Introduction to transition metals and the d-block elements

In this lesson, we will learn:

• The definition of a transition metal as distinguished from a d-block element.
• The general properties of d-block and transition metal elements.
• How to write electron configurations for the transition metal elements and their common ions.

• Until now we have only really looked at the groups of elements around the edges of the periodic table, like the halogens, noble gases and alkali metals. These groups have very distinct properties that make them a bit easier to study and learn about.
  This module is all about the transition metals and the d-block, the middle section of the periodic table.


• The d-block elements have distinct properties and features just like the alkali metals and halogens. They arent always as clear-cut as the other groups, but generally they are:
• Variable oxidation states, such as Fe (II) and Fe (III)
• Coloured compounds, such as copper (II) or chromium (VI)
• Ability to form complex ions.


• Transition metals and d-block elements are NOT the same! The distinction between them is very important because it is the transition metal properties that are mistaken for all d-block element properties.
• A d-block element is any element found in the central d-block of the periodic table this is simple enough. If its ground-state (the neutral atom) electron configuration ends with a d-subshell (such as 3d^1-10), it is a d-block element.
• A transition metal is any d-block element that forms at least one stable ion with a PARTIALLY FILLED d-subshell..
So, all transition metals are d-block elements, but not all d-block elements are transition metals.


• The difference is easiest to explain with examples using electron configuration.
• Scandium is a d-block element but IS NOT a transition metal. This is because, despite its neutral atom having an outer shell configuration of 4s² 3d¹, its only stable ion is Sc³⁺, which removes all three of those outer shell electrons, leaving it with an [Ar] 4s⁰ 3d⁰ or just [Ar] electron configuration. This is an empty d-subshell, not a partially filled one! Therefore, scandium does not count.


• Zinc IS NOT a transition metal. Again, the neutral atom has an outer shell configuration ending with the d-subshell (3d¹⁰), but its only stable ion is Zn²⁺, which has the outer shell configuration 4s⁰ 3d¹⁰. This is completely filled, not partially! Therefore, zinc is not a transition metal.


• Iron IS a transition metal. Its ground-state electron configuration is 4s² 3d⁶, and it forms two stable ions:
• Fe²⁺ with the configuration 4s⁰ 3d⁶
• Fe³⁺ with the configuration 4s⁰ 3d⁵.
The ground state and both these stable ions satisfy the partially-filled criteria, so iron counts as a transition metal in the d-block.


• To easily distinguish transition metals in the d-block, you need to know the electron configurations for ground-state d-block elements and their common ions. There are a few guidelines to this:
• When filling subshells, the 4s subshell is filled first, then 3d. The 3d sub-shell is made up of 5 orbitals, for a total of 10 electrons. Like in the p-subshell, the electrons fill up a separate orbital first before pairing up.
• For example, reading across the beginning 1^st row of d-block elements we see the following pattern:


• When 1^strow transition metals become ions, the 4s² electrons are removed first.
• For example, Ti²⁺ has an ion electron configuration of [Ar] 4s⁰ 3d², as the 4s electrons are removed first.
This is why the +2 oxidation state is common for transition metals. If a transition metal can form an M²⁺ ion, it is probably by removing the 4s² electrons, without touching the 3d electrons.


• Many ions have electron configurations settled on a half-filled or completely empty d-subshell. If you have a metal ion of an unusual charge/oxidation state and need to suggest the configuration, try to remove electrons from the 3d subshell until it is half or completely empty.
• For example, after Fe (II), irons other main oxidation state is Fe (III), which removes the one paired electron of its 3d⁶ electrons to give [Ar] 4s⁰ 3d⁵ as its ion electron configuration.
• Manganese (VII) is possible because the ground state of Mn is 4s² 3d⁵. Manganese loses control of these seven when it becomes Mn (VII), the type used in a lot of redox reactions.


• Chromium and copper are special cases with their electron configurations. The ground-state configurations of chromium and copper respectively show a half-filled and completely-filled 3d subshell. They get this from the paired 4s subshell which becomes half-filled itself.

The reasons for this are complex; it is just easier to remember that these particular two elements are more stable with a 4s¹ 3d⁵ or 4s¹ 3d¹⁰ ground state.