- Periodic trends
- Atomic radius trends on periodic table
- Atomic and ionic radii
- Mini-video on ion size
- Ionization energy trends
- Ionization energy: period trend
- First and second ionization energy
- Electron affinity: period trend
- Electronegativity and bonding
- Metallic nature
- Periodic trends and Coulomb's law
- Worked example: Identifying an element from successive ionization energies
- Ionization energy: group trend
How atomic radius is defined, and trends across a period and down a group. Created by Sal Khan.
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- Why going down a group makes a bigger radius? I mean, K has more protons than Li so wouldn't there be a greater pull towards the nucleus and therefor a smaller radius?(15 votes)
- the simplest answer is that Potassium has higher valence energy level (energy level 4) than Lithium (energy level 2), which has greater distance from the nuclear thus has bigger radius(7 votes)
- how many electrons are filled out in each subshell?
is 2-8-8-8-8-8-... right?(5 votes)
Shell 1: 2 electrons
Shell 2: 8 electrons
Shell 3: 18 electrons
Shell 4: 32 electrons
Shells 5 through 7: 32 electrons in any known element, however there are additional orbitals available to hold even more electrons, but there is no element with a large enough atomic number to fill up those slots.(31 votes)
- Why doesn't the electrons around the protons push away the outermost electrons? I thought like charges repel, and since they are closer shouldn't they push away the outer shell?(4 votes)
- Sal said that as you move down the rows, the atoms get larger because of more shells. But aren't the amount of protons and electrons increasing as well, and draw them closer to the nucleus?(3 votes)
- While they amount of protons do increase and draw the electrons in, the way the electrons extend outward in a shell formation makes the atom bigger. The electrons in the outer layer are repelled by electrons in the innermost layers (negative and negative repel), producing a large amount of layers like an onion.(6 votes)
- Of the following element which has the smallest first ionization energy(3 votes)
- Probably an element in the bottom left corner of the periodic table, like Francium.
The first ionization energy is the energy required for an element to lose one of its electrons. Elements found in the first group (Li, Na, K, ...) all have one electron in their outer shell, so they would quite favorably lose that electron. Larger atoms require less energy because the pull on the electrons is weaker the further away the elctron is from the nucleus. So, a weaker pull for larger atoms, thus smaller ionization energy(5 votes)
- Sal said that the two ways to calculate the atomic radius are the van der Wall's radius and the covalent radius.But the the covalent radius is usually smaller than the van der Waal's radius as the electron cloud overlap.so, wouldn't this create the problem in the actual real value(2 votes)
- There is no "actual real value". Atoms don't have discrete borders, so there isn't a specific "real" way to determine where they end. So, these various methods are useful models. They all take the radius as half the distance between atoms but use different circumstances -- which is why they are different.
So, you can sort of think of an atomic radius as being half the distance that another atom can come to an atom before they repel each other.(5 votes)
- This question has no relevance to this topic but it is of relevance to chemistry.
I was wondering that if Mass is defined as the amount of matter in a body and we also say that electron has a mass of 9.11*10(-31) (I hope you read this in the scientific notation), then wouldn't it mean that electrons, protons and neutrons have some kind of a sub-electron or sub-proton or a sub-nucleus? That electrons, protons and neutrons are not the "building blocks"of matter?(2 votes)
- No, it wouldn't mean that.
Protons and neutrons are made of smaller particles called quarks. Quarks and electrons are not made of anything smaller, so far as we can tell.
Also: mass is not really defined as the amount of matter in a body, but that's another discussion entirely.(4 votes)
- So, JUST the outer shell comes closer to the nucleus or all the shells come closer ?(3 votes)
- as the no. of protons are increasing the no. of electrons are also increasing so how is that increasing the coulomb attraction and thus decreasing the atomic size across the period bco in whatever atom we take there is always 1 proton per electron?(2 votes)
- As we move right to left along the same period in the periodic table, the atomic number increases which means the number of protons in the nucleus also increases. To make the atoms neutral, an electron is also added to the atom. All electrons of an atom feel an attractive force to the nucleus because of the positive charge of the protons. However, they also feel a repulsive force from the other electrons since they all possess the same negative charge. But these two conflicting forces, the attractive and repulsive forces, are not equal in magnitude. The attractive force is stronger than the repulsive and so even the outermost electrons feel a net attraction to the nucleus. Which should make sense since atoms exist.
