S1: and so, between the lecture, the lab, and your discussions for Chem one-twenty-five, that should, be sufficient for you to do well, in Chemistry one-twenty-five (the advanced.) the difference probably comes to some degree in your backgrounds. the purpose of the lecture is to, cover the background material that you need, to get acquainted with to perform satisfactorily in the lab. however, typically there's fifty minute time intervals to do that. and so, you're going to have to decide, uh, after you leave the lecture whether you need to do additional reading on, a particular topic. because i'm just kind of reaching for some kind of common denominator, in terms of how i approach the material. and i'm also, uh covering it based on the feedback i get from you. so therefore it is important before we discuss topics, that you email me, with questions that you have, on that particular topic. so that i can maybe address them here in the lecture, okay as best i can or i will tell you, that it is not going to be addressed in the lecture because you are to discover that in the lab, and it will be discussed in your discussion itself. it is really the discussion is really important, that is where you share and discuss the results of the lab, and, really, focus on the implications of the results. and that is what the exams focus on. now in some cases you're going to have a discussion and you don't have one yet that's graded that's officially on material that we did on the exam, till you finish experiment two. and you may find that, your peer presentations, left you in some cases, (xx) uh feeling, well i'm not sure, about what those results were. it is really important to follow up on that, and in terms of even dropping in on another discussion, going to office hours, talking with your instructor about what happened there. also you will see on the syllabus that, just before exams occur, there will be a review session that i will do and some others will do, where i will highlight what the major findings were for the labs, to make sure that we are all in agreement. you don't have to come to the review session, i'll also put that up on the web. okay, so you have that to look at if you are uncertain about what the data was saying. okay, we'll address this term's, data. but between all those things you should be (directing) (xx) (this course) yourself, to, uh, to do well on those examinations. okay so keep those questions coming. the topic for your, upcoming lab you're starting your first officially graded lab they start on a two, it's called, uh water purity. and to illustrate that i have a cover of Chemical Engineer News, on the left screen, which has a picture of Venice Italy. and in Venice Italy, uh there is a big concern with the water purity, there are many industrial plants surrounding Venice that in particular discharge, heavy metals, into those waters. and those heavy metals in particular cause a lot of environmental damage. they're not only toxic, but they tend to get involved in many reactions including the type that you are studying in experiment two precipitation, reactions. now, when i say heavy metals i don't mean just the atomic form of metals, and i'm talking about, the metal that have higher masses so that they're, over here in the right of the Periodic Table and down into what we call the Transition and Post-Transition Area. and there in particular there's one culprit we're very concerned with mercury. okay and i'm going to be doing an illustration with mercury H-G, uh in a moment, but it is not the metal that actually is in the water that's the concern it's the metal ions. and an ion is a charged form of the metal. actually they charge, positively, and all our metals that go into solution, are positively charged they're called cations. and these cations tend to associate with negatively charged, species, called anions. and in our oceans in particular we have, a lot of, the salt sodium chloride, that contains N-A, that's the, symbol for sodium, and, C-L, the symbol for chlorine, and in solution we find we have the positively charged sodium atom called sodium cation, and it, tends to get attracted and, combine with chloride ion. and this is an artist's rendition of what is the most common salt in our oceans, sodium, chloride. all you'll be dealing with in this lab important, to keep in mind are salts. salts contain, positively charged ions called cations, and negatively charged ions, called anions. the way i think of cations is i picture a cat, two round circles for the eyes with, two, positive charges, (which) by positive charge meaning eye. whatever you can use to remember, what a cation is versus an anion that works for you, great. there's a lot of terminology in these labs and we need to find, methods for, learning, the language. so here we are all these compounds will contain, cations and anions, and we will call them salts. the exception to calling them salts, is the following, if the cation is a hydrogen ion H-plus, then will be calling an acid if it's got O-H-minus we call it a base, and if it's got oxygen O-two-minus, we call it an acid-anhydride. don't