Transcript for:
Chemical Bonding and Structure chapter 6

hello everyone today we're going to discuss ion ionic and molecular compounds and so we want to really think about what is going on with regards to the electrons so in chapter 4 we talked about what's going on inside and outside the nucleus and so now we're going to be talking what's going on outside the nucleus specifically with the electrons and again this is where most of the chemistry is going to happen so for an ionic bond that by definition is a transfer of electrons so we're going to transfer electrons from one atom to another atom so we're transferring those electrons and again that is the definition of an ionic bond is that we're going to be transferring those electrons so this usually happens with a s block metal and a p block nonmetal so that is just the general overview of where ionic bonds typically happen they're usually between an s block element on the periodic table and a p block element on the periodic table so let's go ahead and give ourselves a refresher if you're looking at the periodic table this is generally what you'll see a periodic table will look like the first two columns are the s block the middle set of columns that dip down are the d block and then over here on the right hand side of the periodic table is the p block and then down here the two rows at the very bottom are the f block and for an ionic bond again that is by definition a transfer of electrons that's usually going to happen between a s block metal and a p block non-metal so those two blocks ionic bonds we're not going to be discussing the d block or the f block when we're talking about our ionic bonds so we're going to be looking at ionic bonds between an s block element and a p block element and what also what we want to know how is the periodic table named why do we have the s block the d block the p block and the f block that is because the valence electrons in the shell or in the shell of the letter on the periodic table so all of the elements that are in the s block their valence shell electrons are going to be in the s subshell for all of the or the s orbital for the p block elements on the periodic table their valence electrons are going to be in that p orbital and so that is where the block names come from on the periodic table and again we're going to be primarily looking at the s block and the p block the definition of a covalent bond which we'll discuss later on in this chapter a covalent bond is a sharing of electrons so when you're talking about an ionic bond that is where you actually have a transfer of electrons and if you're talking about a covalent bond you're talking about sharing electrons so let's go ahead and look at table 6.1 and we're going to look at some various types of compounds so we have an ionic compound or a molecular compound and an ionic compound by definition remember this is a transfer of electrons and a molecular compound is a covalent bond which is sharing of electrons and so that is really important those are very different uh characters so the ionic compounds are going to form ions while the molecular compounds form molecules the type of bond that we call this is again if it's an ionic compound then it's an ionic bond if you're looking at a covalent bond those are molecular compounds and again a covalent bond is the sharing of electrons some examples would be of an ionic bond would be sodium chloride and we're going to look at that electron configuration and then a covalent bond would be water or something this is propane again you're not expected to know that but c3h8 is propane the propane structure that we have here and those are some examples of covalent bonds so let's go ahead and look at only right now we're only going to look at ionic compounds and look at how they become charged or why they become charged so again this is a transfer of electrons and we're going to look at sodium chloride so why is the sodium becoming positively charged and the chlorine atom becoming negatively charged making it chloride so we're going to look at sodium and chloride and right now we're only going to look at the chlorine atom we'll come back to sodium later but the chloride uh anion is what i want to talk about first so by definition an anion is the negatively charged species and how do we know that that chlorine is going to be negatively charged inside the nucleus chlorine is number 17 on the periodic table so inside the nucleus you should have 17 protons so right now we have 4 10 11 12 13 14 15 16 and 17. so remember the definition of a chlorine atom is that it is its atomic number and its atomic number is 17 so therefore it needs to have 17 protons chlorine always has to have 17 protons that you cannot vary that that number so inside the nucleus is where we define the number of protons that the atom would have and that is the definition of the actual element itself so inside the nucleus you have 17 protons if you just had cl minus um then outside the nucleus or excuse me if you just had chlorine then outside the nucleus you would also need to have 17 electrons and that is because 17 protons which are positively charged plus 17 electrons which are negatively charged equals the grand total zero charge so the number of protons inside the nucleus equals the number of electrons on the outside of the nucleus and that's for chlorine itself however in sodium chloride what we see is that we have cl minus so inside the nucleus the proton count always has to be the same because that is the definition of the element so cl minus then again we're going to have 17 protons no matter what because the proton count cannot change for an element on the periodic table but we have an extra negative charge for this chloride anion so we know that we have 18 electrons outside of the nucleus and the way that we know that is because we have 17 positive charges and we're trying to figure out how many electrons we have and this needs to equal a total of negative one so i'm going to isolate by i um i by its uh x by itself to figure out exactly what my charge is so i'm going to subtract 17 from the left i'm going to subtract 17 from the right and that's going to give me x is equal to negative 18. so i know that i need to have 18 negative charges and that is exactly what i have for chloride which is 18 electrons and so inside the nucleus you have 17 protons outside the nucleus you have 18 electrons and remember those are negatively charged so that gives you a grand total of a negative one charge you have one extra electron then you do protons so let's go ahead and we're going to talk about the electron configuration a little bit later so let's go ahead and look at a sample problem so we're going to look at um we're going to write the symbol for an ion that has 7 protons and 10 electrons so let's go ahead and just look at a first so inside the nucleus i'm going to go ahead and write 1 2 3 4 five six and seven so by definition the element that i'm looking for has to always have seven protons that number can never change and so number seven on the periodic table is nitrogen nitrogen is the seventh element on the periodic table it will always have seven electrons uh seven protons excuse me outside of the nucleus i'm told that i have 10 electrons so i'm going to go ahead and put them all on this circle around the nucleus even though remember the orbital shapes they're not all in the same subshell 1 2 3 4 5 6 7 8 9 and 10 electrons so inside the nucleus i have seven protons and outside the nucleus i have 10 electrons so i have let's go ahead and count how many this proton in this electron will go ahead and cancel each other out this proton and this electron will cancel each other out this proton and this electron will cancel each other out this proton and this electron will cancel each other out this proton and this electron are going to cancel each other out this proton and this electron will cancel each other out and then last but not least we'll look at this proton and this electron will cancel each other out so as of now all the protons are balanced now we have an extra negative one negative two and negative three charge so that is how nitrogen because remember it's number seven on the periodic table that can't change is going to have a negative three charge and that is because you have three extra electrons three extra electrons then nitrogen uh itself the neutral nitrogen so go ahead and try and work on b so what i will tell you is that b again is number 20 on the periodic table so that should be a a good start for everyone and we also want to remember that inside the nucleus you have 20 positively charged species and outside the nucleus you're going to have 18 electrons so now do you have more positive charge or more negative charge that is something that you'll always need to determine because not every element is always going to be negatively charged some elements of course will be positively charged so remember that this orbital cannot hold all these electrons we just draw it like that or i just draw it like that for simplicity to help uh really understand um or not really understand but to help uh really depict what is going on with the balance of outside the nucleus and inside the nucleus so now let's look at ions and group numbers ions and group numbers we're still in section 6.1 and we're still talking about ions so a transfer of electrons so let's go ahead and look at the periodic table and again we're primarily looking at a transfer of electrons and for the transfer of electrons we're mostly talking about an s block element and a p block element so that's really important so the s block is here and the p block is here and of course this is the d block and we're not really interested in that so for the s block you have column one and column two for the p block you have column three column four column five column six and then i need to add on two more columns column seven and column eight so for the s block and the p block and i want to keep the colors the same as i had on my previous slide so we're looking at the s or a slide a few slides ago so we're looking at the s block and the p block all elements in column one will form a plus one charge all the time a hundred percent of the time and we'll look at the electron configuration in a second all of the elements in column two are going to form a plus two charge magnesium calcium things like that in column three we're only going to be looking at aluminum and aluminum is going to form a plus three charge aluminum only then we're going to look at carbon and under carbon is silicon for column four those do not have a specific number number five on the periodic table which is nitrogen and phosphorus we're gonna primarily be looking at nitrogen that's going to form a negative three charge oxygen is in the sixth column of the periodic table that's going to form a negative two charge and fluorine chlorine bromine and iodine which are halogens are going to form