|Ch.1 - Intro to General Chemistry||2hrs & 53mins||0% complete||WorksheetStart|
|Ch.2 - Atoms & Elements||2hrs & 49mins||0% complete||WorksheetStart|
|Ch.3 - Chemical Reactions||3hrs & 25mins||0% complete||WorksheetStart|
|BONUS: Lab Techniques and Procedures||1hr & 38mins||0% complete||WorksheetStart|
|BONUS: Mathematical Operations and Functions||47mins||0% complete||WorksheetStart|
|Ch.4 - Chemical Quantities & Aqueous Reactions||3hrs & 55mins||0% complete||WorksheetStart|
|Ch.5 - Gases||3hrs & 47mins||0% complete||WorksheetStart|
|Ch.6 - Thermochemistry||2hrs & 21mins||0% complete||WorksheetStart|
|Ch.7 - Quantum Mechanics||2hrs & 35mins||0% complete||WorksheetStart|
|Ch.8 - Periodic Properties of the Elements||1hr & 57mins||0% complete||WorksheetStart|
|Ch.9 - Bonding & Molecular Structure||2hrs & 5mins||0% complete||WorksheetStart|
|Ch.10 - Molecular Shapes & Valence Bond Theory||1hr & 31mins||0% complete||WorksheetStart|
|Ch.11 - Liquids, Solids & Intermolecular Forces||3hrs & 40mins||0% complete||WorksheetStart|
|Ch.12 - Solutions||2hrs & 17mins||0% complete||WorksheetStart|
|Ch.13 - Chemical Kinetics||2hrs & 22mins||0% complete||WorksheetStart|
|Ch.14 - Chemical Equilibrium||2hrs & 26mins||0% complete||WorksheetStart|
|Ch.15 - Acid and Base Equilibrium||4hrs & 42mins||0% complete||WorksheetStart|
|Ch.16 - Aqueous Equilibrium||3hrs & 48mins||0% complete||WorksheetStart|
|Ch. 17 - Chemical Thermodynamics||1hr & 44mins||0% complete||WorksheetStart|
|Ch.18 - Electrochemistry||2hrs & 58mins||0% complete||WorksheetStart|
|Ch.19 - Nuclear Chemistry||1hr & 33mins||0% complete||WorksheetStart|
|Ch.20 - Organic Chemistry||3hrs||0% complete||WorksheetStart|
|Ch.22 - Chemistry of the Nonmetals||2hrs & 1min||0% complete||WorksheetStart|
|Ch.23 - Transition Metals and Coordination Compounds||1hr & 54mins||0% complete||WorksheetStart|
|Transition Metals||22 mins||0 completed|
|Transition Metals Properties||32 mins||0 completed|
|Coordination Complexes||29 mins||0 completed|
|Naming Coordination Compounds||22 mins||0 completed|
|Coordination Isomers||9 mins||0 completed|
|Transition Metals Electron Configuration|
|Oxidation States of Transition Metals|
|Crystal Field Theory|
|Colors of Complex Ions|
Concept #1: Coordination Compounds
Hey guys. In this new video we're going to take a look at coordination compound. Now, we're going to say here the most important or most prevalent aspect of transition metal chemistry is the interesting types of compounds they form, we refer to these compounds as coordination compounds or coordination complexes and important things to know about them is that these coordination complexes or compounds they are usually colored. So, they form very vibrant colors, reds, greens, blues, purples different types of colors because of the transition metal involved also. So, they're usually colored and paramagnetic, so the transition metal that's in the center of them usually has at least one electron that is unpaired. Now, we're going to say here, with Na coordination complex, here this is our coordination complex that we're talking about, it's made up of different things, we're going to say the coordination complex there is at least one thing that we call a complex ion and we're going to say complex ion is a species that is made up of a metal cation, so the transition metal in the middle, that is connected to molecules that are neutral or it can also be connected to anions and, negative ions, these neutral molecules that are connected to the metal or these negative ions that are connected to the metal are called ligands. Now, this is important to understand, the metal cation is going to act as an electron pair acceptor, we stay is acting as an electron pair acceptor, we're basically saying that it's acting as a lewis acid and then here the ligands, which could be neutral molecules such as ammonia or water or they can be negative ions such as chloride ion or fluoride ion, these ligands they're acting as electron pair donors. So, they have lone pairs on them that can donate electron to the metal cation since they're acting as electron pair donors, they're acting as Lewis bases, when you guys get to that section, if you haven't yet in terms of my videos about acid and base identification, realize that there are different types of acids and bases out there here, we're talking about lewis acids and bases and here we're talking about how they donate electron pairs. Alright, so our complex ion, when we're looking at this right here, this is our complex while this our coordination complex or connect coordination come on compound, it breaks up into a complex ion and then also in order to maintain the overall neutrality of the compound, we use counter ions. Alright, so how you look at this thing and figure out how it breaks up? Well, we're going to say here, first is we should calculate what is the oxidation number of nickel here, here, we're going to say that nickel is x, ammonia has no charge cuz it's a neutral covalent compound. So, oxidation is number 0, chlorine here is minus 1. So, it'd be x plus four ammonias, each one is 0, plus two chlorines, each one is minus one equals 0, the charge of a compound, this cancels out, so this is x minus two equals 0. So, x equal plus two. So, what we should realize here, this thing breaks up into two ions, it breaks up into the nickel connected to the four ammonias, we find out that the oxidation number of nickel is plus two, ammonia is 0, so the overall charge of this thing is two plus, then we have these two chlorines that are also involved, they break free as well, there's two of them, each one minus one. So, we're starting out with our coordination complex on the left side of the arrow, it breaks up into two ions, this part here, this is our complex ion, and again your complex ion is made up of basically two major things, it's made up of your metal cation and then it's made up of your ligands, ligand or ligands, here our ligands are ammonia, which are neutral but there could have also been ammonia with bromine, there could be more than one type of ligand attached to the transition metal in the middle and then here these Cl minus is here, there's two of them. So, overall, this is minus two. So, this negative two charge overall would cancel up this positive two charge overall and that's why our coordination complex in the beginning was neutral. So, since it's giving us a neutral compound when they're together these represent our counter ions.
