# Relationship between the size of an ion to its atom

### Mini-video on ion size (video) | Khan Academy if the atoms lose electrons, the ions are relatively smaller than the regular one and they become positively charged. What is the relation of lattice energy to size and charge? will mean a smaller ion (when compared to its atom) because the Coulombic attraction between the. The ionic radius is the the distance from the nucleus to the outermost electrons in a ion. While the atomic radius is the distance from the nucleus.

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.

• What is the difference between the ionic radius and the atomic radius of an element?
• Mini-video on ion size

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.

### Explain the relationship between the relative size of an ion to its atom and the charge on the ion

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.

## 7.3: Sizes of Atoms and Ions

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. 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. In a similar approach, we can use the lengths of carbon—carbon single bonds in organic compounds, which are remarkably uniform at pm, to assign a value of 77 pm as the covalent atomic radius for carbon.

If these values do indeed reflect the actual sizes of the atoms, then we should be able to predict the lengths of covalent bonds formed between different elements by adding them. A similar approach for measuring the size of ions is discussed later in this section. Covalent atomic radii can be determined for most of the nonmetals, but how do chemists obtain atomic radii for elements that do not form covalent bonds?

For these elements, a variety of other methods have been developed. This is somewhat difficult for helium which does not form a solid at any temperature. Periodic Trends in Atomic Radii Because it is impossible to measure the sizes of both metallic and nonmetallic elements using any one method, chemists have developed a self-consistent way of calculating atomic radii using the quantum mechanical functions.

### What is the difference between the ionic radius and the atomic radius of an element? | Socratic

The sizes of the circles illustrate the relative sizes of the atoms. The calculated values are based on quantum mechanical wave functions. Web Elements is an excellent on line source for looking up atomic properties. For all elements except H, the effective nuclear charge is always less than the actual nuclear charge because of shielding effects. The greater the effective nuclear charge, the more strongly the outermost electrons are attracted to the nucleus and the smaller the atomic radius.

Atomic radii decrease from left to right across a row and increase from top to bottom down a column. The atoms in the second row of the periodic table Li through Ne illustrate the effect of electron shielding. Although electrons are being added to the 2s and 2p orbitals, electrons in the same principal shell are not very effective at shielding one another from the nuclear charge.

In contrast, the two 2s electrons in beryllium do not shield each other very well, although the filled 1s2 shell effectively neutralizes two of the four positive charges in the nucleus. Consequently, beryllium is significantly smaller than lithium.

Similarly, as we proceed across the row, the increasing nuclear charge is not effectively neutralized by the electrons being added to the 2s and 2p orbitals. The result is a steady increase in the effective nuclear charge and a steady decrease in atomic size.

The Atomic Radius of the Elements. The atomic radius of the elements increases as we go from right to left across a period and as we go down the periods in a group. The increase in atomic size going down a column is also due to electron shielding, but the situation is more complex because the principal quantum number n is not constant. In group 1, for example, the size of the atoms increases substantially going down the column. It may at first seem reasonable to attribute this effect to the successive addition of electrons to ns orbitals with increasing values of n. However, it is important to remember that the radius of an orbital depends dramatically on the nuclear charge. That force depends on the effective nuclear charge experienced by the the inner electrons. In fact, the effective nuclear charge felt by the outermost electrons in cesium is much less than expected 6 rather than This means that cesium, with a 6s1 valence electron configuration, is much larger than lithium, with a 2s1 valence electron configuration.

The effective nuclear charge changes relatively little for electrons in the outermost, or valence shell, from lithium to cesium because electrons in filled inner shells are highly effective at shielding electrons in outer shells from the nuclear charge. The same dynamic is responsible for the steady increase in size observed as we go down the other columns of the periodic table.

Irregularities can usually be explained by variations in effective nuclear charge. Note Electrons in the same principal shell are not very effective at shielding one another from the nuclear charge, whereas electrons in filled inner shells are highly effective at shielding electrons in outer shells from the nuclear charge. Identify the location of the elements in the periodic table.

Ionic Size - Which ion is bigger?