a negative one charge and then here you have your noble gases also in the p block and those are going to form no charge at all so you can look at the periodic table trends are extremely important and know that based on the column number column one is always going to form a plus one column two is always going to form a plus two for column three a we're only going to be looking at aluminum that's gonna form a plus three charge column four follows no trend column five is going to form a negative three charge column six is going to form a negative 2 charge and column 7 is going to form a negative 1 charge so let's look at why that is why is call we're going to start with column 7. why does column 7 always form this negative one charge and the same is true for every single trend why does oxygen form a negative two charge let's look at the electron configuration so we're going to look at chlorine since our example was sodium chloride we're going to look at the electron configuration of chlorine itself so we have the 1s orbital the 2s orbital and then the p orbital which has the three subshells and then we have the 3s orbital and we have the 3p orbital so that is again electron configuration chapter 4. so let's go ahead and look at we know chlorine has 17 protons we're looking at neutral chlorine so therefore we need to have 17 electrons and we're looking at the electron configuration so let's go ahead and put in our 17 electrons we have electrons 1 and 2 3 and 4 5 6 7 8 9 10 11 12 13 14 15 16 and 17 and what we'll notice is that we have one empty slot which is uh great so now let's go ahead and look at argon argon is its closest noble gas argon is number 18 on the periodic table so we're going to write those exact same electron uh uh excuse me the orbital so you have your 1s your 2s your 2p your 3s and your 3p so we need to fill in 18 electrons for argon so 1 2 3 4 5 6 7 8 9 10 11 12 13 14 15 16 17 and 18. and what we'll note this is full it's fully full which is great and argon remember is a noble gas so it's not reactive and that's because all of the shells are fully full all of the subshells are fully full now let's go ahead and look at neon which happens to be number 10 on the periodic table it's the noble gas above argon so there's no noble gas in between argon and neon so we're going to look at neon so that's going to have a 1s a 2s and a 2p we'll just go ahead and put our 3s and our 3p to really drive home this point so neon has 10 electrons so 1 2 3 4 5 6 7 8 9 and 10. so again we're the question that we're really asking ourselves is why do all of the elements in column 7 form a negative one charge and we're looking at chlorine specifically chlorine 17. so we want to look at the noble gas configurations of argon and neon to figure out why chlorine likes to make a negative one charge chlorine is number 17 on the periodic table if remember we have one empty slot if we fill that slot with a single electron then its electron configuration is going to look like argon which is a noble gas and not reactive and then it would be fully full in order for chlorine to look like neon it would need to lose all of these electrons and that's too many electrons to lose that's not going to happen so chlorine is never its electron configuration can never look like neon however if you fill one if you add one single one more electron and i'll add it in orange you add one more electron in to that empty slot then now your electron configuration looks like argon so chlorine is going to want to become cl minus and that is because its electron configuration will be fully full just like argons so now let's go ahead and look at oxygen and i'm going to duplicate this page i'm going to go ahead and erase my chlorine but i want to be able to compare oxygen with the electron configurations of both neon and argon so i don't need to rewrite the whole thing and i'm going to go ahead and just erase this crossed out and i will erase this because again we're looking at the electron configuration for oxygen we want to compare that to both argon as well as neon so oxygen has an it's number six on the periodic table so it has six protons and neutral oxygen will also have six electrons so that would be uh your 1s orbital your 2s orbital your 2p orbital and then we'll just go ahead and write them in case if we need them your 3s orbital and your 3p orbital we may or may not need these orbitals so let's go ahead and fill in our six electrons for neutral oxygen so one two three four five and six in order for oxygen to look like excuse me oxygen is number eight on the periodic table it has six valence electrons but we want to look at all of the electrons so oxygen is number eight on the periodic table it has six valence electrons sorry about that so what we want to do is we want to put in our electron count so we should have eight electrons and now we do so we have eight electrons total so in order for our oxygen atom to look like our neon electron configuration wise we need to add one electron and a second electron so let's go ahead and draw those in orange so you need to have one extra electron in this subshell and one extra electron in this subshell so now this oxygen atom is going to have 10 electrons so your oxygen is still going to have eight protons but now 10 electrons so you have two extra negative charges so that becomes oxygen two minus and its electron configuration will now look like neon which is exactly what we would expect and that is why elements will either gain or lose electrons so now let's go ahead and look at an example where we're going to lose an electron or maybe more than one so remember column one on the periodic table is always going to form a plus one charge column two is gonna form a plus two charge we're not looking at the d block at all and over here is the p block and again column three four five six seven those are your columns column three we're only going to look at aluminum and that's gonna be a plus three charge column four doesn't make any special trends column five is minus three column six is minus two and column seven is minus one so we looked at why these are going to form a minus two and a minus one charge the same is going to be true for that minus three it's trying to get to its closest noble gas so now let's look at what's going on with these plus charges so we're going to look at lithium and lithium is number three on the periodic table and that means it has regular lithium not ionized number three on the periodic table that means it's going to have three protons and three electrons again we're looking at neutral lithium so let's go ahead and rewrite our electron configurations for the noble gases because everything is trying to become like a noble gas so we're going to write our electron configuration for helium which is number 2 on the periodic table and that's 1s and it has 2 electrons neon is the next noble gas in line and that's going to have 10 electrons we've already looked at that electron configuration so 1 2 3 4 5 6 7 8 9 and 10 electrons so let's go ahead and look at lithium lithium is going to have 1s and then 2s and then we'll just go ahead and put the 2p orbital we might need it we might not so lithium has 3 electrons 3 protons three electrons so i'm going to go ahead and put those in my three electrons in order for lithium to look like neon all of this would need to be full so that would mean we would need to add one and then two three four five six seven electrons and that's not going to happen that's too many electrons but if we lose this one single electron now it looks like helium its electron configuration will look like helium and that is a good thing so lithium if we think about what's going on with lithium positively charged lithium is always going to have three protons because remember that is the definition of an element is its proton count that cannot change outside the nucleus we're going to have two electrons because remember we lost this third electron in orange so inside the nucleus you have three positive charges outside the nucleus you have two negative charges so three minus two is going to be plus one so your overall charge is going to be lithium plus one and that's how that is represented and so all of the elements in the first column of the periodic table will lose one electron so it looks like its closest noble gas and all of the elements in column two of the periodic table is going to lose two electrons in order to look like its closest noble gas so that is a really in-depth understanding of what's going on with ions super important to know it's always the electrons that are changing so now let's go ahead and move on to section 6.2 which is all about ionic compounds and again we're still looking at ionic compounds so we're looking at a transfer of electrons they're being transferred if the electrons are being transferred from the s block metal to the p-block metal non-metal because remember electrons are negatively charged and the s block wants to lose electrons to become like its closest noble gas while the p block wants to gain except for the exception is aluminum the p block wants to gain electrons to look like its closest noble gas so we're looking we're still looking at a transfer of electrons because we're still looking at ionic compounds and electrons are being transferred from the s-block metal to the p-block non-metal and we just looked at those electron configurations so remember electrons are negatively charged and so the electrons are being transferred because they're going to get lost from the s-block metal and the p-block non-metal is going to gain that electron so now let's go ahead and look at some chemical formulas that depend on charge so now we know how to determine the charge but so the charges sometimes can never change so chemical formulas depend on charge and so again we're going to draw our periodic table here we have the d block which we're not interested in the s block and the p block and you always have a plus one and a plus two for the p block aluminum is gonna form a plus three row four no trend row five is going to form negative three row excuse me column uh six is going to form a negative two column seven is going to form a negative one and column eight is going to form a zero so if we're looking at sodium chloride where will you find these elements on the periodic table sodium is in column one of the periodic table and chlorine is in column seven of the periodic table so because sodium is in column one it's always going to form a plus one charge chlorine is in column seven of the periodic table so that's always gonna form a negative one charge so you have a plus one and a negative one plus one and a negative one and that's going to equal a grand total charge of zero which is exactly what you want things do not want to be charged so you just write your chemical formula is nacl a lot of times you'll see me write it as in a plus and cl minus and that is so that we can really understand what's