Periodic Trend: Ionic Radius - Online Tutor, Practice Problems & Exam Prep

Understanding the concept of ionic radius is crucial for grasping periodic trends in chemistry. The ionic radius refers to the distance from an ion's nucleus to its outermost electron shell. Unlike other periodic trends, it does not follow the general top-right direction on the periodic table. Instead, the ionic radius is influenced by the number of electrons in an ion. Cations, which are positively charged ions, are typically smaller than their neutral atoms since they lose valence electrons, reducing the number of electron shells. Conversely, anions, which are negatively charged ions, are larger than their neutral atoms because they gain electrons, increasing electron-electron repulsion in the outer shell. Therefore, the more electrons an ion has, the larger its ionic radius becomes. This concept is essential in predicting and explaining the properties of ions in various chemical contexts.

Periodic Trend: Ionic Radius

For a majority of the periodic trends, looking at the periodic table can tell you if they're increasing or decreasing. In all our periodic trends, we're heading towards the top right corner and heading in that directional either increase or decrease that particular periodic trend. Now ionic radius doesn't fit into this idea. Ionic radius does not involve looking at the periodic table or Let's say ionic radius equals the distance between an ion's nucleus and its outer shell, and the periodic trend is this. The ionic radius increases as the number of electrons of that ion increases.

So we're going to say here when we're looking at ionic radius, we're looking at ions and ions can either be positive or negative. For positive ions, which are cations, we're going to say cations tend to be smaller than their neutral parent form. So what do I mean by this? Well, if we take a look here at lithium, the lithium atom is I1S2S1. It has one electron in its valence electron shell, right? Because its second shell only has one. When I become lithium ion I lose 1 electron, I lose it from that valence shell. So now the lithium ion no longer has 2 shells, it only has one shell that contains electrons and we can see that there is a drop in the size for that particular ion. So just remember the neutral form of the element is bigger than the cation form. So the general trend is losing an electron or electrons causes a decrease in your ionic radius.

Now if we're looking at the anions, anions tend to be larger than their neutral parent forms. Oxygen is S1S2S2P4 here. These electrons are in the second shell, and there's six of them total. When we become the oxide ion, we gain two more electrons. So now my outer shell has eight electrons. It may not be as apparent, but the anion is slightly bigger than the neutral form because Y we said earlier the ionic radius increases as the number of electrons increases. So basically more electrons equal larger ionic radius.

So we're going to say gaining an electron or electrons causes an increase in your ionic radius. So just remember we don't look at the periodic table to determine the trend in ionic radius. Instead, we look at the total number of electrons that ion has. The more electrons it has, the bigger its ionic radius.

Periodic Trend: Ionic Radius Example 1

So here for this example question, it says which of the following statements is false. All right, so the cation has a smaller ionic radius than the calcium atom. Remember atom means it's neutral. Calcium is in Group two way, so its charge is 2. Plus remember we said earlier that cations lose electrons which causes a decrease for the ionic radius of the positive ion. So yeah, since calcium is a positive ion, it will in fact have a smaller ionic radius than the neutral form. So this is true.

Sulfide ion has a larger ionic radius than these sulfur atom. Sulfide I'd means it's an anion, It's gained electrons. Remember we said that anions are larger than their neutral form because more electrons equals larger ionic radius. So this is true. The cadmium ion has a larger ionic radius than the cadmium atom. So cadmium has a charge of two. Plus remember cations are smaller than their neutral form because they have less electrons. So this is false and our answer?

Bromide ion. I mean it's a negative charge, it's an anion. Anions are larger than their neutral forms because again, more electrons equals larger ionic radius O. The only statement here that's false would have to be option C.

Periodic Trend: Ionic Radius Example 2

Here we're going to say when ranking ionic radius you must determine the total number of electrons for atoms and ions. Now more electrons equals larger ionic radius. But recall isoelectronic species are atoms or ions that have the same number of electrons. In those cases what do we do? Well if we take a look at this example it says arrange the atoms and or ions in order of decreasing ionic radius. So that means that we need to go from largest ionic radius to smallest ionic radius.

So here if we take a look in step one, it says determine the total number of electrons for each element or ion. Just remember the higher the number of electrons then the greater the ionic radius. So if we take a look here, iron, when it is neutral, it has 26 electrons because its atomic number is 26. Here it's lost two, so it's going to have 24 electrons remaining. Manganese has an atomic number of 25. So when it's neutral it has 25 electrons plus one means it's lost 1 S It has 24 electrons as well.

