|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 & 30mins||0% complete||WorksheetStart|
|Ch.5 - Gases||3hrs & 47mins||0% complete||WorksheetStart|
|Ch.6 - Thermochemistry||2hrs & 28mins||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: Analyzing the Transition Metals
Guys is what we are going to say here that the transition metals represent elements found basically in the d block of the periodic table. Now, let me take myself out of the image and let's talk a little bit about this periodic table. So, we know that we have our representative elements, these are elements in groups 1a 2a all the way to 8a, they're your representative or main group elements. Now, inside of this pit here we have our transition metals and what we need to realize here is that we also have within this pit our lanthanides, which is this row and your actinides. Now, we call these two roles your inner transition metals, why they call the inner transition metals? because if you take a look at this periodic table. Notice how the number goes from 56 and then all of a sudden to become 71, what happens here is we go 56 then 57 all the way to 70 and then we come back up here to 71, which means between these two elements, we actually find this entire row and the same thing can be said for, we go 88, 89 all the way to 102 then we come back up here to 103, okay? So, here your actinides are found between these two. So, they're found within the transition metal pit. So, that's why they're called inter transition metals.
Now, we're going to say here that this pit is also referred to sometimes as our B elements, so notice that our main group elements are 1a, 2a, 3a, all the way to 8a in this pit though we can say that this is 3B, bB, 5b, 6b, 7b. Now, these three here collectively they're 8b, we're going to say, this is 8b1, 8B2, 8B3, this is actually 1b and 2b, it's just a different system of classifying these different types of groups that exist for the transition metals, and in some periodic tables you may actually see it written as basically, this is Group 1, 2 then they actually count this as 3, 4, 5 all the way to 12 and this here would be 13, 14, 15, 16, 17 and 18. So, you might see this classification as well. Now, here we're going to say where is the main group elements show similar chemical behaviors because of their valence electrons, transition metal similarities are treated differently because here we don't look at their group number to determine the valence electron number, here we're going to say transition metals show great chemical similarities in both of their horizontal periods. So, remember you're periods all your rows and it also shows similar chemicals from properties within their vertical groups or families, okay? So, remember your group is your column, it's your family also. So, one of the reasons why transition metals are different is because we have our d orbitals and our f orbitals. So, we're going to say in the gradual addition of electrons to transition metals new electrons are added to their inner core electrons. So, we're adding them to our d or f orbitals, and these orbitals don't participate in chemical bonds and we're going to say for transition metals each transitional electron is added to the d-block orbitals while for lanthanides and actinides they're added to the f-block orbitals. So, we revisit things such as electron configuration, we'll be talking about how adding electrons how it affects your tone and shape, how it affects your atomic size, things called effective nuclear charge, a lot of common types of periodic trends that we've talked about in the past but now strictly applying them to transition metals. So, here we're just opening up the door, entering into the realm transition metals before we go more and more in depth in terms of different things we haven't seen before.
Concept #2: Properties of Transition Metals
Hey guys. In this video we're going to take a look at some of the common types of properties dealing with transition metals. So, here we're going to say like most main group elements, transition metals possess similar physical properties. So, some of the physical properties that they share are common to many types of metals, for instance their luster or shine metals in general tend to be shiny, they also have high densities, we're going to say here they're good at electrical and thermal conductivity, we're also going to say they possess high melting points and also hardness. Now, when it comes to conduction there are some metals that work better than others, for instance, we're going to say that silver or AG, possesses the greatest electrical conductivity. So, silver is usually a great way of basically transferring electrons from point A to point B, we're also want to say coming in second would be copper. Now, when it comes to melting point, we're going to say that tungsten or W possesses the highest melting at 3400 degrees Celsius. Now, you'll learn that, well, you probably know from common knowledge that old light bulbs used to have a tungsten filament in them before we move on to more efficient means of using light bulbs, we're also going to say Tungsten's were also used as huge, huge containers for melting of like hot iron or within old school factories, you may actually, if you are in certain towns, you may still see huge furnaces, huge containers made of tungsten, we're going to say that while blank is the only metal that is a liquid at room temperature. So, we say that this is mercury. Now, on the hardness scale, we're going to say here that iron and titanium are strong or hard metals, meaning it would take incredibly high temperatures and a long amount of time for us to be able to melt them. Now, we're going to say here that copper, silver and gold are considered to be soft metals, which makes sense, if you think about it in history coins were usually made up of what, copper, silver and gold, they're malleable easier to melt and fashioned into different types of coins with different types of images on them, that's why a lot of money back in the ancient times was made up of these different types of metals, these soft metals. Now, oxidation states, we're going to say, remember that transition metals possess variable on charges, certain transition metals can have multiple charges, remember, we refer to them as type 2 metals, type 2 metals just means they have more than one charge, manganese for instance has many different charges so it's a type 2 metal, but there are some transition metals that possess only one charge. Remember, that zinc is always plus 2, cadmium is always plus 2, silver is always plus 1, if they don't have variable charges, they only have one charge. So, we refer to them as type 1 metals, type 1 metals have only one charge, type 2 metals have more than one charge. So, transition metals possess variable charges and so the use of what we call reducing agents, reducing agents have to be used in order to identify which particular charge we're dealing with, these reducing agents donate electrons and based on the number of electrons of the metal ion accepts we're able to determine the oxidation state of that transition metal, those of you who are lucky enough to decide to go into organic chemistry, you'll learn more and more about the different types of reducing agents that are used every day with inorganic systems.