going on with those charges it will help you predict chemical reactions later on in the textbook you'll see it and most often you'll see it written as sodium chloride i do not write ionic charges as that i always write them with the actual charge because it's really going to help me understand what's going on later on um when i'm trying to predict chemical reactions so this is a very good habit to get into is writing your atoms charged in what the corresponding charge species so now let's go ahead and look at magnesium chloride what is this chemical formula i will tell you that magnesium is found in column two of the periodic table so magnesium always has to form a two plus charge because it's in column two of the periodic table chlorine is going to form a negative one charge always so you have a chlorine atom with a negative one charge so your magnesium is going to have two positive charges your chlorine is going to have one negative charge and now if we add all of that up that would be two positives plus one negative charge is going to equal a total of plus one and that is not good we are always looking to have a charge of zero so the only way to make this work would be is if you had two chlorine atoms two chlorine atoms and we write that as a subscript so now we have our magnesium is still two positive charges your first chlorine atom is a negative one charge and your second chlorine atom is also a negative one charge so let's add all of those up you have two red pluses you have a green minus and you have a blue minus and so if you add all of those up you this plus and minus will cancel and this plus and minus will cancel and that's going to give you a grand total of zero charge so it's not mgcl we write that as mg we have one magnesium and we have two chlorines atoms again this is not going to be really helpful when you're trying to predict chemical reactions so instead of writing it like that let's go ahead and write mg2 plus and cl you have two of those and those are each going to be minus one that will help you when you write the when you get in the habit it takes one extra second and when you get in this habit it's really going to help solidify predicting chemical reactions so i always write the charged species even if i'm not necessarily asked to let's do study check 6.3 and we're going to write we're gonna work on a and b together so you'll want to have a periodic table in front of you calcium is in column two so that's always going to form a two plus charge oxygen is in column six so that's always going to have a two negative charge so for your calcium that has a two plus and your oxygen has a two minus your two plus and your two minus cancel each other out so you have calcium oxide cao and then you can draw your charges so let's go ahead and look at our two plus we have two plus charges and let's go ahead and look at our two minus we have two negative charges so the two plus the plus and the minus and the plus and the minus will cancel each other out you would now have a full charge of zero and that's exactly what you want so calcium oxide one calcium for every one oxygen atom now let's go ahead and look at b b we have magnesium and again you'll get comfortable with where all of these elements are in the periodic table this is in column two of the periodic table so magnesium is always going to form a two plus phosphorus is in column five on the periodic table so that's always going to form a three minus so if you wrote that out you would have two plus charges so let's go ahead and color code our plus charges in blue and let's write our three minus charges in red and that is going to add up to an overall charge of negative one and again we're always trying to obtain a zero charge so let's go ahead and figure out how do we calculate what the chemical formula is going to be let's look at our lowest common denominator and your lowest common denominator is going to be six two times three is six and 2 3 times 2 is 6. so all i'm doing is to figure out this lowest common denominator is i'm taking this 2 and i'm multiplying it by this 3 and i'm getting a value a lowest common denominator of six that's one way that you can find uh the or the um one of the lowest common denominators often it will work out uh where if you just multiply those two numerical values together you will find the lowest common denominator it's not always true but it's almost always true so we know our lowest common denominator is six so let's go ahead and first look at how many magnesiums we have so for the case of magnesium we're gonna have 2x is equal to six so right now we're just looking at magnesium which is our two plus charge or i should not say two plus charge let's just think about magnesium magnesium is always going to form a plus two charge we need to figure out how many of those plus twos we have so 2x is equal to six so i just need to divide both sides by two and then x is equal to three so i have to have three magnesium atoms and now let's go ahead and look at phosphorus phosphorus is always going to form that negative 3 charge so i'm going to say 3x is equal to 6 and i want to figure out how many phosphorus atoms i have so i'm going to divide both sides by three and now i will say okay x is equal to two so i must have two phosphorous atoms so i believe my chemical formula should be mg three p two so now let's go ahead and think about what is going on is this the correct formula so magnesium is going to form a plus two charge and i have three i think i have three of those atoms so that's going to be one two three four five six plus charges for the case of phosphorus phosphorus is going to form a negative three charge and i believe that i have two of those so two times negative three would be the first um phosphorus atom would be one two three negative charges and that second phosphorus atom still negative three so one two and three negative charges so now let's see do they cancel out so this positive charge and this negative charge cancel this positive charge and this negative charge cancel this positive and this negative charge cancel this positive and this negative charge cancel this positive and this negative charge cancel and last but not least is positive in this negative cancel so now everything is balanced correctly you have the exact number of positives and the exact number of negative charges so that is the correct chemical formula so it's easiest when you have uh no option magnesium is always going to form plus two and phosphorus is always gonna form plus three excuse me three minus you always find that lowest common denominator now let's talk about naming and write writing ionic formulas which is chapter three uh 6.3 section 6.3 we're naming and writing ionic formulas so we're still talking about an s block metal and a p-block non-metal the cation is your positively charged species and that gets its normal name and then your anion whatever is negatively charged that negatively charged atom is going to get the suffix i d e so that is going to be your suffix so let's go ahead and look at an example so you'll have sodium and chlorine so your sodium again we don't write it as nacl because that doesn't tell you much you know your sodium is positively charged and your chlorine is negative one charge and that's based on where it is on the periodic table that's one reason why the periodic table is set up the way that it is so your sodium is the positively charged species so that's the cation and your chlorine is the anion your negatively charged species so the sodium gets its normal name so we would name uh this excuse me sodium and then your anion gets the suffix ide so this would just be chloride sodium chloride and that is how you name that and so we will be going over this in uh lab in much more detail you're also going to get a lot of homework uh problems naming it's pretty straightforward the positively charged species gets named its normal name and then the negatively charged species get this the suffix ide so let's go ahead and look at an another example so we just looked at magnesium three magnesiums and two phosphoruses and we looked at why that is so your magnesium is going to always form the plus two charge and your phosphorus is always going to form a three minus charge so when we write it just as so it doesn't really tell you much so let's go ahead and rewrite that is mg we have three magnesiums and those each form a plus two charge and we have two phosphorous atoms and each of those form a three minus charge now when we write it out like this we can see that the magnesium is our cation and the phosphorus is our anion so the magnesium is going to get its normal name so we'll just call this magnesium since it's cation and then phosphorus is the anion so we're just going to call this phosphide ide we're going to change that suffix and that is how you name an ionic species an ionic charge so now let's look at an element where you can have more than one charge so remember on the periodic table it's set up very well um in the sense that all of the s block metals either form a plus one or a plus two charge so that charge is not varying your p block again your aluminum is gonna form a plus three charge column four we don't have a trend column five forms negative three column six forms negative two column seven forms negative 1 and your noble gas columns always form a zero charge so these are also not variable what is variable though is this d block those can form many different charges and that is why it doesn't follow the same trend and that is why we don't have a trend down here at the bottom for this d block because each element in the d block may be able to form many different charges so let's go ahead and look at our chemical formula cucl so how do we know what the charges of this copper atom we know that chlorine is always going to form a negative one charge so your cl is always going to be negative one again you want your charges to add up to zero so we're gonna have x minus one is equal to zero because your charges should always add up to zero so we're going to solve for x so we add this plus one to both sides so x is equal to one so now we know that this copper must be a plus one charge so you have cu plus 1 and cl minus 1. let me go ahead and draw that in red so we're going to call this copper and then we're going to open parentheses and use roman numerals one chloride okay copper one nothing gets changed with its name the anion gets the negative charge ide suffix so now let's go ahead and look at cucl2 remember elements in the d block of the periodic table can form multiple charges so this chlorine atom always has to form a negative one charge and you have two of them so that's gonna give you a grand total of negative two so again we know that our overall charge needs to be zero so x minus two is equal to let's go ahead and solve for x so that that way we are able to know what the charge of copper is going to be so we're going to add 2 to the left