking about now right so we know that F orbitals right has really weak penetration compared to D orbitals so in the case of of the lead oxide is the electrons are going to be existed more less in the D orbital right is a little bit more stable than having lead to oxide where we have electrons existing in the F orbital which has very weak um penetration so let um water oide is actually more stable than Le to oxide all right the tables giving the properties of our monoxides now so carbon monoxide silicon monoxide gerum monoxide thin monoxide and Lead monoxide right so our carbon monoxide and silicon monoxide are simple molecular where the rest of them are predominantly ionic in nature our acid based characteristic now right they tend to have less acidic characteristics so acidity also decreases going down the group right for oide so it's so monoxide carbon monoxide is very weakly acid then silicon monoxide is actually neutral and then the rest of them are ampo thermal stability now all of them readily oxidize to their dioxide right it tends to be other than really the lead the lead doesn't really read oxidize this dioxide because we know that lead dioxide is actually less stable than lead oxide so this information here right what will tend to happen is that lead oxide would actually be oxidized in the presence of excess oxygen to lead dioxide and then start decomposing because it is not really found within that structure it's not stable all right so that's the information that we have there so hopefully everybody's understanding it's just going through this information technically this should be just reviewing all right the oxidation stat the syus ask us to look at the oxidation states which which is the last thing of our about this now right oxidation stat of our compounds we know that we exist in 2 plus or we exist in 4 plus right those are the most stable oxidation States right of course we have 1 plus oxidation state but that is less stable than 2 plus and 4 plus unless we're talking about lead you know unless talking about well no lead 2 plus would be more stable let 4 plus yeah so two let me just talk about that right like 2 plus and 4 plus would actually be more stable right cuz lead lead monoxide is actually lead 2 plus all right and it prefers that over it 4 plus stat okay so when we look at stuff like this now when we heat these elements with oxygen right we with the exception of lead they all would form an oxide with a plus4 oxidation state because that more stable right so we have that there so we must note however that the relative stability of 2 plus oxidation state and the stability of the pus oxidation state increases down the group right as the bond changes from Co valent to ionic right so what we're generally seeing here know is that as we go down the group we tend to prefer the 2+ than the 4 plus so oxygen um not oxygen carbon would prefers 4 plus right um silicon would pref prefers 4 plus but as we go down we start to change preference to the 2 plus oxidation state make sense since we go down we prefer 2 plus because four plus will become more unstable because of the inner Perfect all right okay are we okay so far guys and here we're called to look at the standard electrode potentials of these and we can know the feasibility of reaction depending on how positive the electrode potential is so this can help us figure out right um what's happening how we favor it so for the germanium 4 plus to be reduced to Germanium 2 plus that a negative cell potential therefore this is not feasible right and for the tin now it's electric potential is actually positive so it's a little bit more feasible right and then the lead um electrode potential here would be more positive than the team therefore this it would more more feasible right so hopefully we understand how to electr potentials all right so that is what it's telling us remember when I remember when somebody asked me not sure if it was Z who asked right when we were doing electric potentials what is the importance of this this is why we use it we can tell whether or not reactions will be feas so this is how we we can tell why the plus2 oxidation state is favored going down the group more so than the plus all right lovely and we know the acid based nature of our oxides already we know how the acidity of oxides decrease as we go down the group right so carbon dioxide would actually produce carbonic acid right or will not will carbonic acid High concentration carbonate which is ACI right when carbon reacts with sodium hydroxide it also gives us sodium carbonate and water so we know that it's acidic in nature same thing with our silicon dioxide it is acidic all right so silicon dioxide um reacts with sodium hydroxide we're going to get sodium silicate right plus water so the basic character increases going the going down the group and what that means is that the acidic character decreases going down the group okay so the oxides dioxides are more acidic than the monoxides that's one thing that we need to state dioxides are more acidic than monoxides right that needs to be stated all right so the plus4 oxidation state is much more acidic than the plus2 oxidation state so here