Video: Applying Knowledge of Ionic Bonding and Ionization

For statements I and II, state for each if they are true or false. I) Larger ions form weaker ionic bonds than smaller ions with the same charge do. II) Nonmetal atoms lose electrons less readily than metal atoms do. If both are true, state if II is a correct explanation for I.

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Video Transcript

For statements I and II, state for each if they are true or false. I) Larger ions form weaker ionic bonds than smaller ions with the same charge do. II) Nonmetal atoms lose electrons less readily than metal atoms do. If both are true, state if II is a correct explanation for I.

Statement I refers to ions in general. So we have to consider both positive and negative ions. We can get a rough idea of the size of an ion by counting the number of electron shells. So, for instance, we can be confident that an Na⁺ ion is larger than an Li⁺ ion because a sodium plus ion has electrons in the second shell, where lithium ion doesn’t. We can see that trend continue down the alkali metals. For instance, potassium plus has electrons in the third shell, where sodium plus doesn’t.

Now let’s look at some negatively charged ions. We can see the same thing for the group 17 ions, the halides. So fluoride chloride and bromide get progressively larger. Statement I says that as these ions get larger, we expect the ionic bonds they form to get weaker. Let’s take, for example, lithium fluoride and potassium fluoride.

In lithium fluoride, the two ions are quite close together, whereas in potassium fluoride, they’re quite far apart. This means that the force of attraction between potassium ions and fluoride ions is weaker than the force of attraction between lithium ions and fluoride ions. We see exactly the same thing when we swap out the negative ion with a bigger negative ion of the same charge, like swapping fluoride for bromide. So on this basis, we expect statement A [I] to be true. Larger ions with the same charge will have weaker ionic bonds.

We can check this by looking up lattice enthalpies. A lattice enthalpy is the enthalpy of formation of an ionic crystal from the separate ions in their gas phase. The value for lithium fluoride is relatively high, at over 1000 kilojoules per mole. If we swap out the lithium ion for the potassium ion, we can see a reduction in the lattice enthalpy to a little over 800 kilojoules per mole. And if we go back and swap out just the fluoride with bromide, we get about the same value, a little over 800 kilojoules per mole for lithium bromide. These two lower values show that swapping out a small ion for a large ion of the same charge will lead to weaker ionic bonds. So statement I is definitely true.

So, on to statement II. Statement II says that nonmetal atoms lose electrons less readily than metal atoms do. Roughly speaking, we can separate the periodic table into nonmetals and metals across this dividing line. When we say an atom loses electrons less readily, we mean that it takes more energy for that atom to be ionized. For an atom in its ground state, the first ionization is the removal of its outer electron.

Now comparing metals and nonmetals as a whole is going to be quite tricky. So we’re going to narrow it down to metals and nonmetals in the same period of the periodic table. Let’s, for the sake of argument, take period three. On the left-hand side, we have the three metals in period three: sodium, magnesium, and aluminum. And playing for the nonmetals, we have silicon, phosphorus, sulfur, chlorine, and argon.

The next thing we’re going to consider is the configuration of the electrons for each of these elements. We’re removing an electron from the outer shell first, so that’s what we’ll focus on. An atom of sodium has an electron in the 3s subshell, while magnesium has two. An atom of aluminum has a full 3s subshell and one electron in the 3p subshell. An atom of silicon has two electrons in the 3p subshell. So we can see across the period we’re only adding electrons to the same overall shell, the third electron shell.

But as we’re moving left to right across the periodic table, we’re also increasing the nuclear charge. An atom of sodium has 11 protons in its nucleus. An atom of magnesium has 12, and so on across the group until we have 18 protons per nucleus for argon. So what we have in moving left to right, from metals to nonmetals, across the periodic table is similar shielding but a more positive nucleus.

Therefore, in general, the ionization energy of nonmetals is greater than the ionization energy of metals. So the statement “Nonmetal atoms lose electrons less readily than metal atoms do” is true.

Now that we’ve dealt with statements I and II separately, let’s see about the last part. If both are true, we need to decide if statement II is a correct explanation for statement I. Broadly speaking, statement II is about how easy it is to make a positive ion from a given type of atom, while statement I is about the force of attraction between ions once they’ve been made. So statement II doesn’t really deal with the same material as statement I. So it can’t be considered a correct explanation. So for this bit, we simply write “False.” Statement II is not a correct explanation for statement I.

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