Understanding Electronic Structure of Atoms

Hi everybody, my name is Andrei and welcome back to Med School EU. Today we are going over the chemistry iMath specifications and we are going to talk about the electronic structure of atoms of different elements, meaning that we are going to talk about the electronic configuration. So first of all, what is electronic configuration and the definition of it is that it is an arrangement of electrons in an atom. So in the previous video we talked about the atomic structure, we talked about atoms having a nucleus that contains protons and neutrons and the electrons are going to be flowing around the nucleus in its shells or orbits or whatever you like to call them in your country. And now we're going to talk about the arrangement of the electrons, where they are positioned, where do they typically vibrate, where do they typically go, and what kind of implications that has on the atom and how we are going to label that and depict that, meaning the electronic configuration. So first thing we're going to do is talk about the principal quantum number. And the principal quantum number represents the shells. So if we've got n1 or n is equal to 1, that represents the first shell that's right here. So this is n equals 1. This shell is going to be n equals 2 and so on. So atoms are going to have multiple shells. Some atoms will just have one shell like hydrogen, for example, and helium. The second group of elements are going to have two shells and the third group are going to have three shells and so on. But each shell could have a different number of maximum electrons that it can hold because remember electrons are negatively charged so if there's too many of them clustered in one shell or too many of them clustered in one region they are going to repel each other so electrons like to spread out they like to spread out onto their different shells based on the energy that they have And that first shell is going to be represented by the principal quantum number labeled by the letter N. And so the first shell could have a maximum of two electrons. And this is why only hydrogen and helium are represented to have just the first shell. Now the second shell would have a maximum of eight electrons. The third shell has a maximum of 18 electrons. And the fourth shell has the maximum of 32 electrons and so on. So the further you go from the shell, the higher the energy of the electrons. Because it takes more energy for them to get away from the positively charged nucleus than to be nice and close to the positively charged nucleus. So therefore those electrons will be highly charged, the ones that are on the outer. Shell or the valence shell as we call it. Now another thing to understand is that each of these shells, they're not going to be like orbits just like we have the sun and the planets circling around it in an orbit that never changes. That's not typically what happens with electrons. I mean electrons are going to be going around in all kinds of different directions just jumping around like this all around this one nucleus that the atom has. It's not going to be stationary going in a... circle that is constantly like that. So this is not an accurate depiction of how electrons behave. However, a more accurate depiction of how electrons behave is determining in which spots they will reside in for most of the time. Not all the time, but most of the time. And that is done by quantum sub-shells. And I'm not going to draw them. They're a little bit complicated to draw because some of them are in 3D and 5D. shapes. So I'm going to show you an image of that. So this is what these quantum subshells look like. Now the first one is called S. This first one that's just a sphere. So this is the S subshell. The next ones are going to be in these 3D depictions. So they're coming from three different orientations. This is going to be the P subshells. and as you can see there's three different ones we've got there. This is the d subshell in the 5d five dimensions and the f subshells. So the energy levels will be going from s to p so p is higher than than s but p is going to be lower than d in terms of energy and d will be lower than f f containing highest level of energy and what these subshells really do is they hold electrons and the the s subshell is going to hold a maximum of two electron within its sphere now that doesn't mean that these electrons cannot escape or it is an actual physical sphere These are not physical spheres. These are not actual things that there's barriers and all these bubbles all around the nucleus. That's not what's happening. This is basically just a depiction of taking millions of photographs of electrons flying by and most of them are going to appear in that circle or in this sort of sphere or in this orientation and that is how we determined these sort of subshells with with just an experiment like that. Now if we are talking about the the p sub shell well they're going to have one electron each in each of these little bubbles the yellow bubble and this bubble here each one of them has is going to have its own number of electrons so it's going to be 2 for the s it will be 6 for the p it's going to be 10 for the d and in F is going to be 14 electrons a maximum of 14 electrons that could be held within those spheres. Now just knowing the quantum subshells and knowing the principal quantum number here's something we need to understand. So in the first subshell when or in the first for the first principal quantum number meaning the first shell or the first orbit we are going to have two electrons, a maximum of two electrons in the s subshell. So only the s subshell will be within the first orbit and that is the reason why the first one can only hold eight electrons. Now in terms of the second we are going to have eight electrons as mentioned previously. Now why is that? Well because these eight electrons are going to come from two Electrons from the S subshell. So we're going to have another S subshell for the second orbit. And on top of that, they will be clustered with the six electrons from the P subshell. And remember, the P subshells, there's three of them, and each one has two electrons. This is why, altogether, we're going to have a maximum of eight electrons in the second shell. Now moving on to the third shell we are going to have 18 electrons in the third shell. Why is that? Because we'll have two from the S subshell and combined they will hold a maximum of 18 electrons