in this video we're going to look at the structure of the atom and then carry that into a discussion on some isotope related things first thing we want to do when we think about the parts of an atom is actually go through and do a real quick comparing and contrasting of the different parts we have in atoms i chose to do this in table form so we have three main particles that go into making up an atom protons the neutrons and the electrons and we're going to discuss these three different subatomic particles with respect to relative mass relative charge and then finally the location where would we find these at now first thing i kind of want to point out is we're not going to talk about the actual mass of these or the actual charge of these instead we're talking about relative mass and what relative mass does is relative mass allows us to compare and contrast these very easily without having to worry about the super crazy small numbers that are associated with these tiny bits of matter the numbers are so small when you talk about the mass of a proton or a neutron or electron it almost doesn't have any meaning so we instead of talking about the actual mass in grams or any other unit we kind of have a new unit that we want to introduce for this so we say a proton has a relative mass of 1 amu net amu stands for atomic mass unit it's just a way for us to get down to that small scale of matter without having to use crazy small numbers that are really tough to kind of get a feel for so a neutron weighs just slightly slightly slightly more than a proton but i'm still going to list this relative mass as one amu an electron i listed the mass as zero amu's and electrons do have a mass so that zero is kind of in quotations because it's not true it's just so much smaller than a proton or a neutron that the mass of the electron doesn't contribute anything to the mass of the overall atom the actual mass of the electron is about one thousandth the mass of a proton and that's just so small you don't have to worry about that mass compared to the other ones it's too small to worry about it's kind of insignificant it doesn't contribute anything to kind of put this in a non-chemistry example it'd be imagine you step on a scale and all of a sudden a butterfly comes and lands on your shoulder is that butterfly going to change the weight you see on the scale well probably not it's very very small but does that butterfly have its own independent mass that we can measure sure it does it's just that when we combined your mass and the butterfly's mass it's pretty much just your mass and that's what we're seeing for the protons and neutrons compared to the electrons the electron just doesn't have enough weight to really worry about its mass compared to the other things same thing we did with mass we're going to charge more and talk about relative charge so we don't deal with crazy small numbers so the relative charge of a proton we say is positive one neutrons are electrically neutral no quotes needed there electrons have the same magnitude or size charge that a proton does but it's the opposite sign we also want to talk about the location of each of these protons and neutrons are located in what's called the nucleus meanwhile the electrons are located outside of the nucleus so we're going to talk about some models of the atom coming up next so this shard is all well and good you can compare and contrast these different subatomic particles using this chart but what this chart doesn't do a good job on is really giving you a good visual for what does an atom really look like how do these properties kind of play out when we're actually thinking about an atom so that's what we want to look at next what does an atom look like kind of on the left what i drew was the classic model of the atom you got a circular shaped atom with a small nucleus in the middle kind of very similar to a human type cell even there's a problem with this model the problem with this model is that this model gets the proportions very very wrong this type of model greatly over exaggerates the volume of the nucleus compared to the rest of the atom even if i were to make the nucleus the tiniest of dots that would still greatly over exaggerate the actual volume that we see in that nucleus so if this model doesn't work because the proportions are completely wrong how can we really get a good visual of what an atom looks like you know even even if the smallest dot makes the nucleus too big how can we really get an appreciation for what does this atom look like so i want to try to do this with hopefully some things you're a little bit familiar with i want us to think about a different way of visualizing the atom with some things you're hopefully a little bit more familiar with so on the left i have a professional sports stadium kind of an overhead shot of that i chose baseball you could do this with soccer football it pretty much all works the same a large sports stadium so envision that that entire large sports stadium is the entire atom now on the right another object you're familiar with is imagine one single small p i want you to imagine taking one of those small single p's and walking out into the middle of that sports stadium and dropping that p right into the middle of that stadium now you have a visual of that p sitting right in the middle of that sports stadium that p represents the nucleus the rest of the stadium represents the rest of the atom that gets you much closer to the actual proportions in there the nucleus is so incredibly tiny compared to the rest of the atom just like the p is so incredibly tiny compared to the rest of the sports stadium the nucleus only represents a small small small size volume of that actual atom it's like the p is only a small small small part compared to the entire sports stadium so when you're thinking about an atom i think this is a clever visual to kind of get you thinking and getting things in about the right proportions of how tiny the nucleus is compared to the rest of that atom you have two very very different worlds inside versus outside the nucleus inside the nucleus protons and neutrons you have an incredibly large mass in there that's where all the mass is but a very very small volume we think about what density is mass divided by volume we have a very large mass but a very small volume that is going to be extremely dense not only will it be extremely dense it's almost unimaginably dense if you would look at the density of a nucleus it would be about a million times more dense than ordinary matter that you manipulate so something like a gold ring imagine that gold ring weighing a million times more than it did that would give you an idea of that actual density inside nucleus besides also being very dense you