add 2 to the right so now x is simply equal to 2. so now we know that this has to be c u two plus and c l we have two of those in each have a negative one charge so we call this copper two because we know that copper has a charge of plus two copper two chloride you are not allowed to change how many of the atoms you have in a compound you cannot change that so you cannot say oh i'm just going to go ahead and say i have cu2 cl2 that is not um allowed you always need to actually solve for x to figure out the charges so let's go ahead and look at study check 6.5 so we have manganese s3 sulfur is under oxygen on the periodic table so our manganese can have various charges because this is in the d block your sulfur you have three of those and each of those have a negative two charge so uh that has to be equal to three times negative two is equal to negative six so your sulfur portion is going to be negative six so now we need to figure out okay what is this manganese and i'm going to go ahead and delete uh my pin so that we can look at the chemical formula so this time you have two x you have two manganese atoms minus six is equal to zero because again you want your formal charge to be zero we're going to add six to both sides to get rid of it on the left hand side so that will cancel each other out so now you have 2x is equal to positive 6. we're going to solve for x so we're going to divide by 2 on both sides and now x is equal to positive 3. so you know that this manganese has to be plus three charged so you have manganese and we have two of those and that has a plus three charge and then you have your sulfur and you have three of those and each have a negative two charge so what we do is we call this manganese because that's what it's asking for manganese three sulfide so you have to solve your equation to figure out what as the actual charge on the manga means and in this case it happens to be three so let's go ahead and look at 10 fluoride what is the value for tin so we're gonna go ahead and look at 10 fluoride so snf4 so we know that this fluorine atom has to have a negative one charge and you have four of those so four times negative one is equal to negative four and so our 10 is going to be uh that's what we're looking for so you have x minus four is equal to again you want your total charge to be zero because there's no charge on this uh group of atoms so you just solve for x in this case we just need to add four to both sides now x is equal to four so that must mean that we have 10 4 fluoride so that is how you figure out elements who what the charge is if you have more than one possibility let's go ahead and again there will be lots of practice naming these compounds in the homework so let's go ahead and talk about polyatomic ions which are again we're still talking about ions so we're still talking about transfer of electrons so section 6.4 is polyatomic ions and so we're going to go ahead and look at some names and this is on the next page and for this class you can have i should say this handout this table for the homework and exams that's fine it's an introductory to chem class i'm not going to ask you to remember all of these so you're going to go ahead and we're looking at polyatomic ions and that means many atoms so polyatomic poly many atomic atom atoms in the ion in the charge so many atoms many atoms with the charge that's what that means polyatomic ions so you can have sulfur which would be for the case of sulfur you can have so4 2 minus or so3 2 minus so42 minus is ate and so3 2 minus is ite and again you can use this table for the homework for phosphorus you can have po4 3 minus and that's phosphate or you can have po3 3 minus and that's phosphite and those are very different because again you have the phosphorus is not changing but the number of oxygen atoms you have are changing and again you cannot change the number of atoms that you have for the case of your carbonates you can have co32 minus or hco3 with a negative one charge and that would be carbonate or hydrogen carbonate and again that's that hydrogen that is going to be different if you have chlorine that is attached to uh four oxygens and it has a negative one charge that we call that per chlorate if you have a chlorine atom with three oxygens and one negative charge that's just chlorate if you have a chlorine atom with two oxygens and one negative charge we call that chlorite and then if you have a chlorine with one oxygen atom and one negative charge we call that hypochlorite and again you can use this table for the homework and this will also be provided for you on the midterm and so again here are some more in table 6.8 here are some more examples of the different polyatomic ions many atoms that are charged and their names so you do not need to remember all of these so let's go ahead and look at for example if we are trying to figure out the chemical formula for magnesium nitrate magnesium nitrate let's go ahead and look at we know nitrate is hopefully somewhere on this table so we're going to look at nitro nitrogen and we're looking at nitrate which is no3 so let's go ahead and give our nitrate is no3 with a negative one charge so it has a negative one charge again you can use this table for the amount of charges and the actual name so we have magnesium nitrate we want to figure out what that is so we know it's no3 and it has a overall charge of negative one magnesium is in column two of the periodic table so this has to have a plus two charge so two plus negative one is equal to plus one and so that is not a charge of zero and again our nitrate has a charge of negative one you cannot change that uh numerical value but what we can do is we can add two nitrates and that would be magnesium and then we're going to open our parentheses and write no3 and each of those have a negative one charge and we have two of those so now magnesium nitrate can be written just as mg open parentheses no3 and then we're gonna indicate that we have two of those nitrates now magnesium has a plus two charge each nitrate has a negative one charge so a negative one and a negative one and that is going to add up to zero and so your chemical formula is balanced so let's go ahead and learn to name again you'll be able to use this table for this class so naming polyatomics let's just look at that same examples you have mg and then no3 and you have two of those so magnesium is going to be your cation because magnesium always has to form a plus two charge so this is your cation so it gets its normal name magnesium and then you have your no3 and that has a negative one charge and so this we know it has to be our anion so this gets the suffix ide i mean excuse me the suffix oh excuse me uh gets the suffix on table 6.8 so again you're going to want to reference table 6.8 so this is magnesium you find you're looking for no3 so you go to your table you are find no3 and you notice that that is called nitrate and so all you need to do is state that this is magnesium nitrate from the chemical formula and again you will have access to using the table as well as using um all of those charges on the exam so you do not need to remember all of those uh all of those names for this class specifically so now we're gonna jump from 6.4 we're going to talk about section 6.7 and 6.9 and then we'll come back and talk about 6.6 and 6.8 together so now we're going to jump ahead a little bit to section 6.7 and this is all about electronegativity so we want to look at uh six electronegativity and as we're looking at the periodic table first let's define what electronegativity is it's how much an element wants an electron that is really important how much an element on the periodic table wants an electron so what we know is based on the electron configuration column one is always going to form call it a plus one charge column two forms a plus two charge the middle of the periodic table we aren't really sure aluminum forms a plus three charge column four nothing and then uh column 5 is going to form a negative 3 charge column 6 is going to form a negative 2 charge column 1 is going to form a negative 1 charge and then column 7 or column 8 the noble gases don't form a charge at all so by definition we're asking or uh by definition electronegativity is how much an element wants a single electron so again based on the electron configuration fluorine wants the electron the most brancium down here on the periodic table wants that extra electron the least it actually wants to get all of the s block want to get rid of a an electron not gain an electron and by definition electronegativity is how much an element wants one single electron so as we're incr as we're going across a row we're increasing in the amount of the element wanting a single electron as we're going down a column we are decreasing the amount an element wants an electron so francium really wants to get rid of that electron and fluorine really wants that electron so francium is the least electronegative element on the periodic table and fluorine is the most electronegative element on the periodic table so that's the overall trend so as you go up a column and from left to right you're increasing the electronegativity and that becomes really important when we're looking at molecular compounds so let's go ahead and look at fluorine's electron configuration 1s 2s and 2p fluorine is number 9 on the periodic table so we're gonna go ahead and fill in nine electrons one two three four five six seven eight and nine so fluorine really wants to fill this box with one electron and that is why fluorine is the most electronegative element on the periodic table it only takes one electron to become like neon let's go ahead and look at lithium lithium is number three on the periodic table and i'm choosing lithium because i don't want to write out the electron configuration of francium it's really long so lithium is number three on the periodic table and so it's going to have three protons and three electrons so we're going to go ahead and write 1s 2s and 2p and remember lithium we already saw this electron configuration in very extensive detail so if you were to just get rid of that electron then it becomes like helium and that would be lithium plus so lithium does not want an electron in that spot and remember the definition is of electronegativity is how much an element wants an electron so lithium is trying to get rid of electrons and that's why they're least electronegative so that is section 6.7 so now let's go ahead and look at section 6.9 and again we're going to put 6.5 6 and 8 all together and so 6.9 is all about intramolecular forces and we have uh excuse me intermolecular forces so intra means within a molecule so we're looking at in um let's see yes intra excuse me molecular forces is and i just want to verify excuse me intermolecular forces that means between two molecules intras within the same molecule and intermolecular forces are outside more than one element so we're looking