in the third shell. Now let's talk about how we're going to represent them in terms of the electron configurations. How can we represent elements and atoms with different number of electrons? with a configuration. Now the configuration will be as follows. So this represents the number of electrons. The s right here is going to of course represent the sub shell that is involved and the number in front of the sub shell will represent the principal quantum number. So in essence that principal quantum number simply represents the number which shell we're talking about or which orbit we are specifically referring to. So which orbit we refer to, the subshell that we're referring to, and the number of electrons in that specific subshell. That is going to be the depiction of the electron configuration that gives the exact location of the electron that we are talking about. Now when we put all of these things together, that gives us an electron configuration. for different elements. So let's show a diagram that is going to depict the energy levels of each of these configurations. So here's the kind of representation that will be used. This is kind of like a guide on the energy levels and you're going to go across like this. Of course following the arrows as soon as the arrow ends you are going to go to the next arrow that will provide you with the sort of instructions on how to go. So it's going to be 1s then 2s then 2p, 3s, 3p, 4s, 3d, 4p, 5s, and so on. And that's going to represent the energy levels that will be covered. Because obviously the electrons want to reside in the lowest energy levels possible. But if those ones are already occupied by other electrons, then they're of course going to take the next one. They're not going to just automatically start with one of the furthest ones because they're high energy. No, electrons don't really gain energy all of a sudden. They are going to reside in their lowest energy based on the availability if it's vacant or if it's taken by some other electrons already. So let's do as an example, we're going to run an example of sodium. Now if we look at your periodic table, sodium is the 11th element. So it means it has 11 electrons. Now how do we depict that? Well, we are going to have the 1s. two because the first subshell is going to be completely filled it's got its two electrons then we're moving on to the next one's got 2s and it will also have two because remember s can maximum have a number of two and there's 11 electrons to do so we're at four electrons so far next we're gonna have a 2p6 so now we're at 10 electrons remember p can hold six And the next one is going to be a 3s but this one will only have one as we only need 11 electrons here. So this is an actual electron configuration of sodium. Now if we are going up to elements that are bigger than 18 so have more than 18 electrons. Sometimes they're going to be represented by their noble gases. And what do I really mean by that? Well if we have an element let's say manganese that this element is is number 25 so it's got 25 electrons more than 18. Now typically it will be represented here as we have argon. Now why is that argon has 18 electrons and argon is a noble gas that has a full fulfilled octet meaning it has a maximum number of electrons in its valence shell. and therefore it's going to be used as kind of like a shortcut to get a representation of your electron configuration for an element that is higher than 18 electrons. Because imagine doing something that's like 100 or 112 electrons. You're not going to go and draw this whole thing forever, right? You are probably going to use the next available noble gas that is out there. So after That is going to be 1s2, 2s2, 2p6, 3s2, then it's going to be 3p6, 4s2 and we stop at the d because d can have 10. So here we've got 18, then it's going to be 3d5 and 4s2. Next, we might be used to explain the ions as well. So for example if we have an ion of calcium which is 2 plus we could have a representation that that goes like this. I mean 2 plus meaning it lost two electrons because it's plus electrons are negatively charged and if you lose two negative charges you're going to have a positively charged overall. That's something we will discuss in greater detail. However here we no longer have 20 Electrons, we have 18 electrons just like argon and so it's going to be 1s2 2s2 2p6 3s2 3p6 and that's going to be the end of that. And so what does this really give us? Well this is the same as argon because the ion for calcium you know, will lose two electrons to become stable and it's going to represent the electron configuration of argon. Now another quantum number that will represent which electron we're talking about is going to be the electron spin. So some electrons are going to spin clockwise and some will spin counterclockwise. And so We are going to represent them with a box form representation. So let's take a look at two examples here. And the reason they spin in opposite directions is again because of of electron to electron repulsion because electrons do not attract, they repel each other since they're both negatively charged. So let's talk about nitrogen here. configuration of nitrogen with the box form. Nitrogen has seven electrons and so let's label the electrons present. This first one is going to be an electron up and down. So it's going to be fully filled. It's the first, you know, the 1s2. Here's going to be your 2s2 that will have two electrons, one spinning clockwise, one spinning counterclockwise. And this will be our 2p3. So how this is going to be structured is that they're going to be completely spread out. The reason for that is because of electron... to electron repulsion so the electrons want to like I mentioned spread out as evenly as possible so they do not interact with each other and as you as you can see that they're going to spin in the same direction as well now let's do another example with oxygen here that has eight electrons now of course here we got our 1s2 so one going clockwise one counterclockwise 2s2 same thing and we've got the 2p4 so we're going to have first fill in our clockwise or our counterclockwise direction for all of these subshells for all of the angles of the p subshells all of the different structures and then finally this one's going to be completely filled because of the extra electron. So this concludes our video for today on the electronic structure of atoms. In the next video we are going to take a look at the periodic trends.