Nickel has an atomic number of 28. Here it's just lost two electrons, so it has 26 remaining. And then finally here we have zinc. Zinc has an atomic number of 30, So when it's neutral it has 30 electrons +2 means now it's lost to so it only has 28 electrons remaining. All right, so we can arrange this somewhat from largest to smallest. We know that zinc 2 plus ion will be the largest because it has the most electrons followed by nickel.

But how do we break the tide between iron 2 plus and manganese plus one? That's where Step 2 comes into play. If the atoms or ions have equal electrons, meaning they're isoelectronic, then the more negative the charge, the larger the ionic radius. O we're going to say for an ionic for isoelectric species we just said the greater the negative charge than the greater the ionic radius. So a -3 charge would be bigger than a -2 charge, bigger than a -1, bigger than 0, bigger than plus 1 + 2 + 3.

Now of course you could have some species that have a charge that's outside of -3 or beyond, plus three as well. Same rule would apply. So both iron 2 plus and manganese plus one have 24 electrons, but manganese 1 is bigger than +2. So we'd say that manganese plus one ion is bigger than iron 2 plus ion. So just remember when it comes to ions, more electrons equals larger ionic radius. If they're tied, then we look at the charge to break that tie. The more negative the charge, the larger dianic radius will be.

Arrange the following atoms and/or ions in the order of increasing size:Br ^–, Kr, Rb⁺, Sr²⁺.

• A. Kr < Br^- < Sr²⁺ < Rb⁺
• B. Kr < Sr²⁺ < Rb⁺ < Br^-
• C. Sr²⁺ < Rb⁺ < Kr < Br^-
• D. Rb⁺ < Sr²⁺ < Kr < Br^-
• E. Br^- < Kr < Rb⁺ < Sr²⁺

Arrange the following isoelectronic series in order of decreasing radius: F^–, O^2– , Mg²⁺, Na⁺.

• A. F^– > O^2– > Mg²⁺ > Na⁺
• B. F^– > O^2– > Na⁺ > Mg²⁺
• C. O^2– > F^– > Na⁺ > Mg²⁺
• D. O^2– > F^– > Mg²⁺ > Na⁺

For an isoelectronic series of ions, the ion that is the smallest is always 

• A. The ion with the fewest protons.
• B. The least positively charged ion.
• C. The ion with the highest atomic number.
• D. The ion with the most neutrons.
• E. The ion with the most electrons.

Here’s what students ask on this topic:

What is the ionic radius?

The ionic radius refers to the measure of an atom's ion in a crystal lattice. It's essentially the distance from the center of the nucleus to the outer edge of the electron cloud of the ion. When an atom loses an electron to form a cation, its ionic radius decreases because the loss of an electron reduces electron-electron repulsion and allows the remaining electrons to be pulled closer to the nucleus. Conversely, when an atom gains an electron to become an anion, its ionic radius increases due to the increase in electron-electron repulsion, which causes the electrons to spread out more.

Ionic radii are important because they affect the physical and chemical properties of ionic compounds, including their lattice energies, melting points, and solubilities. The size of the ions influences how closely they can pack in a crystal lattice and how they interact with each other, which in turn affects the compound's overall stability and reactivity.

Why does the ionic radius decrease across a period?

The ionic radius decreases across a period on the periodic table primarily due to the increase in the nuclear charge without a corresponding increase in the number of electron shells. As you move from left to right across a period, protons are added to the nucleus, and electrons are added to the same principal energy level. However, the inner shells remain the same, and thus there is no additional shielding effect to counterbalance the increased nuclear charge.

This increased nuclear charge pulls the electrons closer to the nucleus, resulting in a smaller atomic radius. In ions, this effect is also present; cations (positively charged ions) have lost electrons, further reducing electron-electron repulsion within the atom and making them smaller. Anions (negatively charged ions) gain electrons, but the increase in negative charge is usually not enough to overcome the pull of the increased nuclear charge across the period, so they also tend to be smaller than their preceding atoms, although they are larger than their cation counterparts.

Why does the ionic radius increase down a group?

The ionic radius increases down a group in the periodic table primarily due to the addition of electron shells. As you move from one period to the next in a group, each element has an extra shell of electrons compared to the one above it. This increase in the number of electron shells leads to a larger distance between the nucleus and the outermost electrons.

Additionally, with each added shell, there is a decrease in the effective nuclear charge experienced by the electrons in the outermost shell because the inner shells partially shield the outer electrons from the pull of the nucleus. This shielding effect allows the outermost electrons to be less tightly bound to the nucleus and therefore the ion's radius to expand.

Moreover, as the ionic radius increases, the electrons are held less tightly due to the increased electron-electron repulsion within the larger cloud of electrons. This further contributes to the increase in ionic size as you descend a group in the periodic table.