- 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
Atomic and ionic radii are found by measuring the distances between atoms and ions in chemical compounds. On the periodic table, atomic radius generally decreases as you move from left to right across a period (due to increasing nuclear charge) and increases as you move down a group (due to the increasing number of electron shells). Similar trends are observed for ionic radius, although cations and anions need to be considered separately. Created by Jay.
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- I dont understand why anion of an element is bigger? As we know by a trend in modern periodic table that across a period the number of valence electrons increase by one but still the atomic size decreases,so why does this not apply for the example of anion being bigger as there also only one extra valence electron is getting added?(15 votes)
- The trend you mentioned is so because as there is an additional proton with an increase of an electron. Since there is no electron shielding for the same period element, the additional proton pulls all the electron closer. Thus the radius is shorter as you go right the periodic table. However, it is not the same for ions. An anion means the number of proton stays the same while an additional electron comes in the orbital. The positive charge does not increase, so the radius will be larger due to the stronger electron repulsion. And vice, versa, a cation will have significantly smaller radius because an electron goes away while the positive charge stays the same.(40 votes)
- OK so I understand how a lithium atom would be smaller if an electron is removed b/c that's eliminating a whole energy level, but why would a beryllium atom be smaller if it became a cation and an electron was removed, since... a remaining electron would still be present in the outer shell? Do I have a knowledge gap here?(5 votes)
- When an electron is lost, the other electrons feel a stronger attraction to the nucleus. Does that make sense?(10 votes)
- How do you even measure that small a distance? ( Atomic radius )(3 votes)
- Probably the most common way to determine these distances is using a method called X-ray crystallography. Along with being able to measure distances between atoms it allows us to determine the structure of a molecule. In effect being able take a picture of the molecule.
This method involves first creating a crystal composed of the molecule with sufficient size, high purity, and a rectangular prism shape. Which in my experience is the most time consuming. Afterwards we place the crystal in an instrument where X-rays are directed toward it and the molecules in the crystals scatter those X-rays. Those scatterings essentially give us pictures of the molecule which we can combine to yield the structure of the molecule and allows us to determine the distances between the atoms.
Hope that helps.(11 votes)
- what does pm mean?(3 votes)
- Post meridiem, or after midday. Na just kidding.
In a chemistry context pm stands for picometers. It’s a unit of length equal to 10^(-12) m, or a trillionth of a meter. We have length this small to measure the sizes of atoms.
Hope that helps.(7 votes)
- I do not quite understand why the added electron to a neutral chlorine ino would increase the size of the ion. Since the electron configuration for the chloride anion is the same as the noble gas Ar, and the atomic radius generally dicreases from left of the periodic table to the right of the periodic table, so why won't the added electron makes the chiride anion smaller in terms of size?(2 votes)
- If you have a neutral chlorine atom and all you do is add an electron, then you’ve added to the repulsive force felt by the electrons. You’ve decreased the effective nuclear charge felt by the electrons towards the nucleus and so they feel less attractive force towards the nucleus and the valence electrons orbit farther from the nucleus resulting in a larger atomic radius.
This is different from the trend of decreasing atomic radii as you move left to right along a period. As you move left to right, you’re changing the type of element the atom is which means you’re adding an extra proton each step to the right. While you are also adding an extra electron, the extra proton results in a net increase in the effective nuclear charge because the attractive pull of a proton is greater than the shielding of an extra electron in the same shell. If the effective nuclear charge for elements increases as you move to the right, the electrons feel a greater force of attraction for the nucleus and the valence electrons orbit closer resulting in a smaller atomic radius.
The chlorine ion example is keeping the same number of protons but adding an electron. So, while a chloride ion has the same electron configuration as a neutral argon atom, they have different radii because of the different number of protons in the nuclei.
Hope that helps.(6 votes)
- Will this help in grinding for IChO in Canada?(3 votes)
- KA would most likely help. At the very least it would serve as an introduction to many of the topics on the IChO. It would be best to seek out additional chemistry teaching materials too. Studying from a variety of sources and becoming proficient in them would certainly help your odds.
I have to warn you though that the IChO exam is formidable for secondary school students. The difficulty of the material they test on is what you would expect at a college or university chemistry course. Here on KA they can give you a good start in general chemistry with their AP chemistry material, but the IChO also tests organic chemistry (also found on KA), physical chemistry, and spectroscopic techniques like proton NMR, C-13 NMR, and IR.
Hope that helps.(3 votes)
- Why exactly does electron repulsion increase in anions once there is an extra electron that is added? I did some research and many sources say that this is because once an electron is added, there is an increase in shielding. But when you add an electron, you're adding a valence electron, not a core electron. So why exactly does shielding increase?(2 votes)
- Shielding happens not only from core electrons, but also from electrons in the same shell. So valence electrons shield other valence electrons from the nucleus. This should make sense since all electrons have negative charges, so an electron in feels a repulsive force from a nearby electron in the same shell.
Shielding for effective nuclear charge is often first introduced as just originating from core electron electrons. But in reality shielding is more complex. Using a simpler definition of shielding works fine as an introduction, but eventually becomes problematic for more complex problems. We use a set of guidelines called Slater’s rules to give a more accurate shielding value which takes into account the shielding from electrons in the same shell.