worry i'm not going to examine you now on, which are salts acids and bases, i just want you to know this for a fact. and another important fact is when these ions combine they combine in such a way that they neutralize each other's charges, so that the total positive and negative charge, is neutral. <P :06> and in the diagram... that i will put back on over here, i want you to note, that there's really no beginning and end to these compounds, they go on and on and you cannot find a finite piece that's just one N-A and one C-L, the N-A-pluses are surrounded by, anions above it and below it and to the side, and this is the only thing we can say about it is it is a one-to-one ratio. (so) there may be a billion sodium ions to a billion chloride ions when we write N-A-C-L we're really writing what we call an empirical formula just to give you the basic ratio of one ion to, another. one of the things that characterizes these compounds is that, they are held together very strongly in place through electrostatic forces. so that if you take a typical salt and you put it in in on your kitchen stove, and you turn the heat up all the way, the most it's going to do is pop around, in that pan, make, perhaps some noise, but it's not going to liquefy. not with the source of heat we have in our, typical kitchens, why? because it takes about a thousand degrees centigrade of a heat source to get these ions flowing past one another. so one of the things you should begin to suspect when you go into a lab and see a bottle containing a solid at room temperature is that it may be, an ionic, compound. it's one of the characteristics, of ionic compounds, that they tend to be solids at room, temperature. now one of the things you need to do and this is on that worksheet that, uh you may have picked up. again i will put the worksheets for the new topics, at the beginning of lecture on a stool, in each of the aisles. you can also download the lecture handout from the web page. and also when we give a completed lecture on a given topic you can go to the web and find my lecture notes, up there and comments. one of the things you need to be able to do is pick the right formulas for ionic, compounds. because you're going to be combining cations and ani- anions in solution, so we need to be able to know how to use the Periodic Table to predict their charges. so we're going to discuss the most common, charges that they have. to begin with, you need to know, your metals and your nonmetals so let me just take this off momentarily and put my other chart back on, you know most of our elements, are, metals, uh they are, to the left of the stepladder here on the Periodic Table and they're represented in blue on this uh, chart. and the elements in yellow are our nonmetals, and then we have just in that intermediate area what we call metalloids they're not particularly nonmetallic they're not particularly metal. but the important thing for determining charge, is the family, that the element is in and whether it's a metal or a nonmetal. okay? so, let's look at the metals, no let's look at the nonmetals, to begin with the yellow those elements in yellow. you'll see that those, nonmetals are (informally) families four-A five-A six-A seven-A eight-A. and the families, are the columns of the Periodic Table. those family numbers describe the number of, electrons in the outermost, shell of the element, and it's those electron(sic) that determine the chemical properties the chemical behavior of the element. and for some magic reason elements like to have a complete shell of electrons, which means, if they have less than eight, they will attempt to gain electrons to get eight if they're a nonmetal. okay? now some of you i can see are smiling and saying well probably i know this stuff already but fine be patient, this is the difference in background here too, gotta know this stuff. okay so if i'm an element in family seven-A, like F, which is fluorine has seven outer electrons, i'm going to gain one that's going to upset my electrical nu- neutrality by giving me a negative one, charge, and i'm going to be an anion with negative one charge and that's true for everything in family seven. same thing (works) in family six except for family six starting with O at the top oxygen on down, i have six outer electrons which (if) i can gain two, that's going to give me a charge of minus two. and, that formula that recipe works for all the nonmetals so essentially you're taking the family number and you're subtracting eight, to get the charge. they are the simplest ones, and they generally do not in any way, uh do anything differently, than either F minus four three two or one. now when we delve into the metals it gets, more, complex. and i'm going to go back to, your, chart that you have in your handout, uh to talk about the metals. and i'm going to look at family one-A, that's the first column, far left family two-A, skip this whole midsection, for now which is the transition families, and go to over to the far right family three-A and four-A and five-A, all of which contains, met- some metals. the metals, in those