at inter molecular forces so between two so intermolecular intermolecular is between two molecules intra so if it is a tra that is within the same molecule so we are looking at intermolecular forces so we're looking at forces between two molecules so let's go ahead and look at we have three different types of molecular forces so the first is the dipole dipole and that is when you have an element or excuse me a molecule that has a varying electronegativity so the chlorine atom is the side that's negatively charged and that making that hydrogen atom positively charged and you have a positive partial negative and a partial positive charge um the dipole dipole is going to be between that partial negative and the partial positive and again we know that chlorine is going to be the overall negatively charged species because it wants the electron density uh based on electronegativity and hydrogen therefore has to be the positively charged species and again just like a magnet you're only going to interact between uh opposite charges so hydrogen bonds are something probably the most important inter molecular force so hydrogen bonds and that happens between a hydrogen atom and a nitro in this case it will be a nitrogen atom and so a hydrogen bond is going to be between a hydrogen atom and some other element so it has to be between hydrogen and nitrogen for example in this case it can be between hydrogen and oxygen or it can be between hydrogen and fluorine and those are your only cases of a hydrogen bond is it has to be between a hydrogen atom and either a nitrogen an oxygen or a fluorine it cannot be any other element on the periodic table and then because the textbook goes over it i'll go over a dispersion force these are very unimportant and this would be between two nonpolar molecules so we're going to go ahead and look at ethane which is you do not need to know its name but it has a chemical formula of c2h6 so you have two carbon atoms and you have three hydrogens you are not expected to know how to draw this at this point but i will tell you that f8 looks like this but this carbon atom is the exact same as this carbon atom so there's no electronegativity difference so that is between two nonpolar molecules so we're gonna need to have two ethanes because again in a reaction you always have thousands and thousands of these molecules and again you have this blue carbon and this green carbon and there's no electronegativity difference here either and so you the only force that you would fill between these two fa molecules would be a dispersion force and so again these are pretty irrelevant they're not very common but those are the three types of intermolecular forces forces between two molecules so now we're going to circle back and talk about molecular compounds um and this is much different than ionic compounds so now we're going to talk about section 6.5 which is molecular compounds and this is all about the sharing of electrons so this is typically two elements or two atoms that are coming together the first one would be the p block and the second element is also going to be the p block so before we've seen uh for ionic compounds we saw an s block metal and a p block nonmetal here we're going to have a p block nonmetal and a p block nonmetal so now let's go ahead and look at an example of this so co2 would be a great example and we're going to go ahead and name this carbon and then dioxide so carbon is found in the p block oxygen is also found in the p block these are going to share electrons and we'll talk about how um how that happens later so right now we just want to go ahead and learn how to name so you have your carbon the first atom that's written just gets its normal name and then the second atom that's written gets the suffix ide and then you specify how many of those atoms that you have so carbon and then you have two oxygens dioxide so let's go ahead and draw our list or write our list so if you have two atoms it's dye if you have three atoms it's tri if you have four atoms it would be tetra five is penta six is hexa seven is hepta eight is octa so we have seven at hepta h-e-p-t-a and eight is octa so let's go ahead and look at an example another example of when we're going to use these prefixes so if you wanted to name a sulfur and six fluorines the first atom uh you notice is in the p block the second atom is also in the p block so that is going to tell you that it is a molecular compound so the first element gets its normal name sulfur and then we have six fluorine atoms so we're going to go ahead and name that hexa so sulfur hexa fluoride because you change the suffix to be ide indicating that it's negatively charged so if you have br2o this would be we have two bromine atoms so dibromo and then we have one oxygen atom so that's just going to be oxide dibromo oxide is how you would name that so again you will naming these is much more straightforward there's not quite so many things to remember as when you're naming a molecular compound and again you'll have lots of practice so let's go ahead and look at section 6.6 which is lewis structures so we're going to look at 6.6 and we're also going to look at 6.8 together which is three dimensional shapes so i i want to teach these together i think it's really important and again we're talking about molecular shapes so lewis structures we're talking about molecular compounds so if we look at the periodic table we've already looked at the case where we have an s block and a p block and we are not looking at the d block at all um except for to name that's the only time that we're going to look at the d block was is to name so you have columns three four five six seven and eight and remember column three we're going to look at aluminum forms a plus three charge column four follows no trend column five forms a negative three charge column six forms a negative two charge column seven forms a negative one charge and column eight we're not even looking at uh because they're noble gases so what we want to know is for when we're trying to draw the actual structure is we need to have the same number of bonds as our column number so that column number the periodic table is set up for you extremely well because that column number is extremely critical in really uh being able to to navigate um how what these charges are very quickly so it's no coincidence that the periodic table is set up uh the way that it is so with the element is going to make the same number of bonds as its column number so it's for example let's look at methane and that has a chemical formula of ch4 so this carbon is in column four of the periodic table so it's gonna make four bonds so we're gonna have a carbon and then our hydrogen is in column one of the periodic table so that's going to form one bond so one two three and four so each hydrogen is making one bond and each carbon is making four bonds so that's really the only way to write methane as you can imagine the structures might get quite crazy so let's go ahead and look at sample 6.12 and we're going to look at phosphorus trichloride so phosphorus is in column 5 of the periodic table and chlorine is in column seven of the periodic table so chlorine is going to want to make seven bonds and when we're talking about bonds we write we say bonds for shorthand but it can have or two lone pair of electrons uh excuse me or two electrons which is called a lone pair of electrons so it's going to either make some combination of one bond two bonds three bonds four bonds five bonds six bond seven bonds or eight pons or a pair of electrons that also counts as a a pair of electrons it counts as one bond and i'm going to put that in quotes because it's not officially a bond so the way that you figure this out is the center atom is the least electronegative atom so your central atom is going to be the least electronegative atom so uh between phosphorus and chlorine phosphorus is the least electronegative that's why i wanted to teach electronegativity first and then i'm going to go ahead and put a chlorine atom a chlorine atom and a chlorine atom around the phosphorus so uh we are looking for the three for the structure we're not going to try and figure out the three-dimensional shape yet but remember lewis dots phosphorus is in column five so it needs to have and this again is from chapter four so it needs to have five dots so phosphorus is going to have one two three four and i can pair up any dot that i want to so i'm just going to pair up this dot here that has nothing that i drew nothing above the phosphorous atom and the chlorine atoms can be around the phosphorus any way that you want them to but you should only ever have one central atom and that's going to be the least electronegative atom your chlorine is in column seven of the periodic table so chlorine should have seven dots so by chlorine atom each chlorine atom you're going to have one two three four and then i'm just gonna pair five six and seven because again a lone pair of electrons a pair of electrons can count as one bond so now this is a pair of electrons a pair of electrons and a pair of electrons and so that's why i paired the ones up that uh that are most likely not going to make a bond so that was my first chlorine atom my second chlorine atom i'm going to go one two three four and then i'm going to pair up five six and pair up seven and then this chlorine atom on the left hand side i'm gonna go one two three four five six and seven so now it looks like each chlorine atom has seven dots so as of now i'm going to go ahead and draw a bond between this phosphorus and the chlorine atom this phosphorus and chlorine and this phosphorus and chlorine and that will give me phosphorus and chlorine chlorine and another chlorine atom and then this chlorine has three lone pairs of electrons and this chlorine atom has three lone pairs of electrons and phosphorus has one lone pair of electrons and i'll go ahead and draw those in black so uh so that is how you draw phosphorus trichloride now things are going to get a little bit more complex when we're trying to figure out the uh molecular shape so the textbook splits this up into two different sections and so um so that is just how you draw the the overall shape so let's go ahead and talk about how you figure all of this stuff out because it's not really useful unless you know all of the rules so the very first thing that you want to do is you want to calculate the number of electrons that you have the actually the very very first thing that you want to do is figure out the least electronegative atom and that goes in the center that's your center atom if you can only have one central atom so these are the rules in order to figure out the three-dimensional shape um and to figure out a more complex problem like this phosphorus trichloride so the textbook doesn't teach it the way that i'm about to but i think that this way is much more uh straightforward so the very first thing that you want to do