Concept #3: Electron Configuration of Transition Metals
Hey guys. In this video we're going to take a look at electron configurations of transition metals. So, remember electron configurations are just representation of how exactly do electrons in a way line up within different orbitals of a particular element.
Now here, just a quick recap of what we've done before when it comes to electron configurations, just remember here that our trench metals are here in the d-block, we have our 3d. Notice that we go from 4s to 3d, 5s to 4d, here we're going 6s then we go 4f, which is the lanthanide row then we go 5 d then 6p then we go 7s then we do our actinide row, which is 5f then we go back up to 6d. Now, we say that this is our Aufbau diagram, which is another way of thinking about electron configurations, where we go 1s to 2s to go back around 2p 6 3s2, I personally understand how the aufbau could be useful, but honestly, remembering what the periodic table really looks like in terms of where's your s block, your p block, your d block, your f block, is much easier to do. Now remember, we call there are exceptions that exist with the electron configuration of transition metals. Remember, in the neutral form d4 and d9 are not allowed to exist, but here's the thing, more specifically we're going to say these exceptions are normally observed with only the first row transition metals, meaning that we normally only see these exceptions with chromium and copper, if you take a look at in your textbook you'll see that other elements within these columns don't necessarily follow this exception, you may not see them with certain elements, for the most part it's mainly with first row elements and depending on which Edition you're using you might see the exception with silver here, remember, when in doubt always go by what the professor says more than anything this exception is seen again mainly with first row transition metals, copper and chromium specifically, but you could see happen with silver as well. Now here, write the conducts configuration for the following element. So, we have to do it here for chromium, which has 24 electrons. So, let me take myself out of the image guys, let's take a look at chromium. So chromium, remember we're doing the condensed, so that means argon is first, okay? So, it's argon 4s2 3d4. Remember, and remember why the 3d4 because it's 1, 2, 3 and 4 we land on chromium, d4 and d9 are not allowed to exist. So, what really happens here is that one electron from the s gets handed over to the d, the d orbital. So, it's actually argon 4s1 3d5. Now, we've done stuff like this before but specifically why should we talk about it again? more specifically, we have to basically talk about why does this do it exactly, there are many reasons, let's talk about the different reasons why this happens. So, remember s has one orbital whereas d has 5, okay? So, this is it before we corrected it and this is after we've corrected it. So, for this first one we have two in the s. So, that's one up one down and then we have four. So, it's up, up, up and up. Now, there are number of reasons why this happens, one reason why we hand 1 over from the s to the d orbital is because we're going to say this is allowed to happen because 4s and 3d basically have similar energies their energies are so similar that we can count them as just one big thing, you can combine them all together. So, since their energies are very similar, we can say that they're degenerate orbitals. Remember, the word degenerate means same energy, and remember according to hund's rule, orbitals that have the same energy are first half-filled before they are totally filled. So, one reason why it does this is since they're all the same energy we have filled first, we go up, up, up, up, up, up before coming back around. So, one reason that this happens is again the orbitals are degenerate, 4s and 3ds energies are so similar that we just say, we're going to follow hund's rule and half fill all of them first, that's reason one, second reason is we could say here that d orbitals are most stable when they're half filled or totally filled, if we kept it originally like it was here, we have only, we'd only need one more electron to be totally filled in. So, what do we do, we take one of these electrons and we hand it over so d orbitals can be half filled. So, again first rule hund's rule second rule, d is most stable when it's half filled or totally filled, final reason, we can say the final reason this occurs is because what we call electron-electron repulsion. So, if we take a look at this 4s2. So, we're going to say this is called electron-electron repulsion. Remember, in chemistry opposites attract, there's two electrons within orbital, both of them are negatively charged, like charges repel each other, okay? And because they're like charges and they repel each other this repulsion causes an increase in our energy. So, it's very unstable to have this 4s2 totally filled in, when we can so easily hand 1 over to the d, we not only make d more stable by making the d orbitals half filled but we also make the s more stable by getting rid of this electron-electron repulsion. So, these are the three main reasons why this happens, one again because these orbitals are seen as degenerate, they all have the same energy. So, hund's rule, we half fill them first, second reason is that d orbitals are most stable when they're half filled are totally filled and third reason because of electron-electron repulsion, it'd be best to hand 1 over to the d orbital thereby stabilizing it and another way stabilizing the s orbital at the same time. So, there's a lot of reasons but those are the most compelling reasons why this happens. So, remember the exception that we've talked about in the past, apply to orbitals now, specifically about transition metals.