are the reactions for the plus2 and the plus4 oxidation States now mhm because we know the difference in acidity and know we're looking at the aoic oxides now for the monoxides of theic bands right so we know that um as we go down the group basic character increases right so germanium would be less basic than our thin oxide which be less basic than our lead oxide right so they all will react with acid to form a so notice how they're basic look at how they're basic they're behaving in that way they're reacting with acids create water this is basic characteristics these are basic characteristics right they all react with Alkali to form oxons right so they're amp in in the fact that they can actually react with hydroxides as well to produce an acid acid derivative and an Oxo which which is our Oxon right acid Dera so when our gerum oxid reacts with our hydroxide we get germinate right when our tin reacts we get our state and when we react with lead we get our Plumb bait all right so these reactions are really interesting they're really cool I just love volume 3 all right so it's just something like that right we get those ions and we get water all right for our plus4 oxidation States now we know that plus4 oxidation states are more acidic that's what I stated above so with our plus4 oxidation States now right we know that our carbon dioxide and our silicon dioxide are going to be acidic but with our otter oxides now what what happens right we're going to get the same thing just balance the reaction that's all you have to do there plus4 oxidation state reactions are the same thing as plus2 ox State reactions for our OTC oxides all we have to do is balance because if we have two more now of oxygen we just balance oxygen going across the equation right so in this case now our Sil our um State now right our um tin dioxide rather are is going to react with our acids give us um our salts and water all right so they're basic in nature there we know that they're tic right and in this case now right with our hydroxides with our bases we're going to get our ions our germinate our state and our plumbate all right so there Amic please know that group four dioxides react with alkalides to produce Oxo ions right so group four otter oxides right react with our alkal us Oxo ions all right so we have that now so the monoxides now right are neutral so our plus um two monoxides right are generally neutral um and are neither acids or Alkali right while the rest of the monoxides are oser right so it's just good to know that our plus two oxides of carbon and silicon are generally neutral right they're slightly acidic weakly acidic right but they're generally neutral all right the last thing of this topic now use of ceramic based silicon so silicon dioxide now right is a good thermal indicator sorry not indicator insulator right and it has a very high melting point due to the many strong calent bonds within silicon within the Silicon um dioxide giant molecular structure right so it's generally used for the manufactur of glass and Porcelain all right so we generally use it um we mold it right in clay so silicon dioxide is actually present in clay as well we mold the different um shapes and we bake it in oven in large furnaces and ovens right in order to create our porcelain vases and those type of stuff right so as bestos here right an example of some of of of silicate right silicate we know what silicate is Right silicate is our SI I3 ion right so our silicate right would form long changes with magnesium calcium or iron ions will form what we call as bestos right tal as well tal is a very um is basically silicate combined with magnesium ions so tal is very basically a very soft rock right type of rock Clays know that we use Clays to create the different types of um Ceramics right Clays um are also basically sheets of silicate right about half of a silicon atom right are repl with am aluminium atoms right so that's why clay has this reddish color because of alumina right but we know that it's alumin aluminium silicate right which is Clay right so we have that there and we use the clay to create multiple different types of Ceramics as when we heat Clay on the high temperatures we can get a hardened brick like consistency right and the last part last part of our syllabus ask us to look at the ceramic applications of silicon dioxide right so silicon dioxide due to the fact that it has high thermal stability through its large crystalline structure they can we can use it to really design structures right um for their properties such as hardness strength and ability to withstand these high temperatures so they're used in the building of heat shields for spacecrafts right and for armor protection in military vehicles so we can use silicates right to really build these different types of structures based off of its thermal stability its strength and its hardness all right so it has a hardness similar to that of diamond right it's really hard okay so it's using um heat shields for spacecrafts right and armor protection that you can see on tanks and those type of stuff right armored vs so that is the use and that is our review of all of our group four um elements