So, that's how we basically look at something and we break it down into its components, I know there's a lot involved here, this complex ion is made up of a bunch of things together, not as simple as your basic ionic compound, but it's still based on the principles of them. Alright, so now that we've talked about the coordination complex and how it's made up of a complex ion and the counter ion, we have to talk about the coordination number involved. Now, we're going to say the coordination number is the number of ligand atoms bonded to the central metal cation and here basically the number of ligand atoms that are usually attached to the central element, usually two, four or six ligands are attached to the central metal these are the most common numbers, there is on rare occasion very rare eight but here we don't worry too much about that one, that one is way out there in terms of the different types of coordination compounds we can form. So, remember your coordination number is the number of ligands attached to your metal cation. Now here, we're going to say the coordination number is based on the size or charge of the metal cation and it's also based on its electron configuration because remember these coordination complexes are paramagnetic. So, we have electrons that are not paired up, the more unpaired electrons you have the greater the chance of the number of bonds you can form because remember forming bonds means you gain electrons, electrons that you can use to pair up with your unpaired electrons and we're going to say here, we're going to say that the most common coordination number is six, however two and four also common, six just happens to be the most likely number for a lot of these coordination complexes. So, if we're talking about this we have to talk about geometries. So, we're going to say the types of the geometries allow are based on the coordination number of the central metal ion, the first one is linear in terms of its geometry, this happens when you have a coordination number of two. So, basically we have our metal cation in the center and it's connected to two ligands, these ligands could be negative or they can be neutral. So, a good example here is we could have copper and it could be connected to, maybe two bromines, and we should realize here, that when it comes to the complex ion, we draw brackets around it, so for the complex I'm always going to put brackets around it. Now here, these next to both happen when we have four ligands connected to the metal cation, okay? So, four here could be either tetrahedral or there could be a square planar or planner, how do we determine, which one is going to dominate? Well, basically we look at the metal cation and if the metal cation has an electron configuration that ends with d-10, for example, zinc, zinc is argon 4s2 3d10, if it has a d-10 configuration then it will be tetrahedral, if it has at d-8 configuration for example, if we want to think of d-8 we could think of nickel because nickel is argon 4s2 3d8 then it will be square planar or planner. So, good examples here, I could do zinc connected to four hydroxides, overall charges minus 2, and how am I coming up with the overall charge? Well, we know that zinc is a type 1 metal, it's always plus 2 but we have four hydroxide ions, each one is minus one. So, if you think about it you have plus two here, you have minus four here, really because there's each one is minus one, that's what a charge overall is minus two and then here you could just do nickel with those four hydroxides as well, also minus two overall. Notice how again I'm putting brackets around a complex ion, because that's what you're supposed to do to write it correctly, and then octahedral is when you have six around here. So, we could just do cobalt, we can use now a neutral ligand if you want. So, you have six of them, you can say overall this charges plus three. So, again these are the most common types of coordination numbers here, eight is possible but you're not going to see it within your book so don't worry too much about that, 6 is the most common one, the ligands can be the neutral or negatively charged and as we go more and more into a coordination complex, we'll be taking a look at different shapes they can have, the different types of isomers that exist as well as how do you name these different types of molecules, it's a bit different from the naming of normal ionic compounds and covalent compounds. So, it's going to be a whole new set of rules that you have to remember. So, just remember coordination complex looks at the complex line with its counter ion, in the complex ion we have our metal ion in the center connected to ligands, which can be either negative or neutral, the counter ion is just anions that are used to balance out the overall charge of the complex ion that overall were neutral. Remember these principles when looking at coordination complexes.