Example #1: Electron Configuration of Transition Metals
Hey guys. In this video we're going to take a look at figuring out the electron configurations of certain types of transition metals. So, here for this first example, we have to figure out the condensed electron configuration of the following element. So, we're looking for the electron configuration of tungsten. So, if we take a look at a basic periodic table, we see that tungsten is right here. So, remember, we're going to do the condensed electron configuration. So, we're going to say here that the last noble gas I passed before I get to tungsten is xenon. So, we have to pass xenon before we get to tungsten. So, we're xenon then remember, this is your sixth row or sixth period so this is 6s. So, this will be 6s2. Now, notice how the number goes from 57 and all of a sudden becomes 72, right? So, what does that mean? That means we're going to have to go here down at this row here, your lanthanides row, right? So, this is the 4f. Alright, because the lanthanides are for f. So, for f 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 4f14 and then we have what? we have back up here, this row so this is 3d, 4d, 5d. So, this would be 5d, 1, 2, 3, 4.
Now, remember for transition metals normally we'd say that d4 and d9 are not allowed to exist but we're going more definite transition metals and remember this rule is mainly followed for the first row of transition metals, we don't apply here to tungsten, this will be the correct configuration for tungsten, we apply those exceptions mainly for chromium and for copper, both in the first row transition metals, you maybe see the exception with silver but for the most or the others we do not see them. So, just take to heart and remember that. So, here we see our answer xenon is s2 4f14 5d4. So, you come back can we write that out. So, xenon 6s2 4f14 5d4 would be our answer. So, remember that exception as many for first row transition metals. Now, this next example here, what I want you guys to do is try to do it on your own, pause the video try to see if you can get on your own and come back and see if you get the same answer, here we're dealing with the ion form. So, just make sure that you do the neutral form first and then recall, what does 4 positive mean in terms of the new electron configuration that results. So, once you do that come back and see how I approach this example question.
Example #2: Electron Configuration of Transition Metals
Alright guys, so hopefully guys pause the video and you attempted to do this question. So, what we want to do here is we want to do manganese neutral. So, manganese neutral, look at your periodic table would be argon, right? It'd be 4s2 3d5. So, that's what you should got for it when it was neutral. Now, 4 plus means are losing 4 electrons. Remember, you always lose your electrons from the outer shell, the outer shell is one with the larger end value. So, here these electrons, these two electrons are in the fourth shell, these 5 electrons are in the third shell, we're going to lose them from the fourth shell first, so we lose two from 4s. So, they'll be gone but we have to lose two more from 3d, losing two more would give us at the end argon 3d3, so this is the answer you should have gotten for manganese 4 plus ion. Now, that you've seen that example try to do the last example on your own, come back and see how I approach that question.
Example #3: Electron Configuration of Transition Metals
Alright guys. So, let's try to do this one, let's do the neutral form of iron first. So, iron when it's neutral it is argon 4s2 3d6, 6 positive means we lose 6 electrons. So, we're going to have left is argon, you lose the first two from the 4s, meaning we have to lose four more. So, we take them from the 3d. So, doing that give us 3d2. So, 3d2 would be your answer here. So, this would be argon 3d2, hopefully you're able to do all these electron configurations. Remember, the exception that we talked about mainly is for first row transition metals, we don't tend to see that the further down we go in terms of those two columns within the transition metal pit, remembering these key exceptions is key to getting the correct answer at the end.
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