So in the previous video we talked about the atomic structure, we talked about atoms having a nucleus that contains protons and neutrons and the electrons are going to be flowing around the nucleus in its shells or orbits or whatever you like to call them in your country. And now we're going to talk about the arrangement of the electrons, where they are positioned, where do they typically vibrate, where do they typically go, and what kind of implications that has on the atom and how we are going to label that and depict that, meaning the electronic configuration. So first thing we're going to do is talk about the principal quantum number. And the principal quantum number represents the shells.

So if we've got n1 or n is equal to 1, that represents the first shell that's right here. So this is n equals 1. This shell is going to be n equals 2 and so on. So atoms are going to have multiple shells.

Some atoms will just have one shell like hydrogen, for example, and helium. The second group of elements are going to have two shells and the third group are going to have three shells and so on. But each shell could have a different number of maximum electrons that it can hold because remember electrons are negatively charged so if there's too many of them clustered in one shell or too many of them clustered in one region they are going to repel each other so electrons like to spread out they like to spread out onto their different shells based on the energy that they have And that first shell is going to be represented by the principal quantum number labeled by the letter N. And so the first shell could have a maximum of two electrons.

And this is why only hydrogen and helium are represented to have just the first shell. Now the second shell would have a maximum of eight electrons. The third shell has a maximum of 18 electrons.

And the fourth shell has the maximum of 32 electrons and so on. So the further you go from the shell, the higher the energy of the electrons. Because it takes more energy for them to get away from the positively charged nucleus than to be nice and close to the positively charged nucleus. So therefore those electrons will be highly charged, the ones that are on the outer.

Shell or the valence shell as we call it. Now another thing to understand is that each of these shells, they're not going to be like orbits just like we have the sun and the planets circling around it in an orbit that never changes. That's not typically what happens with electrons. I mean electrons are going to be going around in all kinds of different directions just jumping around like this all around this one nucleus that the atom has. It's not going to be stationary going in a...

circle that is constantly like that. So this is not an accurate depiction of how electrons behave. However, a more accurate depiction of how electrons behave is determining in which spots they will reside in for most of the time. Not all the time, but most of the time.

And that is done by quantum sub-shells. And I'm not going to draw them. They're a little bit complicated to draw because some of them are in 3D and 5D. shapes. So I'm going to show you an image of that.

So this is what these quantum subshells look like. Now the first one is called S. This first one that's just a sphere.

So this is the S subshell. The next ones are going to be in these 3D depictions. So they're coming from three different orientations.

This is going to be the P subshells. and as you can see there's three different ones we've got there. This is the d subshell in the 5d five dimensions and the f subshells. So the energy levels will be going from s to p so p is higher than than s but p is going to be lower than d in terms of energy and d will be lower than f f containing highest level of energy and what these subshells really do is they hold electrons and the the s subshell is going to hold a maximum of two electron within its sphere now that doesn't mean that these electrons cannot escape or it is an actual physical sphere These are not physical spheres. These are not actual things that there's barriers and all these bubbles all around the nucleus.

That's not what's happening. This is basically just a depiction of taking millions of photographs of electrons flying by and most of them are going to appear in that circle or in this sort of sphere or in this orientation and that is how we determined these sort of subshells with with just an experiment like that. Now if we are talking about the the p sub shell well they're going to have one electron each in each of these little bubbles the yellow bubble and this bubble here each one of them has is going to have its own number of electrons so it's going to be 2 for the s it will be 6 for the p it's going to be 10 for the d and in F is going to be 14 electrons a maximum of 14 electrons that could be held within those spheres. Now just knowing the quantum subshells and knowing the principal quantum number here's something we need to understand.

So in the first subshell when or in the first for the first principal quantum number meaning the first shell or the first orbit we are going to have two electrons, a maximum of two electrons in the s subshell. So only the s subshell will be within the first orbit and that is the reason why the first one can only hold eight electrons. Now in terms of the second we are going to have eight electrons as mentioned previously. Now why is that? Well because these eight electrons are going to come from two Electrons from the S subshell.

So we're going to have another S subshell for the second orbit. And on top of that, they will be clustered with the six electrons from the P subshell. And remember, the P subshells, there's three of them, and each one has two electrons.

This is why, altogether, we're going to have a maximum of eight electrons in the second shell. Now moving on to the third shell we are going to have 18 electrons in the third shell. Why is that? Because we'll have two from the S subshell and combined they will hold a maximum of 18 electrons in the third shell. Now let's talk about how we're going to represent them in terms of the electron configurations.

How can we represent elements and atoms with different number of electrons? with a configuration. Now the configuration will be as follows. So this represents the number of electrons.