Example #1: Coordination Compounds
So here they ask to find the geometry of the following complex ion. So, here we have zinc connected to four ammonia molecules, these are the ligands, and the overall charge is plus 2. Now, remember zinc its electron configuration is argon 4S2 3d10 and remember, we said that if you have a d-10 configuration your shape would be tetrahedral. So, here the shape, geometry would be tetrahedral and if we wanted to draw this out correctly, we have zinc in the center. Remember, tetrahedral means you are connected to four groups. Now, traditionally you probably see it like this NH3 here, NH3 here, NH3 here and here it's the n that's connected and that n is connected to three hydrogen's. So, here I draw it backwards to show the connection is between the zinc, central element, and the nitrogen because it's the nitrogen that has the lone pair being used to make the connection. Now this is not the best way to draw tetrahedral, technically the best way to draw it will be drawing it like this, we still have a zinc in the center, we still have one ammonia molecule up here, but then the remaining three, we draw kind of like that, this would be the more correct way of drawing this connection and here because it has a charge we put it in brackets and they'd be a plus 2 charge on the outside. So, that'd be the correct way of drawing it, we've seen this one, let's see if you can draw the next one, pause the video really quickly, attempt to do it on your own then come back and see if your answer matches up with mine.
Example #2: Coordination Compounds
Alright guys, hopefully you pause the video when you attempted to do it on your own, this one is fairly simple, we have gold here but there are only two ligands two bromide ions. So, if your coordination number is 2 then your only shape can be linear, it's when the coordination number is 4 that we have to decide, does it follow a tetrahedral shape or does it follow a square planar or planner shape. So, here we have gold in the center. Remember, linear 180 degree in terms of its bond angle, okay? So, we put this in brackets and the negative charge on the outside, so that would be the way you should have drawn it and remember because there's a charge you put brackets around it ,and the way the formula is presented to us should have brackets in it. Now, now that we've done this one try to attempt to do this practice one on your own, come back and see how best I draw. Now, remember just like tetrahedral there is an ideal way to draw this particular shape, draw it first the way you would like to see it and then we'll see if your shape matches up with my shape. So, good luck guys.
Practice: Determine the geometry for the following complex ion: [Cr(NH3)4Cl2]2+
Concept #2: Ligands
Hey guys. in this new video we're going to take a look at ligands. So, we're going to say a ligand can be thought of as just simply a lewis base because remember, a Lewis base is an electron pair donor. So, what happens here is that the transition metal in the middle is a cation. So, it's positive so it can easily accept negative electrons, the metal cation in the center is the Lewis acid, which is an electron pair acceptor, the ligand, which can be either neutral or negative donates its lone pairs to the central metal. Now, I'm going to say ligands can be characterized by the number of elements in the molecule that can donate a lone pair. Now, we're going to say here, that these compounds use their lone pairs to grab on to these metal cations and therefore they're referred to as chelating agents. So, chelating agents. Now, chelating agents, we're going to say that Chela is Greek and it means crabs claw and basically what happens here is that the lone pair kind of act like teeth that kind of chump into the metal cations, when it's donating its lone pairs actually surrounding and holding on to that metal cation performing a strong bond. So, that's where the term chelating agent comes from. Now, we're going to say ligands that possess only one element, able to donate a lone pair, referred to as monodentate ligands, when we say monodentate, monodentate means 1-2, because again, we think of them as claws or teeth that connect to the metal ion, common ones here we have water of course only one lone pair is being donated to form the bond not both, here X, X just stands for halogen. So, we're talking about chlorine, chlorine, bromine and iodine, technically since they're negative we're talking about fluoride chloride, bromide and iodide. Now, when you get to organic for those of you are brave enough to go into organic you'll learn that X is just the default symbol to represent all halogens. So, take note of that, when you see x in chemistry religious refers to some type of halogen, next we have our cyanide ion. Now, nitrogen also has a lone pair but it's the carbon that possesses the negative charge, it's the carbon that's going to be attacking with its lone pair, here we have our hydroxide ion, we have ammonia.
