5.8. Trends within the d-Block

Today, we will look at the trends in atomic and ionic radii for transition metals. Can anyone explain why we see a decrease in atomic radius moving from Scandium to Zinc?

Is it because of the increasing number of protons that pull the electrons closer?

Exactly! As the effective nuclear charge increases, it draws the electrons in closer. This is why the atomic radius decreases. Can someone explain how the ionic radius behaves when we form cations and anions?

Cations are smaller because they lose electrons, which reduces electron-electron repulsion?

Correct! And anions are larger because they gain electrons, increasing electron-electron repulsion. Let's remember this with the acronym 'CAL,' which stands for 'Cations Are Little' and 'Anions Are Larger.'

That's a good way to remember it!

Let's summarize what we learned: moving across the d-block, atomic radius decreases due to increasing Z_eff. Down a group, the atomic radius increases due to additional energy levels. And for ionic radii, cations are smaller than their neutral atoms, while anions are larger.

Now, let's discuss ionization energy. Who can tell me how ionization energy changes across the d-block?

I think it increases as you go to the right because the atoms are pulling the electrons in more due to the increased charge?

Correct! That’s the general trend. It's worth noting, though, that there are dips between certain elements, especially between the d⁵ to d⁶ and d¹⁰ to d¹¹ configurations because half-filled and fully filled subshells are more stable. Can anyone guess why we see this?

Maybe because it's easier to remove an electron from an unstable arrangement?

Exactly! This stability makes it harder to remove those electrons. Keep this in mind: 'Falling’ meaning lower energy when heading into a more stable configuration helps remember the lows and highs of ionization energy!

Got it! So the acronym 'FRESH' helps us—'Falling to Rediscover Energy Stability High.'

Great job! So we learned that ionization energy increases across the d-block but decreases down a group due to increased electron shielding. Additionally, bear in mind those key anomalies as we discuss properties in chemistry.

Let’s wrap up today with oxidation states. What do you notice about the oxidation states of early transition metals compared to late transition metals?

Early transition metals can have higher oxidation states like +6, while late transition metals usually have +2.

Exactly! Early transition metals, such as Titanium, Vanadium, and Chromium exhibit higher oxidation states often due to the availability of d electrons. This is really about how these states maximize ligand field stabilization energy. Can anyone summarize what happens with Copper and Zinc?

Copper prefers +1 and +2 oxidation states, while Zinc is stable only in +2 state, right?

That’s right! For Copper, +1 state is stable because it avoids partially filled configurations, while Zinc remains in the +2 state due to fully filled d-subshells. Let’s remember with the mnemonic 'Cousins for Stability,' implying Copper's unpredictability and Zinc's consistency in their oxidation states.

That makes sense and is easy to remember!

To conclude, we’ve reviewed how oxidation states vary significantly across the d-block and discussed their stability in light of electron configurations. Remember, higher oxidation states are more common for early transition metals, while late transition metals prefer lower states!