families have either one two three four or five outer electrons, and tend to lose those electrons, why? because when they do, then they expose the next inner shell of electrons which are complete, and they so- u- uh tend to be stable. so in that case, the, charge will be the same as the family number, okay? so family one-A there's one outer electron we lose one, negative charge now you're positively charged by one. family two two outer electrons you lose them, you end up positively charged two. same thing over in family three, when it loses electrons, positive three family four positive four family five positive five. okay? now i'm giving you their primary, oxidation states there are always exceptions in chemistry, however, ninety-five percent of the time these will work, for you. now when we get into the transition groups, it is more complex except for family, one-B and two-B all the transcript- transition elements are labeled family uh uh, uh family B, those in, one-B at one outer electron and two-B, two outer electrons, so they either have a plus one or a plus two. they're the simplest, to learn. there is an exception in that, copper, C-U here in one-B, tends to, also show, uh quite normally a plus two. that is the exception to to to the general simplistic rule that i'm giving you. now, i'd like you to look at families three-B four-B five-B six-B and seven-B, so the only one we wouldn't have is just the family eight, there. there, you're going to have, a primary oxidation state or charge these are synonyms, i've brought in another term now, the charge of an ion is also called its oxidation state we will learn. the primary charge is typically plus two. the plus two comes from the fact, that all these transition elements start with these two electrons, just by adding electrons, across the series so you start with the one for example of K and the two of C-A and they add the electrons to it. but, you don't have to worry about that and i'm not going to ask you about electron structure (where denominators) come from. simply, they all have oxidation states of plus two, but they also have charges that are the same, go up to their family number. to be explicit, in family three-B, if i am S-C which is scandium, i will have an oxidation state of plus two, or three. these all have multi-charges that's what makes them complicated. if i am, in family four-B, T-I for example titanium, i would be plus two, plus three or plus four, maximum charge the same as the family number. and that is true of all these elements, in families three-B to seven-B, plus two, all the way up to the family number. so we do one more M-N manganese in family seven-D can be, plus two three four five six or seven, okay. now, the eight-D group... i don't want to get into but i will say, ninety percent of the time if you pick plus two, it's going to be correct. okay? you will get into details of that when you ask questions about those details and the things you want (heard) but not in this course. okay? now, sometimes you're going to have a cation, combined, with other than, an element, as a nonmetal. so for example, the usual thing is a metal, for example w- uh we just saw sodium ion, family one-A plus one, with C-L, in family seven-A, so this is minus one. so we simply write N-A-C-L, and the typical names, of salts, are taking the name of the metal, sodium, the name of the nonmetal but change the ending of the nonmetal, to I-D-E. so instead of chlorine it becomes chloride, sodium chloride. occasionally you're going to see, a, cation combined with, an anion, where the anion has more than one, level. you are to assume, that when you see a group of non-metals clumped together, that they are an ion, themselves and the nonmetals are bonded together, okay, and they exist as one unit. you can, figure out the charge of something like N-O-three by simply going again to the Periodic Table, nitrogen is in family five, now because it is bonded to oxygen we put it in the metal row so (your old) family five is a positive charge plus five, oxygen is in family six, it's acting as the nonmetal so, it is minus two, so if we have, plus five for nitrogen, minus two for oxygen there's three of them that's minus six, plus five and minus six gives us negative one. you can figure them out using the Periodic Table, but it's up to you to, to practice, these, uh arrangements okay? and in this case it will have a different ending than I, and typically when you see oxygen it's gonna be ate, A-T-E, but, again that's not the purpose of Organ- Chem one-twenty-five, to make these compounds but you do need to know how to write a, uh formula. oh and we have on camera some formulas, for example, uh here's your N-A-C-L and these are all chloride salts, so you want to be able to, figure out what the formula for barium chloride would be, you have to find it on the Periodic Table first, barium, is here in family two-A, that means it's plus two, so it's chloride we're going to need two, of those, C-Ls, to balance that plus two charge, and there's the B-A-C-L-two, and we always put a subscript beneath the element to indicate the number of ions, that