Now here, this one is basically similar to cyanide ion except now we have the possession of a sulphur in this compound, we're going to say here that when it comes to the structure this is called your thiocyanate ion, and here's the thing, if the sulfur or the nitrogen can be the attacker, not both it's one or the other. So, that's why it's monodentate here, this is our nitrite ion and we should realize here about our nitrite ion is that it's the nitrogen or the oxygen that can donate and technically your nitrite ion has resonance, we commonly talk about Lewis dot structures, certain compounds can do resonance. So, another way I could have drawn this is I could have slowed down that nitrogen in the center but it could have been the oxygen that's on the right that was single bonded and then the oxygen on the left now is double bonded. So, here it's still one or the other, the nitrogen or the oxygen. So, these are monodentate ligands here, ligands that possess two elements, able to donate a lone pair referred to as bidentate, we don't say didentate, didentate does not exist, okay? It's bidentate. So, here bi means 2, what we have here is we have our oxalate ion, which will sometimes see written as C204 2-, so this is one way we could have drawn it, this also is resonance as well, we could draw it a different way, we could draw still those carbons in the center but now it's the oxygens on top that are both single bonded and then the oxygens on the bottom are double bonded, okay? Or, we could just mix and match, maybe this one here single bonded and maybe that one there single bonded, so it all depends how you wanted to look at it, we just realized oxalate ion is also resonance stabilized, here this other one that we have here, this is called ethylene diamine, let me take myself out of the image guys. So, that's ethylene diamine ethylene just means you have two CH2 groups and we're saying a mean here, we're talking about nitrogen's. Remember, nitrogen likes to make three bonds, it's each one only making one bond right now. So, each one will have two hydrogen's involved, they're in group 5 ay so they have 5 valence electrons. So, that's where the lone pairs come from. Alright, here we're going to say that bidentate and Polident 8 ligands because they have more than one element with a lone pair they can form rings, when they attach to the metal cation. So, we're going to say that bidentate and polydentate ligands, which will be on the next page, they form rings, in the complex ion, later on we'll be seeing examples of this. Now here, if our ligand possess more than two elements that are able to donate a lone pair, then they're referred to as Polydentate ligands poly, meaning many 2. So, here we have triphosphate. Now, notice although these oxygens here are also negative, it is not them that donate their electrons, it's these three here that donate the electrons, then we have here diethylene. So, di means two, remember ethylene means ch2, ch2. So, here goes one ethylene here goes a second ethylene, that's why it's diethylene, tri means three, three, what three amines, here goes an amine, here goes an amine and here goes an amine. Notice that here at the end, nitrogen again wants to make three bonds, each one is making one bond to a carbon, meaning that they need to make two more bonds. So, that's why each one has two hydrogen's, this one here in the center though is already making two bonds, one to this carbon and one to this carbon so it only needs one more hydrogen to have three. Now, finally this last one here, this is called ethylenediaminetetraacetate ion or simply EDTA. So, let's look at the naming ethylene, ethylene because, we're going to say it's ethylene because we have this ethylene part right here, here goes our ethylene, diamine, diamine means we have two nitrogens and then tetra means four, tetra acetate, here's our acetate ion this part right here is acetate, and how many of them are there? there's four of them. Now, technically here there are six groups that can donate electrons, so this is technically a hexadentate ligand, but here's the thing, usually what happens is it's these four, since they're negatively charged they have an excess of electrons. So, they're the ones most likely to donate electrons, okay? These nitrogen's here are neutral. So, they're less likely to donate their electrons, they could but more than likely it'll be the acetate oxygens that are negative that are donating their electrons. Now, EDTA, you may say I've never heard of EDTA but the thing about EDTA is EDTA is found in a lot of everyday products, you'll find it in processed fruits and vegetables even if you look at certain shampoos in the in the grocery store, if you look at certain ones they have EDTA in them as an active ingredient, you'll see it in mayonnaise, you'll see it in salad dressings, your seat and sweeteners, it's in a lot of things that we either use to clean ourselves or to clean our homes or to even eat, here's the thing, why this is found on everything? basically, for example, if you have a can of fruits, right? what happens here is that during the process of packaging certain foods, there are trace amounts of metal shavings that can be found in the food, which isn't good, so how do you make sure you basically soak up these metal ions so that we don't get sick from them? you put a little bit of EDTA in these processed cans, what happens here is that the EDTA, it binds to any traces of metals left behind in the food and what it does here is it gets rid of those metal ions from your food and at the same time it stops your food from spoiling too quickly, so it acts kind of like a preservative, it also acts as a way of protecting us from free-floating metal ions. So, there's huge science involved in terms of this compound EDTA, it's not in a lot of things. So, it's a very, very, very important ligand that you may not have heard of until today but it's found in almost everything that we use. So, remember we're talking about ligands here, ligands are just simply lewis bases, they connect to the metal ion in order to form our complex ion. Remember, your complex ion is made up of your metal ion in the center, which acts as a lewis acid and then your ligand that's around it, these ligands can either be negatively charged or they can be neutral as in case of ethylenediamine. So, just remember the different types, we have mono, bi and poly, EDTA being the most popular polydentate ligand.
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