is figure out the least electronegative atom and you put that in the center and we'll we'll look at phosphorus trichloride again the second thing that you want to do is you want to get a total electron count so this is for each atom in the molecule and remember we're talking about molecular compounds so this these set of rules only apply when you have two p block elements on the periodic table you need a total electron count so phosphorus is in column five on the periodic table so has five valence electrons and chlorine has a total of uh seven valence electrons on the periodic table so if we're looking at our example of phosphorus trichloride phosphorus has one excuse me has five valence electrons so phosphorus needs five electrons chlorine each chlorine atom has seven electrons and you have three of those so that gives you a grand total of 21 electrons so your overall structure should have five electrons plus 21 electrons which gives you 26 electrons and i will tell you as a hint that number should always be even and we'll see examples of when it's not and how it how to work with that so we're multiplying it by this three because we have three chlorine atoms so when we draw we're going to put our least electronegative atom in the center and that phosphorus is going to have five valence electrons so one two three and then four and five valence electrons and then each chlorine atom needs to have seven valence electrons so i'm going to go ahead and put a chlorine to the left a chlorine to the right and a chlorine below my phosphorus it doesn't matter how you put them around the phosphorus atom you can make a wide variety of different shapes but you can only have one central atom and in this case it happens to be the least electronegative in all cases it's going to be so each chlorine atom needs to have seven electrons so i'm going to go ahead and put my first electron my second electron my third electron my fourth electron my fifth electron my sixth electron and my seventh electron so that goes it takes care of the first chlorine atom my second chlorine atom's still seven one two three four five six and seven and then my last chlorine one two three four five six and seven so i should have a grand total of 26 electrons so let's go ahead and count our green electrons first so you have two three four and five green electrons which is exactly what you would expect for phosphorus and now we're going to count our blue electrons and we would expect to have uh 21 blue electrons so let's figure out do we have 21 blue electrons one two three four five 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 blue electrons so we have a grand total of 26 electrons which is exactly what we calculated so now that is correct so now what we can do is we can actually go ahead and draw our bond between this chlorine and phosphorus this chlorine and phosphorus and this chlorine and phosphorus and this is where you really want to pay attention so we have our phosphorus and our chlorine our chlorine and our chlorine and i'm going to go ahead and keep uh my electrons you're allowed to draw the bonds as two dots or a line it doesn't matter but i'm going to go ahead and draw out all of the dots so that we can keep track of our electronic bookkeeping so i need to draw out all of these electrons and this time i drew out the bond between those uh that phosphorus and the carbon atom as phosphorus and chlorine excuse me that's 26 total electrons so the very first step that you do is you figure out who is the least electronegative and that atom goes in the center then you need to do a total electron count for each atom in the molecule that is the next thing that you need to do is an electron count for each atom in the molecule so that's what we did and we figured okay we have 26 electrons that cannot change here's our actual structure we have 26 electrons we're good to go the next thing that we need to do is we need to count the formal charge and you want this to be zero and we'll talk about when it's not zero later on so phosphorus is and this is going to be your column number minus the actual number of electrons so let's go look at phosphorus first so phosphorus is in column five so phosphorus is going to be column five now how many electrons do you actually have around that phosphorus atom you have one two three four and five electrons so that's good so five minus five is equal to zero so the formal charge on this phosphorus atom is zero so that's great because you're going for a zero charge when possible for the chlorine atom each chlorine atom should have it's in column seven of the periodic table because remember you're going to go uh based on your column number minus your actual number so this chlorine atom we're going to go ahead and look at the chlorine atom on the left first how many electrons are actually around that chlorine one two three four five six seven so seven minus seven is equal to zero so this chlorine atom has a grand total charge of zero the same is true for this one and this one you can count the electrons if you want to but you'll find that each chlorine atom is in column seven and each chlorine atom has a negative tr uh actually has seven electrons around it so it has a negative a charge of zero a net charge of zero so that is exactly what we want to do so you have a formal charge of zero so that's the next thing that you ask then the third thing that you ask and often these two things are taught together but formal charge is completely separate than octet so the formal charge is the actual number of electrons around an element so that's why it helps to draw your dots the actual number of electrons number of electrons around an element the next thing that you need to do in order to figure out the three-dimensional shape is now you need to account for your octet rule and we'll talk about what this is in one second your octet rule so we said okay we believe that our phosphorus should be in the center and our chlorine atom is we have three chlorine atoms and each of these chlorine atoms has three lone pairs of electrons and our phosphorus has a shape that looks like that or an electron configuration that looks like that okay so it's the exact same structure that we had on the previous page now we have to account for the octet rule and this is for bonds and pairs of electrons so remember the definition of a molecular bond is the sharing of electrons and so we only talk about sharing of electrons when we're counting the octet rule not formal charge so you have to count your formal charge and that's based on electrons your formal charge is solely based on the actual number of electrons around the molecule the octet rule takes into account the bonds and pairs of electrons and the octet rule octet means eight so there are specific atoms that need to have eight around it so let's go ahead and talk about this phosphorus first phosphorus is not one of the elements that requires um eight but you have in this case there is an octet so you have two four six and this second pair counts as eight because again remember when you're talking about a molecular bond you're talking about bonds that are sharing the electrons so this phosphorus atom thinks that both electrons in this bond belong to the phosphorus atom because they're being shared and you're not counting for formal charge you're counting for the octet rule two very different rules now for the case of chlorine so the phosphorus has its octet is satisfied and that is good so now let's go ahead and look at the chlorine atom is the octet rule satisfied so remember that you count here's two electrons so for two four and six so now we're going to count this bond because again it's being shared between chlorine and phosphorus that counts as eight so you have two four six and eight so this chlorine atom yes it has a satisfied octet and you're good to go but you have to count for everything so this chlorine has two four six and eight so its octet is satisfied and then this chlorine atom on the left two four 6 and 8 and it has its octet is satisfied so you have to count formal charge and octet rule separately to see if you have the correct structure so we believe that we do have the correct structure so let's go ahead and look at phosphorus and then you have chlorine and you have all of these atoms around the p so the phosphorus atom the very first thing uh the axe method and that's on page right here here's your axe method the textbook also shows this a little bit different but a is equal to the number of center atoms you have and you will be able to use this table x is equal to the number of groups you have and e is equal to the number of lone pairs okay so we're gonna be using table one you'll be able to use this it will be provided for you so you're not expected to memorize all of these shapes but you will memorize them in future chemistry classes but again this is a very introductory chemistry course so uh i'm not going to expect you to remember all of these things uh just know how to navigate this table so we're going to look at our phosphorus trichloride our pcl3 phosphorus is in the p block chlorine is in the p block so we know that this is molecular compound and that means sharing of electrons and the electrons that it's sharing are these electrons here these electrons here and these electrons here those are that where the electrons are being shared and you count those in your octet so our axe method our a is the center atom it's always going to be one for this class a will always be one your center atom so a is equal to the number of center atoms you have in this case it's always going to be one in this class x is going to be the number of groups you have around your center atom so you have one chlorine and i'll go ahead and highlight you have one chlorine two chlorine and three chlorine atoms so your x is equal to three that's the number of groups you have around your central atom and then your e is equal to the number of lone pairs on the center atom only so the number of lone pairs we're looking at this phosphorus atom only the number of lone pairs that you have around that phosphorus group is going to be you have one lone pair so your e is equal to one so your a is equal to one your x is equal to three and your e is equal to one so what shape is that um this handout will tell you you're looking for an a is equal to one and a is equal to one x is equal to three and e is equal to one so that is this right here and that is going to give you a shape of trigonal pyramid and that has a bond angle of 109.5 degrees apart you do not know what that looks like yet so the shape of this is trigonal pyramid and it has bond angles of 109.5 degrees apart if you want to know what that looks like uh trigonal pyramid then what you'll want to do is to actually visualize shapes in 3d use mole view and i'm going to go ahead and grab mold view really quickly so that that way you can use it or i'll show you how to use it so just bear with me one