The electrons orbit the nucleus in a variety of orbital shapes and do not completely shield other electrons from the nucleus. The net charge which electrons feel towards the nucleus after subtracting out the repulsions of the electrons is known as effective nuclear charge. The newly added electrons going left to right in a period are being added to the same electron shell as the existing valence electrons. While they do provide some repulsion, they do not block out the entire positive of a proton. So each added proton adds a little more positive charge felt by the valence electrons than is shielded by the additional electrons.
Hope that helps.(3 votes)
- why does the size of the atom become smaller as the protons increase??...because even if the +ve attracting charges have increased the -ve charges have also increased with them...so the effect would remain the same!!(2 votes)
- The outer most shell remains the same but now you have more protons and more electrons, so the pull toward the nucleus with more positive and more negative parts is even stronger! Effect: the outer most shell is attracted to the nucleus and the atom size decreases.
I think you are confusing yourself, thinking that if the ratio of protons to electrons is the same, they should still have the same strong pull when in fact, it is not the charge per electron that matters but the total number of charges involved(3 votes)
Voiceover: Let's think a little bit about the notion of atomic size or atomic radius in this video. At first thought, you might think well this might be a fairly straight-forward thing. If I'm trying to calculate the radius of some type of circular object I'm just thinking about well what's the distance between the center of that circular object and the edge of it. So the length of this line right over here. That would be the radius. And so a lot of people when they conceptualize an atom they imagine a positive nucleus with the protons in the center right over here then they imagine the electrons on these fixed orbits around that nucleus so they might imagine some electrons in this orbit right over here, just kind of orbiting around and then there might be a few more on this orbit out here orbiting around, orbiting around out here. And you might say, "well okay, that's easy to figure out the atomic radius. I just figure out the distance between the nucleus and the outermost electron and we could call that the radius." That would work except for the fact that this is not the right way to conceptualize how electrons or how they move or how they are distributed around a nucleus. Electrons are not in orbits the way that planets are in orbit around the sun and we've talked about this in previous videos. They are in orbitals which are really just probability distributions of where the electrons can be, but they're not that well defined. So, you might have an orbital, and I'm just showing you in 2 dimensions. It would actually be in 3 dimensions, where maybe there's a high probability that the electrons where I'm drawing it in kind of this more shaded in green. But there's some probability that the electrons are in this area right over here and some probability that the electrons are in this area over here, and let's say even a lower probability that the electrons are over this, like this over here. And so you might say, well at a moment the electron's there. The outermost electron we'd say is there. You might say well that's the radius. But in the next moment, there's some probability it might be likely that it ends up here. But there's some probability that it's going to be over there. Then the radius could be there. So electrons, these orbitals, these diffuse probability distributions, they don't have a hard edge, so how can you say what the size of an atom actually is? There's several techniques for thinking about this. One technique for thinking about this is saying, okay, if you have 2 of the same atom, that are- 2 atoms of the same element that are not connected to each other, that are not bonded to each other, that are not part of the same molecule, and you were able to determine somehow the closest that you could get them to each other without them bonding. So, you would kind of see, what's the closest that they can, they can kind of get to each other? So let's say that's one of them and then this is the other one right over here. And if you could figure out that distance, that closest, that minimum distance, without some type of, you know, really, I guess, strong influence happening here, but just the minimum distance that you might see between these 2 and then you could take half of that. So that's one notion. That's actually called the Van der Waals radius. Another way is well what about if you have 2 atoms, 2 atoms of the same element that are bonded to each other? They're bonded to each other through a covalent bond. So a covalent bond, we've already- we've seen this in the past. The most famous of covalent bonds is well, a covalent bond you essentially have 2 atoms. So that's the nucleus of one. That's the nucleus of the other. And they're sharing electrons. So their electron clouds actually, their electron clouds actually overlap with each