The s right here is going to of course represent the sub shell that is involved and the number in front of the sub shell will represent the principal quantum number. So in essence that principal quantum number simply represents the number which shell we're talking about or which orbit we are specifically referring to. So which orbit we refer to, the subshell that we're referring to, and the number of electrons in that specific subshell.

That is going to be the depiction of the electron configuration that gives the exact location of the electron that we are talking about. Now when we put all of these things together, that gives us an electron configuration. for different elements. So let's show a diagram that is going to depict the energy levels of each of these configurations.

So here's the kind of representation that will be used. This is kind of like a guide on the energy levels and you're going to go across like this. Of course following the arrows as soon as the arrow ends you are going to go to the next arrow that will provide you with the sort of instructions on how to go. So it's going to be 1s then 2s then 2p, 3s, 3p, 4s, 3d, 4p, 5s, and so on.

And that's going to represent the energy levels that will be covered. Because obviously the electrons want to reside in the lowest energy levels possible. But if those ones are already occupied by other electrons, then they're of course going to take the next one.

They're not going to just automatically start with one of the furthest ones because they're high energy. No, electrons don't really gain energy all of a sudden. They are going to reside in their lowest energy based on the availability if it's vacant or if it's taken by some other electrons already. So let's do as an example, we're going to run an example of sodium. Now if we look at your periodic table, sodium is the 11th element.

So it means it has 11 electrons. Now how do we depict that? Well, we are going to have the 1s.

two because the first subshell is going to be completely filled it's got its two electrons then we're moving on to the next one's got 2s and it will also have two because remember s can maximum have a number of two and there's 11 electrons to do so we're at four electrons so far next we're gonna have a 2p6 so now we're at 10 electrons remember p can hold six And the next one is going to be a 3s but this one will only have one as we only need 11 electrons here. So this is an actual electron configuration of sodium. Now if we are going up to elements that are bigger than 18 so have more than 18 electrons.

Sometimes they're going to be represented by their noble gases. And what do I really mean by that? Well if we have an element let's say manganese that this element is is number 25 so it's got 25 electrons more than 18. Now typically it will be represented here as we have argon. Now why is that argon has 18 electrons and argon is a noble gas that has a full fulfilled octet meaning it has a maximum number of electrons in its valence shell.

and therefore it's going to be used as kind of like a shortcut to get a representation of your electron configuration for an element that is higher than 18 electrons. Because imagine doing something that's like 100 or 112 electrons. You're not going to go and draw this whole thing forever, right? You are probably going to use the next available noble gas that is out there.

So after That is going to be 1s2, 2s2, 2p6, 3s2, then it's going to be 3p6, 4s2 and we stop at the d because d can have 10. So here we've got 18, then it's going to be 3d5 and 4s2. Next, we might be used to explain the ions as well. So for example if we have an ion of calcium which is 2 plus we could have a representation that that goes like this. I mean 2 plus meaning it lost two electrons because it's plus electrons are negatively charged and if you lose two negative charges you're going to have a positively charged overall. That's something we will discuss in greater detail.

However here we no longer have 20 Electrons, we have 18 electrons just like argon and so it's going to be 1s2 2s2 2p6 3s2 3p6 and that's going to be the end of that. And so what does this really give us? Well this is the same as argon because the ion for calcium you know, will lose two electrons to become stable and it's going to represent the electron configuration of argon.

Now another quantum number that will represent which electron we're talking about is going to be the electron spin. So some electrons are going to spin clockwise and some will spin counterclockwise. And so We are going to represent them with a box form representation.

So let's take a look at two examples here. And the reason they spin in opposite directions is again because of of electron to electron repulsion because electrons do not attract, they repel each other since they're both negatively charged. So let's talk about nitrogen here. configuration of nitrogen with the box form. Nitrogen has seven electrons and so let's label the electrons present.

This first one is going to be an electron up and down. So it's going to be fully filled. It's the first, you know, the 1s2.

Here's going to be your 2s2 that will have two electrons, one spinning clockwise, one spinning counterclockwise. And this will be our 2p3. So how this is going to be structured is that they're going to be completely spread out.

The reason for that is because of electron... to electron repulsion so the electrons want to like I mentioned spread out as evenly as possible so they do not interact with each other and as you as you can see that they're going to spin in the same direction as well now let's do another example with oxygen here that has eight electrons now of course here we got our 1s2 so one going clockwise one counterclockwise 2s2 same thing and we've got the 2p4 so we're going to have first fill in our clockwise or our counterclockwise direction for all of these subshells for all of the angles of the p subshells all of the different structures and then finally this one's going to be completely filled because of the extra electron. So this concludes our video for today on the electronic structure of atoms. In the next video we are going to take a look at the periodic trends.

Transcript for:Understanding Electronic Structure of Atoms

Transcript for:
Understanding Electronic Structure of Atoms