are in that compound. so here the ratio would be, two C-Ls for every one B-A. so specifically there're two C-L minuses for every one B-A-two-plus. and be careful when you start to write your prelabs and your lab, uh reports that you say chloride ion rather than chlorine or just chloride you'll lose points on the exam, if you do not distinguish between atoms and ions. to a chemist it is a big difference between lead and the lead ion. they're totally different species. okay and this one here aluminum, you have to know where it is on the Periodic Table, and on exams you'll always get a Periodic Table. here is aluminum, over in group three-A, so it's plus three, the formula has to be A-L-C-L-three to be electrically, neutral. alright if you look on your worksheet you'll see we're gonna do a little demonstration that has to do with the lab. <P :04> and that involves... taking, uh a water source so this could be, uh water in Venice, and we're going to put, a, salt of a heavy metal, into that water source. so there goes a little mercury chloride, on one side of that, pond or lake or whatever it is of course it's a petri dish so it's an, an artificial, environment, and on the other side we're going to put some, potassium iodide. the amazing thing that occurs in water is that these, salts, that take a thousand degree minimum of heat, to start melting, immediately start actually coming apart and going into solution so what happens is the the actual ionic crystal, the ions start to dissociate from each other, and what is happening is the same kind of thing that might happen in Venice, uh and Ed, Burton our demonstrator is pointing to it, we're getting uh this sludge, forming, right in the middle of, our lake. uh, and this is a solid coming out of solution, it's because these ions actually can migrate this is the demonstration of the fact that these cations and anions actually migrate, in water and move. they can go upstream downstream whatever, and you can get the positive charges the new combinations to the negative charges and, so on, so that in particular the mercury-two ion, in addition to being attracted to the chloride ion is also attracted to this iodide ion. so we started at mercury-two chloride but now we have this new possibility, of forming an iodide compound, with mercury, and indeed that's what that orange kind of sludge is in the middle of the, petri dish. uh, similarly the potassium ions, will be uh, both attracted to the iodide ions and, the, chloride ions. so any positive ion will be attracted to a negative ion negative to, positive. and what really happens here is that each of the ion combinations have, specific solubilities, in, uh, a water solution. for example there is a tremendous difference in solubility between potassium iodide and mercury-two iodide, and this is given in terms of the number of moles required moles that it can, uh dissolve in a liter of water, but if you look at this (one) a very very small number of moles of mercury-two iodide specifically two-point-two-times-ten-to-the-minus four moles, that can stay in solution, but if it's potassium iodide seven-point-seven moles, can stay in solution in a liter. so as soon as the iodide ion makes contact with mercury, we cannot keep as much in solution that comes out of solution, and of course the quicker they make contact the faster it comes out of solution. and so right now uh, on camera we have that same reaction where we're giving them very quick contact the mercury chloride and potassium iodide, and, this is what you tend to do in the lab you pour them together in direct contact and we get this, uh orange colored precipitate, that's the same stuff we saw in, the pond. and so, we say that what has happened, to this ion interaction, is, precipitation. and, when we, write that chemically, we have different ways we can symbolize it, uh we can write P-P-T, rather than spelling out that long precipitation word, or you can put next to H-G-I-two which is the precipitate an arrow pointing down showing it's coming out of solution, or also another synonym, is to put next to the H-G-I-two, brackets S, close the brackets meaning a solid is being formed. any solid coming out of solution is precipitation doesn't matter what it is a metal a salt or whatever, in this case it happens to be a salt. and the opposite of precipitation of course is something going into solution, because it's soluble, and uh, you can see on camera over here, that uh this is an example here in the test tube, of that mixt- same mixture that's been allowed to sit for a period of time, and what happens is that precipitate will actually fall to the bottom creating a sludge, and then we get a liquid above it and that liquid above is called supernatant. you'll get two phases supernatant, and the precipitate uh form. so now we come to very first part, first part of the lab that's the background. some terminology and stuff to get ready for. and i will say this, uh we're not aiming this week to get that far in this lab. just parts one and two. why? well, as you know it's been very confusing and hectic so far. why? because the university