second and i'm going to switch the screen over to mole view so now you can see mole view so we're just going to go ahead and they might have switched their website a little bit it doesn't look like they did and i can link this webpage into the canvas shell just click close and then you're going to trash this because we're not really interested in looking at that molecule we had phosphorus in the center so we're going to go ahead and put that phosphorus atom in the center and let me trash that again and put the phosphorus in the center and let's see if i can actually zoom in a little bit and then put my phosphorous atom there we go so you can see a little bit better so i'm just clicked on phosphorus and i'm going to put that in the center of my page it's a little temperamental then i have three chlorine atoms one two and my three chlorine atoms you don't have to get the shape right on mole view so this is what it looks like in 2d you do not need to include that lone pair on this phosphorus atom you click 2d to 3d and it will convert that for you and now you can see that it is trigonal pyramidal and so we call this trigonal so you have your center atom so trig is one two and three and then that's where its name trig comes from and then pyramid this is more of a pyramid shape because they're not all in the same line you can see that it has this puckered shape where if i put all three of these green dots the chlorine atoms in the same line this phosphorus is puckered up and the like being at the top of a pyramid so that's where trigonal pyramidal comes from so now let's go ahead and talk about when you have double or triple bonds so let's look at a more complex example so we're going to look at a more complex example so when you have double or triple bonds i will link that mole view page if it's not already into your canvas shell so now let's look at uh double and triple bonds oh before we start talking about double and triple bonds excuse me let me talk about this octet rule because this is very important so your octet rule is for the following atoms your octet rule is for carbon nitrogen oxygen and fluorine those are the only atoms on the periodic table that need to have the octet rule exceptions to this rule if your center atom is not carbon nitrogen oxygen or fluorine you do not need to count the octet rule believe it or not most of the time your central atom will be um offshore often will be one of those so your exceptions are following the following beryllium makes a duet and beryllium is number four on the periodic table beryllium makes a duet boron is going to make a sex tag which that means instead of uh instead of um instead of eight bonds four uh boron should make six bonds the other exception to this rule is hydrogen and hydrogen forms a duet and that should form uh one bond two electrons one bond excuse me beryllium it doesn't form a duet that's not what that's called it is called uh i guess a quartet and that is two bonds so beryllium forms two bonds a quartet boron forms a sextet which is a total of uh six electrons or three bonds and then hydrogen forms a duet which is one bond or two electrons so whichever way that you want to think about it in terms of bonds or in terms of electrons those are exceptions to the rule aluminum forms a sextet is what we call it often a sextet uh often aluminum will form a sextet and that is either three bonds or six electrons anything above fluorine so if you have silicon if you have phosphorus if you have sulfur if you have chlorine bromine iodine all of these can have what we call expanded octets which means more than four bonds or eight electrons and those are most often the elements that you'll see so the octet rule is specifically for carbon nitrogen oxygen and fluorine those four elements on the periodic table exceptions to the octet rule is where you would need to have beryllium will form a quartet which is two bonds four electrons boron and aluminum will often form a sextet which is three bonds or six electrons and then hydrogen maximum can form a duet which is one bond or two electrons so now let's go ahead and look at a more complex example and this is uh again this is going to be a much more complex example and you'll have lots and lots of practice three-dimensional shapes is extremely key extremely critical so let's go ahead and look at hcn hydrogen cyanide hcn so out of these three elements on the periodic table carbon is going to be the least electronegative so i'm going to put carbon in the center i'm going to put my hydrogen to the left and i'm going to put my nitrogen to the right so that's the very first thing that you do figure out who goes in the center the least electronegative atom the second thing that you do is you need to do your electron count so for hydrogen hydrogen is in row or excuse me column one of the periodic table so hydrogen gets one electron and i'm going to put it to the right so i can make that bond carbon is in column four of the periodic table so carbon needs to have four electrons so one two three and four and then last but not least we have nitrogen which is in row five column five of the periodic table so nitrogen needs to have one two three four and i'm going to pair up this one it doesn't matter which one you pair five so now i'm going to go ahead and draw a bond between my uh blue dot and my red dot and then my red dot and my green dot and what i'll notice is that i have an unpaired electron here and an unpaired electron here so you have two unpaired electrons so that's not good for the case of nitrogen you have an unpaired electron here and an unpaired electron here so that's also not good so all i'm going to do is i'm going to go ahead and pair this electron up and i'm going to pair this one up and see if that will solve my problem so i'm going to go ahead and draw h and then my carbon in the center and my nitrogen and i'm going to draw my hydrogen dot is blue i'm going to keep my carbon dots as red but remember i moved that one up so i'm going to go ahead and put those dots accordingly and then same for my nitrogen i'm going to i put my two pairs here and i have two pairs here so then the next thing on my checklist that i need to do is count formal charge and remember formal charge is for the actual number of electrons so you take your column number to count formal charge for each atom you take your column number minus the actual number of electrons around that center excuse me around the element the atom that you're talking about so you take your column number minus your actual number is equal to your formal charge so in the case of hydrogen hydrogen has one dot around it so hydrogen is in column one of the periodic table it has one dot around it so your formal charge of hydrogen is zero your formal charge for carbon carbon is in column four of the periodic table you have one two three four dots around that carbon atom so four minus four is equal to zero so based on formal charges so far so good now let's look at nitrogen nitrogen you have one two three four and five so you take column number minus the actual number and that is equal to zero so all your formal charges add up so you ask yourself is this the structure so the very next thing that you need to do is you need to count for the octet so i'm going to go ahead and put my hydrogen my carbon and my nitrogen and then i'll put my blue dot and then i'm going to go ahead and put my red dots for my carbon as so and then for nitrogen i'll go ahead and put my green dots so the rule is or the exception to the rule is that hydrogen needs to form a duet so for hydrogen the it's satisfied it forms a duet there's two electrons around that hydrogen atom you're good to go so now let's count for this carbon atom so in my carbon atom i'm looking does it follow the octet rule you have two electrons three four oh no you only have six six electrons that's not an octet so that cannot be the chemical structure let's go ahead and look at for nitrogen as well nitrogen you have two four and six six electrons and remember that is for the octet rule it is always going to be four the octet rule has to be satisfied for carbon nitrogen oxygen and fluorine in this case we have a carbon atom and a nitrogen atom and neither of those form an octet so your formal charge is good but your octet count is bad so this cannot be the structure luckily we have some double and triple bonds that we're allowed to make so let's go ahead and draw this out one more time so you have h and c and n and i'm going to go ahead and draw at my hydrogen i know that that's what that looks like my nitrogen atom i had like so and then my carbon atom i had two dots up here a dot here and a dot here so all i need to do is i'm going to move that electron there and i'm going to pair it up with this green electron and that will give me h and c and n and so then i'm going to have a single electron there and i'm gonna have one blue electron from that hydrogen atom and then i'm gonna have i moved in one green electron i already had a green and a red electron paired and then i have a single electron left but the just based on your electron count you have a single electron up here and a single electron up here you cannot have single electrons so that also cannot be your structure because you cannot have single electrons so let's go ahead and move that red electron in and pair it up with that green electron and see what happens so i'm going to go ahead and flip the page because i don't want to cram anything too much on this page but remember i'm moving those i'm pairing up this red electron in this green electron so all i need to do is write h and c and n and my hydrogen had its blue electron and that was paired up to a red electron then i moved the two electrons that were on top of that carbon atom in between the carbon and the nitrogen and i already had one down there so then i moved the two electrons that were on top of that nitrogen atom down to make a carbon nitrogen bond and then i had two electrons left over so i'm going to go ahead and pair everything up and see what this looks like so now let's count formal charge remember formal charge is an electron by electron basis so for hydrogen hydrogen still has one electron around it and it's in column one of the periodic table so hydrogen has a column number of one right now it currently has one electron around it so your formal charge of hydrogen is zero so hydrogen checks out for the case of carbon carbon is in column four of the periodic table i currently have one two three four red dots around carbon so four minus four is equal to zero so your formal charge of carbon is still good even though you moved all those electrons around in the case of nitrogen nitrogen is in column five of the periodic table around nitrogen i have one two three four and five dots so i still my formal charge on every single atom is still equal to zero so we're good to go there so now we need to count octet and remember octet is a pair of electrons or