Using only the periodic table, predict the most stable ion for Na, Mg, Al, S, Cl, K , Ca, and Ga

Using only the periodic table, predict the most stable ion...

Key Concepts

Ion Stability Prediction

This concept involves using an element's position on the periodic table to predict which ion it will most likely form based on achieving a noble gas electron configuration. Typically, elements will lose or gain electrons to obtain a full valence shell. For example, metals in groups 1 and 2 tend to lose electrons to form cations, while nonmetals in groups 16 and 17 tend to gain electrons to form anions. This approach relies on the octet rule and the general periodic trends in electron configuration.

Periodic Trends in Ionic Radii

The periodic table provides guidance on how atomic and ionic sizes change across periods and down groups. Generally, moving across a period leads to a decrease in atomic radius due to increasing effective nuclear charge, whereas moving down a group increases the radius due to the addition of electron shells. When ions are formed, the loss or gain of electrons will also modify ionic sizes, making it possible to predict relative radii based solely on an element’s position.

Effective Nuclear Charge

Effective nuclear charge is the net positive charge experienced by an electron in a multi-electron atom. It plays a crucial role in determining both atomic and ionic sizes. In a series of ions that are isoelectronic, a higher effective nuclear charge will pull the electrons closer to the nucleus and result in a smaller ionic radius. Conversely, a lower effective nuclear charge means less pull on the electrons and a larger ionic radius.

Isoelectronic Series and Comparative Ionic Sizes

Many ions formed by different elements can end up having the same number of electrons, leading to an isoelectronic series. Within such a series, the ionic radius is primarily determined by the number of protons in the nucleus; more protons mean a greater effective nuclear charge and, therefore, a smaller radius. This concept allows us to rationalize and arrange ions from largest to smallest based on their corresponding nuclear charges.

Comparison with Empirical Data

While the periodic table offers strong predictive power concerning ion formation and trends in ionic radii, empirical data such as that presented in figures or measured radii can exhibit deviations due to factors like coordination effects and experimental conditions. Comparing predictions based solely on periodic trends with empirical results helps validate the model and highlights any nuances that must be considered in real-world applications.

Transcript

00:01 Problem 3 says to use the periodic table to predict the most stable ion for each of the atoms shown here, and then to arrange these from the largest to smallest ionic radius, and then compare your answers to figure 8 .8.

00:15 So for most main group elements, there's a single ion, and these atoms either lose or gain electrons to take on a noble gas configuration.

00:26 Basically, the noble gas that they're closest to in the periodic table.

00:31 So here i've provided the three noble gases that are of interest to us.

00:38 And these will be the noble gases that each of these atoms will either lose or gain electrons to become more alike.

00:49 So whether atoms gain or lose electrons sort of depends on electronegativity.

00:57 So starting with these over here, sulfur and chlorine, their closest noble gas is argon here.

01:08 And so because chlorine is in the next group over, you can see that the easiest way for chlorine to become like argon is to gain a single electron.

01:26 So gaining one electron makes it argon -like.

01:31 Similarly, sulfur here can become more argon -like by gas.

01:36 Gaining two electrons.

01:38 So those are the ions for sulfur and chlorine.

01:44 Now for these groups over here, it's a little bit trickier because they're actually closer to the previous periods, noble gas, and the one over here.

01:56 If you count across the main group elements, aluminum, for instance, is 1, 2, 3, 4, 5 electrons away from argonne in this direction, but one, two, three electrons away from neon in this direction.

02:11 So for aluminum and by extension gallium, the easiest way to become noble gas -like is to lose electrons, specifically three electrons.

02:24 So by losing three electrons, aluminum becomes more like neon, and by losing three electrons, gallium becomes more like argon.

02:33 For these over here, though, sodium, magnesium, potassium, and calcium, it's very clear that the easiest way for these to become noble gas -like is just like aluminum and galleon to lose electrons.

02:46 So in this first group over here, sodium of potassium simply need to lose a single electron to become either like neon or argon, respectively.

02:58 And then the same for magnesium and calcium.

03:01 They have to lose two electrons to become more like argon or neon.

03:06 So those are the most common ion, the most stable ionic forms for each of these atoms.

03:14 And now we have to use what we know about radius patterns across the periodic table to determine what the radii for these ions looks like.

03:26 So as a general rule, atomic radius increases as you go down the periodic table and in general, sort of as you go across, but from left to right.

03:42 But that's not necessarily true of ions, because what you have to think about more is whether the nucleus, so the positive charge in the center of each of these atoms, is becoming associated with either more or fewer electrons.

04:00 So as the nucleus becomes larger across a period, losing electrons will make the atomic radius smaller...

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