starts classes in the middle of the week, and so uh we had some labs meeting, uh during our first half week of school and getting an early start and others that didn't start, uh till that first full, second week of school so, therefore, uh we're gonna give you some time this week for all the labs to catch up so we'll be together, so we want you to get through parts one and two of this, but that means that, you should also practice preparing, for discussion which will not be held, for uh another week don't worry about it you're not doing the discussion but, each team is going to be assigned to question one two or three. and questions one two or three are based on parts one or two, which you will complete. so as soon as you complete parts one and two, try a hand at your discussion question prepare it which means manipulate the data, show it to your, uh instructor get some feedback on it, uh and then, if you have some free time fine. okay there will be some labs that will be going, the full time so catch up. okay so what is part one? in part one, uh you look at a reaction, where there's precipitate being formed, and you have to experimentally prove what is the precipitate. now in order to answer that question, you really need to pay attention to the information provided at the beginning of, the experiment. and that is that potassium ions and sodium ions, are highly soluble. we just saw that with the potassium iodide in the demonstration potassium ions, uh have quite a high solubility in terms of moles per liter that stay in solution. both potassium ions and sodium ions are in family one. indeed family one, as a family tends to be highly soluble cations... we use that information to, design reference blanks. okay? a reference blank, to be specific is a sample, that has all the components of your reaction mixture except for one species that you take out. okay so when i'm doing a piece of research, involving precipitation, i take out a species and i replace it with a species whose behavior i know. for example, i can, put into that reaction mission mixture a potassium ion because i know, what it does, it tends to be highly soluble and will not, precipitate. and it helps me to figure out what is going on in the reaction. here's an example of a blank for this reaction, that we would look at mercury chloride and potassium iodide. an example of a blank here would be mercury chloride and potassium nitrate we're not gonna demonstrate this one to you cuz it's just an example. and what have i done? i've taken out the I-minus ion, okay and i've replaced it with a nitrate ion. this would be an example of a blank that i would use to determine what the role of the I-minus ion is. in other words, what happens when i take it out? will i still have a precipitate? or will i not get a precipitate? it determines for me the role of the I-minus whether it's critical to the reaction or not. i know that the nitrate ions and that's those N-O-three-minus anions that are highly soluble, uh will not precipitate. okay so if you look at question number one, uh in your handout <P :05> we have that reaction, and we are, going to repeat it when we perform a reference blank test. and there is what we do in the labs so i have some N-A-C-L instead of H-G-C-L-two and we mix it with potassium iodide. the potassium iodide is common to both reaction mixtures, however, we have replaced the H-G-two-plus, with N-A-plus, and to determine what will happen when it's not there. what we get is we get no reaction. and that's one of two possible outcomes either we're going to get the same reaction hopefully or no reaction. and since we have no reaction now we know that without the mercury-two-plus, we will not get a precipitate. or another way of saying, uh responding to this was we know, mercury-two-plus is critical to precipitate formation. and you'll see on your prelab for the lab coming up experiment two, questions like this on the use of reference blanks. and by the way that prelab is due when you show up at your laboratory. your G-S-I has the right to, let you, stay out of the lab, till you complete that, prelab. okay to go on, (here is) from an exam. same reaction mixture, H-G-C-L-two and potassium iodide, now you will pick one of these three reference blank tests, that you could use to directly test the hypothesis that chloride ions are critical to, the reaction mixture. okay? so we want to look at a mixture, where we have removed, the chloride ions. <P :05> okay and, that means that it's not this one here because you still have H-G-C-L-two in there, and it's not this one because you have replaced and removed the mercury ions with potassium ions. so the correct, reference blank if you're going to be designing these for part one, would be this reaction mixture. okay so, this is assuming that my hypothesis is chloride ions are needed for this precipitation reaction, so that when i take them out, and in fact when you do that <P :07> you still get the original reaction in other words you see the same thing here an orange precipitation. and so my hypothesis turns out to be false, and this question asks you