So if you have a neutral atom, and all you do is add an extra valence electron, you’re adding an extra repulsive force to the existing valence electrons without any increase in the attractive force from the nucleus. The net change is that the valence electrons have less effective nuclear charge be applied to them and so feel less attracted to the nucleus and orbit at a greater distance from the nucleus resulting in a larger atomic radius as compared to the neutral atom.
Hope that helps.(5 votes)
- When going left to right in the d or f block, isn't that adding electrons to a non-valence shell? So why does the shielding effect stay the same?(2 votes)
- Well first elements in the d and f block do have d and f electrons as a part of their valence electrons.
In a simplified version here, they only consider the shielding due to core electrons in lower shells. However, in reality shielding is caused by valence electrons in the same shell as well. Which should make sense since all electrons are negatively charged and they all repel each other. This also means the shield would change as we add electrons.
Hope that helps.(3 votes)
- If two ions have the same number of electrons but one has more protons which will be larger One with more protons or less protons?(1 vote)
- The other replier is correct, I would just like to expand on it.
The ion with more protons will have a smaller atomic radius because the effective nuclear charge felt by the electrons is greater. The effective nuclear charge being the force of attraction felt between the protons of the nucleus and the electrons in orbit around them. If the protons and electrons have a stronger force of attraction between them, they will be able to move closer together resulting in a smaller electron cloud.
Hope that helps.(4 votes)
- 'On the periodic table, atomic radius generally decreases as you move from left to right across a period (due to increasing nuclear charge) and increases as you move down a group (due to the increasing number of electron shells).'
but the nuumber of protons also increases down the group, Isnt it?(2 votes)
- Yes, but the higher shell numbers cause a greater increase in the atomic radii than is countered by the contraction of the additional protons. The net change is an increase in atomic radius.
Hope that helps.(2 votes)
In this video, we're going to look at atomic and ionic radii. And first, we'll start with the atomic radius. So if you think about an atom as a sphere, the idea of atomic radius is simple. You would just take this as a sphere here, and then a sphere of course would have fixed and defined radius. And so that would be one way of thinking about it. The problem is that an atom doesn't really have a fixed, defined radius like this sphere example, because there's a nucleus and then there's this electron cloud, or this probability of finding your electron. So there's no real, clear defined boundary there, and so it's difficult to have a fixed and defined radius. So what chemists do is they take two identical atoms. So let's say these are two atoms bonded together, the same element. And if you find their nuclei-- so let's say that that's their nuclei here-- and you measure the distance between those two nuclei, so this would be our distance d between our two nuclei. If you take half of that distance, that would be a good approximation of the atomic radius of one of those atoms. And so that's the idea behind the definition of atomic radius. Let's look at the trends for atomic radius, and first we'll start with group trends. And so here we have two elements found in group one, so hydrogen and lithium. And let's go ahead and sketch out the atoms first. And so we start with hydrogen, which has atomic number of 1, which means that it has one proton in the nucleus. So here's our nucleus for hydrogen, so one proton. In a neutral atom, the number of protons equals the number of electrons, and so therefore there must be one electron. So go ahead and sketch in our electron here. And we'll make things really simple and just show this simple version of the atom, even though we know it doesn't really exactly look like this. And when we do lithium, atomic number of 3, so that means three protons in the nucleus of lithium. So this is representative of lithium's nucleus with three protons and three electrons. Two of those electrons are in the inner shell. So let me go ahead and show two of lithium's electrons in the inner shell, so that would be in the first energy level. And then we would need to account for one more, so lithium's third electron is in the second energy level or at the outer shell in this example. And so here we have our two atoms. And you can see as you go down a group, you're going to get an increase in the atomic radius. And that's because as you go down a group, you're adding electrons in higher energy levels that are farther away. So in this case, we added this electron to a higher energy level which is farther away from the nucleus, which means that the atoms of course would get larger and larger. So you're adding more stuff to it, so it's kind of a simple idea. Let's look at period trends next. As you're going across a period this way, so as you're going this way, you're actually going to get a decrease in the atomic radius. And let's see if we can figure out why by once again drawing some simple pictures of our atoms. And so lithium with atomic number of 3, so we've already talked about that. So there are three protons in the nucleus of lithium. So I'm going to go ahead and write that in here. So 3 positive charge for the nucleus of lithium. And we have to account for the three electrons. So once again two of those electrons were in an inner shell, so there we go, and then we had one electron in an outer shell, so the picture is something like this. Now, let's think about what's going to happen to that outer electron as a result of where it is. So this outer electron, this one right here in magenta, would be pulled closer to the nucleus. The nucleus is positively charged, that electron is negatively charged, and so the positively charged nucleus is going to pull that electron in closer to it. At the same time, those negatively charged inner shell electrons are going to repel it. So let me go ahead and highlight these guys right here. These are our inner shell electrons. Like charges repel. And so you could think about this electron right here wanting to push this outer electron that way, and this electron wanting to push this electron that way. And so the nucleus attracts a negative charge, and the inner shell electrons repel the outer