other, actually overlap with each other so the covalent bond, there the electrons in that bond could spend some of their time on this atom and some of their time on this atom right over here. And so when you have a covalent bond like this, you can then find the distance between the 2 nuclei and take half of that and call that call that the atomic radius. So these are all different ways of thinking about it. Now, with that out of the way, let's think about what the trends for atomic size or atomic radii would be in the periodic table. So the first thing to think about is what do you think will be the trend for atomic radii as we move through a period. So let's say we're in the fourth period and we were to go from potassium to krypton. What do you think is going to be the trend here? And if you want to think about the extremes, how do you think potassium is going to compare to krypton in terms of atomic radius. I encourage you to pause this video and think about that on your own. Well, when you're in the fourth period, the outermost electrons are going to be in your fourth shell. Here, you're filling out 4S1, 4S2. Then you start back filling into the 3D subshell and then you start filling again in 4P1 and so forth. You start filling out the P subshell. So as you go from potassium to krypton, you're filling out that outermost fourth shell. Now what's going on there? Well, when you're at potassium, you have 19- 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, 19. You have 19 protons and you have 19 electrons. Well I'll just draw those. 1, 2, 3, 4, 5, 6, 7, 8, 9, 10, 11, 12, 13, 14, 15, 16, 17, 18, but you only have 1 electron in that outermost, in that fourth shell, so let's just say that's that electron at a moment, just for visually. It doesn't necessarily have to be there but just to visualize that. So that 1 electron right over there, you have 19, yeah, you have 19 protons. So, you have some, I guess you could say Coulom force that is attracting it, that is keeping it there. But if you go to krypton, all of a sudden you have much more positive charge in the nucleus. So you have 1, 2, 3, 4, 5, 6, 7, 8- I don't have to do them all. You have 36. You have a positive charge of 36. Let me write that, you have plus 36. Here you have plus 19. And you have 36 electrons, you have 36 electrons- I don't know, I've lost track of it, but in your outermost shell, in your fourth, you're going to have the 2S and then you're going to have the 6P. So you have 8 in your outermost shell. So that'd be 1, 2, 3, 4, 5, 6, 7, 8. So one way to think about it, if you have more positive charge in the center, and you have more negative charge on that outer shell, so that's going to bring that outer shell inward. It's going to have more I guess you could imagine one way, more Coulomb attraction right over there. And because of that, that outer most shell is going to drawn in. Krypton is going to be smaller, is going to have a smaller atomic radius than potassium. So the trend, as you go to the right is that you are getting, and the general trend I would say, is that you are getting smaller as you go to the right in a period. That's the reason why the smallest atom of all, the element with the smallest atom is not hydrogen, it's helium. Helium is actually smaller than hydrogen, depending on how you, depending on what technique you use to measure it. That's because, if we take the simplest case, hydrogen, you have 1 proton in the nucleus and then you have 1 electron in that 1S shell, and in helium you have 2, 2 protons in the nucleus and I'm not drawing the neutrons and obviously there's different isotopes, different numbers of neutrons, but you have 2 electrons now in your outer most shell. So you have more, I guess you could say, you could have more Coulomb attraction here. This is plus 2 and then these 2 combined are negative 2. They're going to be drawn inward. So, that's the trend as we go to the right, as we go from the left to the right of the periodic table, we're getting smaller. Now what do you think is going to happen as we go down the period table? As we go down the periodic table in a given group? Well, as we go down a group, each new element down the group, we're adding, we're in a new period. We're adding a new shell. So you're adding more and more and more shells. Here you have just the first shell, now the second shell and each shell is getting further and further and further away. So as you go down the periodic table, you are getting, you are getting larger. You're having a larger atomic radius depending on how you are measuring it. So what's the general trend? Well if you get larger as you go down, that means you're getting smaller as you go up. You get smaller, smaller as you go up. So, what are going to be, what's going to be the smallest ones? Well, we've already said helium is the smallest. So what are going to be some of the largest? What are going to be some of the largest atoms? Well that's going to be the atoms down here at the bottom left. So, these are going to be large, these are going to be small. So, large over here, small over here and the general trend, as you go from the bottom left to the top right you are getting, you are getting smaller.