well my hypothesis is false what would you observe? well you should still have the reaction occurring, because those particular ions you just removed are not the critical species for that precipitation reaction to, occur. very important concepts there will be, some false reference blanks that will be designed, by you. okay? what is a false reference blank? okay that is any reference blank that that is incorrectly designed and it happens in research all the time that's way they make the wrong conclusion about the impact of decaffeinated coffee on our lives, or the impact of aspirin versus, uh Bufferin. because generally they have a test that's incorrectly (designed) and they leave some variables in it that skew the outcome. if you are testing precipitation reactions you have to again use a reference blank that includes a good testing species one that doesn't participate in precipitation reactions. if you're trying to see what sees uh uh causes a cancerous tumor to grow, and you're taking some component of that molecular substance of the tumor and altering it you can't put something into it that's gonna contribute to the reactions, okay? it has to be something you know is a solid little species in there that's going to act in a controlled manner and not, produce other reactions in other words it's not going to cause that cancerous tumor to grow you absolutely know, okay? so, uh, a incorrectly designed blank test is going to, get involved and start producing other side reactions, okay? a good reference blank test is going to produ- produce either no reaction, okay because it, doesn't get involved in precipitation reactions, or, you're going to get the same reaction that you originally saw in our case an orange precipitate. but if you start seeing something new you've got a false reference blank test. we're going to show you an example of that, uh this is instead of, mercury chloride it's copper chloride. we're going to add potassium iodide to it. so we've replaced the mercury, with copper, and what we get_ just let it stand there's it's fine it's fine (like that.) what you get is a slow reaction you'll get this murky muddy looking stuff that looks quite different than our beautiful orange uh precipitate here. so what we've done by replacing the mercury ion with copper ion, is simply put in a species that's producing all kinds of other reactions. okay, what we have to see is either no reaction or the same orange, colored precipitate. students mess this up on exams, all the time, and if it happens in lab when you're running your reference blank tests, all kinds of reactions, you've got to see no reaction or the same outcome of the initial, reaction that you are studying. okay? and uh... um you will see on your handout question two, you should try that on your own it involves false reference blank tests. try it on your own, and then when you're ready, go up to the web, the Chem one-twenty-five site, go under lecture handouts, and you'll find the key, to question two posted, and you can check your reasoning. uh, and that question is from, uh old exams. when you furthermore when you're analyzing these reactions you can compare the properties of the precipitate with samples of known identity, to determine what the precipitate is and verify what it is. one of things that is very useful, you will find in our instrument room, and sometimes in the labs, C-R-C handbook. there're lots of handbooks where you can look up the physical constants, of inorganic compounds all these salts are inorganic compounds. you will find an alphabetical listing, for example this is a listing for mercury compounds, uh starting with where the anion is with F, and here's your iodides, right here. i've (xx) precipitate of mercury iodide note that there are three different forms of the precipitate of mercury iodide and they included an alpha and a beta-two form, and you can find out what the color of these forms are for example, mercury uh H-G-two-I-two is yellow, and we have an H-G-I-two which is red, and another one here that's yellow, and it'll give you things like density, boiling point, solubility in different solvents, so as the course goes on you'll be using this handbook to reverify what you're producing in your reactions. okay part two <P :09> chemists use models all the time to make predictions, and so in that beaker, drawing is an artist's rendition of what precipitation is all about. and an attempt to explain how it can be, that you can put a salt into water, and the salt that's so tightly bound together, goes right into solution and apparently liquefies, yet it will not do so if you put it on a stove and give it a big source of heat. okay and so in the diagram here, you have a salt crystal at the bottom, showing all these anions and cations that are adhering together, and the water molecule, which is also charged and has a negative end by the oxygen part of the water and two positive poles those on hydrogen ends of the water molecule, approaches it so that the negative end of the water approaches a cation. and what it effectively does is neutralize the charge of the cation. as the charge