also have an extremely concentrated positive charge you have a lot of positive charge jammed into that same small area so a very concentrated positive charge outside the nucleus very very different world just the electrons that's going to be a very very small mass but a large volume that is not going to be dense at all not only does it is it not dense in terms of mass and volume it's a very diffuse or spread out of negative charge the negative charge you have a lot of it outside the nucleus all the electrons are negatively charged but there's a lot of room so they're able to spread out that means a very spread out or diffuse negative charge you want to think about it as almost like wispy cloud like everything is very very spread out so you get these two very different and opposing worlds inside the nucleus with this crazy amount of mass but in that very small volume with that extreme density and all of that positive charge jammed up into that same little spot and outside the nucleus not a whole lot small mass spread out over a large volume negative charge spread out over that very very large volume as well so i'll talk a little bit about isotopes and there's a format that we use for describing isotopes and this kind of the format in the big box is the elemental symbol that is where you would actually place the element that you're dealing with if you're dealing with carbon you have a c there if you're dealing with nitrogen you'd have an n there in the upper left of that is where you place the mass number what the mass number is is the mass of the isotope we think about what contributes to the mass of an atom we recognize it's the protons and the neutrons so the mass number will be the same as the number of protons and neutrons in the lower left-hand corner is what's called the atomic number so the atomic number gives you the identity of the elements almost like the elements social security number and the identity of the element is based upon how many protons that atom or that isotope actually has different protons are what's going to give us the different elements there's one more thing you sometimes see in this isotopic format and that's in the upper right hand corner sometimes you see the charge what the charge is is it is the charge of that isotope and if we look at the two things that have charge the two subatomic particles that have charge are the protons and the electrons so the charge is the protons minus the electrons this is kind of important we're going to see a lot of things this semester have charges in it because the electrons are always changing electrons are almost like the currency of chemistry they're getting moved they're you're gaining electrons you're losing electrons meanwhile the protons are almost always staying the same so what this charge does it gives us a little bit of a counting in terms of what's our relationship between number of protons and the number of electrons we kind of go and look at the periodic table the periodic table is not set up in isotopic format here the atomic numbers are in the upper left that increase left to right so the atomic number for carbon is six nitrogen seven eight nine ten because this is an isotopic format we'll talk a little bit about that later in this video and what you see in for the mass numbers is you see average mass numbers below the element now remember what you see on any periodic table is not listed in the format that we generally write isotopes in someone do one real quick isotope problem we'll put a couple more in a different video i'll make sure we kind of get through at least one because it makes sense for some of the other things we want to do with isotopes is i put up an isotope in its standard notation 31 in the upper left 15 in the bottom left negative 3 in the upper right you have a symbol there so using that we want to complete this information so first thing i have the answers in here but how did we get where we got how do we figure out the number of protons well the protons we saw from the previous slides or the same as the atomic number well where's the atomic number at well the atomic number is 15. well that means our atomic number is 15. that's also going to mean atomic number and the number of protons maps that's going to give us 15 protons your number of protons and your atomic number always are going to be the same those are going to completely always match those have to match there's no choice but for them to match okay that's literally the definition of what they are well what else can we figure out here well we can also identify the mass number real quickly that's in that upper left hand corner that 31 well that 31 that we see here well that's the mass number so that's where we got that from and that's just matching the location for what it's trying to tell you we can also do the same thing for the charge negative three well that's the upper right hand corner tell us what the charge is sometimes you'll see the charge written as three minus sometimes you'll see it as minus three it means the same thing it's kind of like first name last name versus last name first name on a form it's telling you the exact same thing so we figure out the protons the mass number the atomic number the charge but we still have to figure out are the neutrons and the electrons let's start with the neutrons and here's how we're going to figure that out there's a relationship between protons neutrons and the mass number your relationship between protons neutrons and the mass number so how do we get the 16 in here well what has to be true for protons neutrons and mass number is the mass has to come from the combination of the protons and the neutrons so the protons and the neutrons combined have to give you that mass number so i had 31 as the mass number 15 of it came from the protons where did the rest of the mass have to come from well it had to come from the number of neutrons that was the only choice in that the last thing we kind of have to figure out are the number of electrons that we have there's a couple different ways we can do that one is kind of more algebraically we can say that charge equals protons minus electrons the charge is negative three the protons are 15. we're looking for the electrons that would be x so i mean we could do this kind of algebraically we could say okay well figure out these to figure out these electrons what we could do is we could say okay we use a little algae say okay negative 3 that's going to equal the number of protons which we already have which is 15 minus the number of electrons x for our number of electrons we could solve this we could get 18 electrons coming out so we could do it that way and that'd be fine kind of an algebraic way of thinking about it there's also a little bit more of a