we can think about this as being a bond the exception is the hydrogen here so hydrogen forms a duet so you have two electrons so your hydrogen is still satisfied based on the octet rule this carbon atom needs to it is one that has to follow the octet so you have two four six and eight because you count the pair of electrons for the octet rule so as of now your carbon follows the octet rule and nitrogen also one of the elements that must follow the octet two four six and then you have this pair for eight eight electrons and your nitrogen also follows the octet so your formal charge checks out and your octet rule checks out so now this is your actual structure and we draw that as h c triple bond n with a lone pair of electrons on that so now let's talk about your axe your a is equal to one center atom always that carbon atom has one group and two groups around it so your x is equal to two and then there are no lone pair of electrons on the carbon atom so your axe is equal to a is equal to one x is equal to two and e is equal to zero so you go to your table and you find a is equal to one x is equal to two and e is equal to zero or you can just drop the e it's the exact same thing so we call that a linear shape but the electron geometry and the molecular geometry are the same that has a bond angle of 180 degrees so we're going to call this linear and your bond angles are 180 degrees so now i'm going to go back to mole view and i want to draw that structure so i'm going to draw my carbon atom in the center it's attached to one hydrogen atom and the carbon got deleted and that's okay because carbon can just be shown as uh the end of a stick and then you have a nitrogen atom and we need to make that into a triple bond and that is the structure of what it actually looks like how you would draw it shorthand here's your hydrogen your carbon atom and your nitrogen we're going to click 2d to 3d and we'll see that here's the structure of hcn so you can see that you have your carbon atom in gray your nitrogen atom in blue and your hydrogen atom in white and these are 180 degrees apart and they're all in a straight line and so three dimensionally that's what it looks like if you rotate this molecule around everything is in a straight line and you can see what it looks like in three dimensions so mole view is really helpful to visualize what the molecule actually looks like in its three dimensional shape so now let's look at an example this is the uh so now let's look at an example with a negative charge so this is the handout from the textbook it's not as clear as straightforward um it doesn't give you as much good information as this handout here which is posted in your canvas shell that you will utilize so let's go ahead and look at something with a negative charge so let's look at clo4 minus so we're going to look at clo4 minus so clo4 and then a minus charge so your chlorine is in column seven of the periodic table you have one of those so that should be seven electrons your oxygen atom is in column six of the periodic table and so that each oxygen atom gets six electrons and you have four of them so that is equal to 24 electrons so 24 plus 7 is equal to 31 electrons and the problem with having an odd number of electrons is that something will be unpaired so this negative charge we'll just add that in so one you have one extra electron now that's going to give you 32 electrons 32 electrons and now you have an even number of electrons and so everything can either make a bond or a pair so that is any time that you have an odd number of electrons there will always be some charge up there to take that into account because you can never have an unpaired electron in molecular i should say you often won't have any unpaired electrons so let's go ahead and look at this the center atom is going to be chlorine and then you have four oxygen so one two three and four so each oxygen atom gets six one two three four five and six one two three four five and six one two three four five and six one two three four five and six excuse me i need to have some of these unpaired to make the bond so each i've uh i have let me count um how many oxygens i have i should have 24 red dots so 2 4 5 6 7 8 9 10 11 12 13 14 15 16 17 18 19 20 21 22 23 24 so i'm missing no electrons i have 24 electrons so i'm good to go there my chlorine atom needs to have seven electrons so i'm going to go ahead and draw those in blue so one two three four i'm just gonna go ahead and draw five six and seven so i notice that i have an unpaired electron here and an unpaired electron here so i'm just going to go ahead and move that electron in i have an unpaired electron here so i'm going to pair it up with this electron and then i have an unpaired electron here and i'm going to pair it up with that electron and i have an unpaired electron here and i'm going to pair it up there so that gives me cl and double bond o double bond o double bond o and double bond o so i need to count to make sure i have the correct number of electrons so i have two four six and seven and then i have again that negative one charge should be let me see i meant i have one extra electron somewhere just give me one second while i'm missing let's see ah i see where it is sorry about that so these are this is already paired so i can't move that electron in but i do have one extra electron from the chlorine so i'm going to pair up that um set of electrons but overall my structure i still get the exact same structure so i'm going to draw my lone pair of electrons on this chlorine on these oxygens so now i'm going to account for formal charge because remember that's the next thing that we need to do so uh again you have an overall charge of negative one so that needs to be taken into account somewhere so we do want to consider that so for oxygen you have we're counting formal charge two four five and six because remember each for formal charge it's only the number of dots around that oxygen atom so those are the there's are where those two red dots are coming from so you take your column number minus the actual number and that is equal to zero so that oxygen is equal to zero and so all oxygens are because they all look the same chlorine is in column seven of the periodic table i currently have one two three four five six seven and eight so eight minus seven minus eight is equal to negative one so my chlorine has a negative one charge and that's fine chlorine can have a negative one charge i'm gonna that is because again you have this negative one that needs to be accounted for so you would expect your some element to have a negative one charge this case it's chlorine you're all good so now let's count for our octet i'm gonna account for my chlorine for a levy count actually for my oxygens first and that is um i'll explain why in one second so 2 4 6 and 8. so each oxygen has its octet is satisfied and remember if you go back to the rules for the octet that is going to be oxygen in this case is one of the elements that has to have an actual octet it's not under this list of exceptions however we see that chlorine is down here that can have this thing called an expanded octet so what is that expanded that means more than that means above you expanded it so let's count the octet rule for the chlorine atom you have 2 4 6 8 10 12 14 and 16. so based on the octet rule chlorine is a 16 at we'll just call it that but that's fine because an expanded octet is greater than eight and that is exactly what we have more than eight so we're good to go there so our expanded octet is totally satisfied for the atom that it needs to be uh chlorine and our regular octet is satisfied for the atom that it needs to be which is oxygen so uh we're gonna just go ahead and indicate that our negative one charge is on that chlorine atom and that is the actual uh structure of clo4 minus so there is lots and lots of practice in this handout and so you'll have all of these different uh examples to go through and you'll also have a lot of uh practice in the homework and i would strongly recommend looking at the three dimensional shapes using mole view which i showed in this lecture so just to summarize i'm going to go ahead and put the summary at the end of the handout so that that way we'll have a summary but it's at the very end excuse me so today's lecture was quite long um and that is because really understanding molecular shapes is a big topic so in today's lecture in section 6.1 we talked about ionic ionic transfers ionic bonds which is a transfer of electrons and this was section 6.1 ionic bonds transfer of electrons followed section 6.1 to section 6. 4. section 6.1 to 6.4 and this is between an s block element and a p block element if it's an s block element and a p block element you're going to have an ionic bond so that and that follows within section 6.1 to 6.4 in section 6.2 we discussed ionic compounds and how the chemical formulas depend on charge okay that's what ionic compounds and we talked about the determining the chemical formula in section 6.3 we discussed naming and writing ionic formulas and you will have those tables that you will be to you able to use in 6.4 we talked about polyatomic which is many atoms ions and again we talked about naming those and so on and so forth then we jumped to 6.7 and we talked about electronegativity and by definition electronegativity is how many how much an atom wants one electron and we looked at electron configurations a lot today and the overall trend was as you are increasing from left to right and a row or as you're increasing up a column that is increasing electronegativity so the overall trend is up and to the right then in section 6.9 we talked about intramolecular forces and we talked about dipole-dipole we talked about dipole-dipole we talked about hydrogen bonding and we talked about dispersion the next chapter we talked about 6.5 which now we're getting into molecular compounds which is all about sharing electrons and this is often between or it's yes most often between a p block element and a p-block element on the periodic table those are going to share electrons so much different than a ionic bond and in uh 6.6 and 6.8 i taught those two sections together because uh it's much easier to teach those together we talked about lewis dots and molecular geometries here we used mole view and that is really important to look at the 3d shape so we always want to think about molecular compounds in their three-dimensional shape here we learned about the octet rule and we learned about formal charge and we want to remember that these are separate rules they're often taught together which they should not be they're separate rules and so we want to think about those separately so that is chapter six it's a very dense chapter it's a lot of information lots of practice and really understanding the difference between ionic molecules and molecular molecules and then thinking about the molecular molecules in their three-dimensional shapes