is neutralized that ion can slip away from the crystal, then go into solution, once it's in solution the ion, then gets surrounded by water molecules. once that happens the ion charge is insulated from any other surrounding ion, charge. those, that process has terms you get more terms, here, there's lotsa terms. the process whereby, the, salt crystal goes into solution is called dissolution, and so here we say N-A-C-L solid, goes to N-A-plus and we put a cube, that means surrounded by water molecules, plus C-L-minus a cube surrounded by water molecules. that surrounding of the ion by water molecules also, is specifically called hydration. so it's the binding of the species to one or more water mo- molecules and it's signified by brackets A cube. very important concept. i think we know it, but maybe not in the whole context of chemistry. dissolving is the opposite of precipitation. okay the, salt can go into solution and dissolve, but it can also come out of solution and precipitate. because that is, a two way street, it is called an equilibrium process. really important. this is so important that i can't emphasize it enough. so precipitation is an equilibrium... system. we go either way, noticed here we have equal it can go either into solution or out of solution as, a, solid. and it seems to make more, a lot of sense doesn't it? that if something is not very soluble, it precipitates quite readily. okay? if something does not precipitate very readily it must be highly soluble. so the point of equilibrium varies for different substances, because their solubility varies. okay? now if you looked at this model this is the question you're asked before you do part two in the lab. you're asked to write down the hypothesis, which asks you to predict, how the cations you'll be testing in lab are going to behave. in other words can you predict which cations, will dissolve well or not dissolve well, or precipitate well or not precipitate well? if you're looking model what would a chemist say? i would say that, the size of the cation is going to affect solubility. you know and how well does it_ and what water molecules can get around it? what about the charge of the cation? well, remember that the water molecule has to neutralize the charge if the salt crystal is to dissolve. so, you would, on that basis see, that the higher the charge, perhaps the more difficult it would be, to get into, solution. with that in mind i want to point out, that in the back of your lab manual there is a Periodic Table of ion size, relevant ion size and you should be familiar with that pattern, because then some of these results will make more sense, to you. notice that as you go down any family, the size of the elements or the size of the ion they're both the same increases... okay? so if there is any, i- impact of size on solubility you should see differences within the family, even though the charge is, the same. okay? and if we go across a row that's a period, i want you to notice that the size of the ions are getting smaller, opposite. finally i'm going to point out for your prelab for part two, and briefly state in this last minute what you're doing in part two. for your prelab, you are told to prepare a grid. and the grid is pictured in the lab manual under part two, tells you exactly what to do. prepare this grid as directed, and you're going to be assigned (with) specific cations to investigate by your lab instructor. once you're assigned to a group of cations you're going to fill in this first column, with the cations you're assigned. so let's say let's say these are the cations you're assigned, you do that at the last minute. so you have this section on a piece of paper. and when you get into lab, you cover it with a transparency, like i have here. that allows you then, to take drops of, your sample that you're testing and put them in each of the grids. for example if i was assigned sodium ions, and all these will be in the form of nitrates why? because they're nitrates are very soluble so any difference in, precipitation tendencies will be due to the cation. so i put sodium nitrate one drop here and then in the next little, grid and so on all the way across, i'm not doing it, okay, and if you have the right side of the transparency there it will not fall off it will stick to it. one side it falls off the other side sticks. so test it before. and then you're gonna take a dropping pipette, and you're going to add the anion, and all the anions will be attached to potassium or sodium. okay? and you're going to just mix two drops together and you're gonna look at the drops to see, wh- when you're getting a precipitate and (whether) they're colored or not. and you're going to put your results into the computer, and voila, when you put your results into the computer, the computer's gonna spit out, a great big piece of data with lots of numbers on how your precipitates were observed and whether they were colored this is a year old yours is gonna have different results, and that's where (your) work begins, (what does it mean?) good luck.
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