chemical way of thinking about this as well and this is where you can kind of just say okay i know protons are positively charged electrons are negatively charged i have more what do i have more of well my charge is negative three i better have more of the negatives that means i need more electrons than protons how many more three more so either one of those techniques you could really use to get it the 18 electrons that you're looking at for this we will do a couple more isotope problems in a separate video so please be sure to watch that separate video where we're going to do some more isotope problems so we've done a lot with isotopes already but what we haven't actually done yet is explain what exactly is an isotope so an isotope is a form of the same element with a different mass well if it's the same element it will have the same number of protons that's what gives each element its identity so if it's the same element with a different mass that different mass has to come from a differing number of neutrons so that's kind of what we're seeing here i put up the three main isotopes of carbon carbon 12 carbon 13 carbon 14. sometimes you won't see the atomic number listed they're not trying to put one over on you or be sneaky it's just that if you're doing chemistry you're going to have a periodic table out you can quickly see what carbon's atomic number is in that periodic table carbon's always going to have six protons the neutrons well it's protons plus neutrons giving you that mass number so the mass number being 12 i better have six neutrons the mass number is 13. i better have seven neutrons if the mass number is 14 i better have eight neutrons there when we look at carbon on the periodic table the 12.011 that we see represents the weighted average of all the isotopes of carbon there are no carbon atoms in the universe that weigh 12.011 none of them zero no carbon weight is 12.011 what that 12.011 represents is the weighted average based upon the abundance of how often you see each specific isotope so without going through the actual calculations i picked one where we can kind of see this easy the carbon 12 it has a relative abundance about 99 almost all the carbon we're looking at is carbon 12. well that's why the average is so close to 12 because most of it is 12. there's a little bit of carbon 13 about a percent what does that do that's going to bump up the average just slightly over 12. we have even a smaller amount of carbon 14 that's going to bump up the average so that's what you're looking at when you're looking at the elements on the periodic table what you're seeing is not the mass of a specific isotope but instead what you're looking at is the weighted average of all of the isotopes with respect to their relative abundance now it's kind of important there's a directionality to this the relative abundance information can tell you the weighted average but it doesn't work backwards just knowing the weighted average does not allow you to accurately predict what the specific isotopes are or what the abundances will ultimately be and what does that mean we'll go through another example showing this but what that means is if i knew the information on the table on the top the isotopes and their abundances i could calculate the 12.01 i can go in that direction i cannot go in the other direction if i just knew the 12.011 i couldn't predict or figure out or calculate what's going on in that table and what we want to do next is we want to do an example that kind of shows you that directionality so the element i want us to look at for this is bromine so bromine has an average mass number of 79.90 that's 79.90 what's that really really close to that is really really close to 80. being it's so close to empty wouldn't it be kind of tempting to expect that one of the isotopes of bromine would have a mass of 80. i think it'd be very tempting to do that not only would that isotope exist but being it's so close to 80 it's almost wanting us to say that it would also have a relatively high abundance maybe 90 95 that's where the mistake comes in this is a flawed way of looking at it you can't work in that direction because if you try working in that direction it just doesn't hold up you can't go in that direction here's why here's why i picked out that example we kind of on the left we're tempted to say one of the isotopes of bromine was 80 and that it was high abundance well on the right we see what actually happens in actuality there isn't a bromine 80 isotope it doesn't exist the actual main isotopes of bromine are bromine 79 at 50.7 percent and bromine 81 at 49.3 percent this makes the average really really close to 80 but an actual bromine 80 isotope doesn't exist so the average ends up being really really close to 80 but you never actually get 80. i want to put this directionality way of dealing with the averages and the abundance of the isotopes in kind of a non-chemical way that i think you're going to be a little bit more familiar with this works the same way as your course grades your course averages in all of your courses i want you to think about if your final average in a course was 93 does that mean you necessarily ever had exactly 93 in an assignment no it doesn't mean that that's that directionality also by looking at the final average you can't say with any certainty what a specific score on any assignment was just because you've got a 93 final average doesn't mean you scored 93 on any test quiz paper report project this works the exact same way like isotopes and average mass numbers if you know all the information about the isotopes including the abundance you can calculate that final average but if all you know is the final average that doesn't tell you anything about the specific isotopes kind of putting it back to that course great idea i want you to think about two people with the exact same final average 93 percent could they have gone about that a very very different way to achieve that final average absolutely one person maybe did really well on projects not so well in the exams maybe somebody else did the exact opposite maybe somebody else did really good on quizzes and papers and not so well on the exams they still ended up with a 93. that 93 average can't tell you anything specific about what actual assignments were that went into that but if you know all the information on the specific assignments you can figure out the average it's just like the isotopes and the average mass numbers the average mass numbers are the final grade the final average the isotopes and the specifics there are like those specific quizzes exams assignments remember there's going to be a separate video on isotopes where we do a couple more examples please remember to make sure you are watching that