electron. And then we call this shielding, because the inner shell electrons are shielding that magenta electron from the pole of the nucleus. So this is called electronic shielding or electron screening. Now, it's going to be important concepts. So now let's go ahead and draw the atom for beryllium, so atomic number 4. And so here's our nucleus for beryllium. With an atomic number of 4, that means there are four protons in the nucleus, so a charge of four plus in our nucleus. And we have four electrons to worry about this time, so I'll go ahead and put in the two electrons in my inner orbital in our first energy level. And then we have two electrons in our outer orbital, or our second energy level. And so again, this is just a rough approximation for an idea of what beryllium might look like. And so when we think about what's happening, we're moving from a charge of 3 plus with lithium to a charge of 4 plus with beryllium. And the more positive your charges, the more it's going to attract those outer electrons. And when you think about the idea of electron screening, so once again we have these electrons in green here shielding our outer shell electrons from the effect of that positively charged nucleus. Now, you might think that outer shell electrons could shield, too. So you might think that oh, this electron right here in magenta could shield the other electron in magenta. But the problem is they're both at pretty much the same distance from the nucleus, so outer shell electrons don't really shield each other. It's more of these inner shell electrons. And because you have the same number of inner shell electrons shielding as in the lithium example-- so let me go ahead and highlight those again. So we have two inner shell electrons shielding a beryllium. We also have two inner shell electrons shielding in lithium. Because you have the same number of shielding but you have a higher positive charge, those outer electrons are going to feel more of a pull from the nucleus. And they're going to be pulled in even tighter than you might imagine, or at least tighter than our previous example. So these electrons are pulled in even more. And because of that, you're going to get the beryllium atom as being smaller than the lithium atom, hence the trend. Hence as you go across the period, you're always going to increase in the number of protons and that increased whole is going to pull those outer electrons in closer, therefore decreasing the size of the atom. All right. Let's look at ionic radius now. And ionic radius can be kind of complicated depending on what chemistry you are involved in. So this is going to be just a real simple version. If I took a neutral lithium atom again, so lithium-- so we've drawn this several times. Let me go ahead and draw it once more. So we have our lithium nucleus, which we have three electrons. So once again I'll go ahead and sketch in our three electrons real fast. Two electrons in the inner shell, and one electron in the outer shell like that. And let's say you were going to form a cation, so we are going to take away an electron from our neutral atom. So we have-- let me go ahead and draw this in here-- we had a three protons in the nucleus and three electrons those cancel each other out to be a neutral atom. And if we were to take away one of those electrons, so let's go ahead and show lithium losing an electron. So if lithium loses an electron, it's going to lose that outer electron. So the nucleus still has a plus 3 charge, because it has three protons in it. And we still have our two inner shell electrons like that, but we took away that outer shell electron. So we took away this electron in magenta, so let me go ahead and label this. So we lost an electron, so that's this electron right here, and so you could just show it over here like that. And by doing so, now we have three positive charges in our nucleus and only two electrons. And so therefore our lithium gets a plus 1 charge. So it's Li plus, it's a cation. And so we formed a cation, which is smaller than the neutral atom itself. And that just makes intuitive sense. If you take away this outer electron, now you have three positive charges in the nucleus and only two electrons here. So it's pulling those electrons in, you lost that outer electron, it's getting smaller. And so the cation is smaller than the neutralize atom. And so we've seen that neutral atoms will shrink when you convert them to cations, so it kind of makes sense that if you take a neutral atom and add an electron, it's going to get larger. And so that's our next concept here. So if we took something like chlorine, so a neutral chlorine atom, and we added an electron to chlorine, that would give it a negative charge. So we would get chlorine with a negative charge, or the chloride anion, I should say. And so in terms of sizes, let's go ahead and draw a representative atom here. So if this is our neutral chlorine atom and we add an electron to it, it actually gets a lot bigger. So the anion is bigger than the neutral atom. And let's see if we can think about why here. So if we were to draw an electron configuration, or to write a noble gas electron configuration for the neutral chlorine-- so you should already know how to do this-- you would just write your noble gas in brackets. So neon and then 3s2, 3p5, so seven electrons in the outer shell for the neutral chlorine atom. For the chloride anion, you would start off the same way. You would say neon in brackets, 3s2. And you'd be adding an electron to it. So it wouldn't be 3p5, it would be 3p6 like that. And so now we would have so this would give us eight electrons in our outer shell, and this would give us only seven electrons in our outer shell. Now, the explanation for the larger size of the chloride anion in most textbooks is, you'll see people say that the addition of this extra electron here, so that means that those electrons are going to repel each other more. You have eight of them instead of seven, and so because they repel each other more, it gets a little bit bigger. And that makes sense, but you'll see some people disagree with that explanation, and I haven't really seen a great alternative offered. And so however you want to think about it, generally the anion is larger than the neutral atom. But in terms of the explanation for that, you could think about it as electrons are repelling each other if you wanted to, despite the fact that people disagree with that. You could think about just more stuff as a really simple way of thinking about it. But again, in general for exams, think about the anion being larger.