Position and electronic configuration

Where iron, copper and gold sit

Think of a coin, an iron bridge and the copper wire in a charger. All of these metals live in the middle of the periodic table, in a wide block between the s block on the left and the p block on the right. This is the d block.

In Class 11 you saw how the table is built. An element belongs to the block of the subshell that receives its last electron. In period 4, potassium and calcium fill $4s$. Then the next ten electrons go into $3d$, one element after another, from scandium to zinc. After zinc, the $4p$ subshell fills.

The same thing happens in the next periods. So the d block has four rows, called series:

• the 3d series, from scandium to zinc;
• the 4d series, from yttrium to cadmium;
• the 5d series, lanthanum and then hafnium to mercury;
• the 6d series, actinium and then rutherfordium to copernicium.

The d block covers groups 3 to 12. It is a bridge between the very reactive metals of the s block and the non-metals of the p block. That is why these elements are called transition elements. Do all ten elements of a series deserve this name? The next part answers it.

Which elements are transition elements?

A transition element has an incomplete d subshell in its atom, or in one of its common ions. Incomplete means that the d subshell holds 1 to 9 electrons, not 0 and not 10.

Test two elements.

• Scandium. The atom is $[\text{Ar}]\,3d^1 4s^2$. The d subshell has one electron, so scandium passes. Its common ion, $\text{Sc}^{3+}$, has no d electron, but the atom is enough.
• Zinc. The atom is $[\text{Ar}]\,3d^{10} 4s^2$ and the ion $\text{Zn}^{2+}$ is $3d^{10}$. The d subshell is full in both, so zinc fails.

Cadmium and mercury behave like zinc. So zinc, cadmium and mercury are d-block elements, but they are not transition elements. Copper is a transition element. Its atom is $3d^{10} 4s^1$, but the ion $\text{Cu}^{2+}$ is $3d^9$.

Why does this matter? Colour, magnetism and many oxidation states all need a partly filled d subshell. Zinc compounds are white, zinc has only one oxidation state, $+2$, and $\text{Zn}^{2+}$ is not magnetic. The next lessons show why.

Filling the 3d series

In Class 11 you learned the filling order, with $4s$ before $3d$. The picture shows the result for the ten elements.

Read it from left to right. From scandium to nickel, each element adds one electron to $3d$ and keeps $4s^2$. Hund's rule puts one electron in each of the five boxes before any pair forms. So the number of unpaired electrons climbs to five at manganese and then falls.

Two elements break the pattern.

• Chromium is $3d^5 4s^1$, not $3d^4 4s^2$.
• Copper is $3d^{10} 4s^1$, not $3d^9 4s^2$.

You met the reason in Class 11. A half-filled or a full d subshell is specially stable, and the $4s$ and $3d$ energies are so close that one electron can move at a small cost.

The general configuration of a d-block atom is $(n-1)d^{1\text{ to }10}\,ns^{0\text{ to }2}$. The heavier rows break the pattern more often, because the $4d$ and $5s$ orbitals, and the $5d$ and $6s$ orbitals, are even closer in energy.

• Niobium, molybdenum, ruthenium, rhodium and silver each have only one $5s$ electron.
• Palladium is $[\text{Kr}]\,4d^{10}$ with an empty $5s$. This is the best-known exception of the second row.
• Platinum is $[\text{Xe}]\,4f^{14}5d^9 6s^1$ and gold is $[\text{Xe}]\,4f^{14}5d^{10} 6s^1$.

Why the 4s electrons leave first

Here is a puzzle. The $4s$ orbital fills before $3d$, so it seems to lie lower. Yet when an atom of the series loses electrons, the $4s$ electrons go first. Does that contradict the filling order?

No. Once the atom has d electrons, the $3d$ electrons sit closer to the nucleus than the $4s$ electrons and are held more tightly. The $4s$ orbital is larger, and it belongs to the outermost shell, $n = 4$. A cation loses electrons from the shell with the highest $n$ first, as you learned in Class 11.

So the rule for ions of the 3d series is:

• remove the $4s$ electrons first, and then the $3d$ electrons.

Here are four examples.

• Iron is $[\text{Ar}]\,3d^6 4s^2$. $\text{Fe}^{2+}$ is $3d^6$. $\text{Fe}^{3+}$ is $3d^5$.
• Manganese is $3d^5 4s^2$. $\text{Mn}^{2+}$ is $3d^5$. $\text{Mn}^{7+}$ has no d electron at all.
• Copper is $3d^{10} 4s^1$. $\text{Cu}^+$ is $3d^{10}$. $\text{Cu}^{2+}$ is $3d^9$.
• Chromium is $3d^5 4s^1$. $\text{Cr}^{3+}$ is $3d^3$.

Here is a shortcut. Every $\text{M}^{2+}$ ion of the series has $Z - 20$ electrons in $3d$. Titanium has $Z = 22$, and $\text{Ti}^{2+}$ is $3d^2$. Check with chromium: $24 - 20 = 4$, and $\text{Cr}^{2+}$ is $3d^4$. It works because both $4s$ electrons are gone and the other $Z - 20$ electrons beyond argon sit in $3d$.

Try the atomic numbers 24 and 29. Then set iron, $Z = 26$, to charge 2 and to charge 3, and see which box loses electrons. The count of unpaired electrons matters in the third lesson, where it decides the magnetism of an ion.

Practise with new numbers

Try it in Python

The program builds the configuration of each element of the series from the filling order and the two exceptions. It also tests the rule for a transition element.

Common mistakes

• Counting zinc as a transition element. Its d subshell is full in the atom and in $\text{Zn}^{2+}$.
• Taking the electrons for $\text{Fe}^{2+}$ from $3d$ first. The $4s$ electrons go first, so $\text{Fe}^{2+}$ is $3d^6$ and not $3d^4 4s^2$.
• Writing chromium as $3d^4 4s^2$ and copper as $3d^9 4s^2$. Each keeps one electron in $4s$.
• Leaving out the $4s$ electrons when you count the unpaired electrons of an atom. Chromium has $3d^5 4s^1$, so six unpaired electrons.

Remember these

• The d block is groups 3 to 12, in four series: 3d, 4d, 5d and 6d.
• A transition element has an incomplete d subshell in its atom or in a common ion. Zinc, cadmium and mercury do not.
• The 3d series fills $3d$ from $d^1$ to $d^{10}$ with $4s^2$, except Cr ($3d^5 4s^1$) and Cu ($3d^{10} 4s^1$).
• Ions lose $4s$ electrons first. Every $\text{M}^{2+}$ ion of the series has $Z - 20$ electrons in $3d$.

Check yourself

Try each one on your own first. If you are stuck, open a hint. Then check with Show the answer.

1. Zinc, cadmium and mercury belong to the d block. Why are they not counted as transition elements?

2. Write the configuration of $\text{Fe}^{3+}$ and say how many unpaired electrons it has.

3. Palladium has atomic number 46. The filling order would give $[\text{Kr}]\,4d^8 5s^2$. What is the real configuration?

4. Which has more unpaired electrons, the chromium atom or the $\text{Cr}^{3+}$ ion?

5. An $\text{M}^{2+}$ ion of the 3d series has eight electrons in $3d$. Which element is it?