d-Block Elements (OCR A-Level Chemistry A): Revision Notes

d-Block Elements

What are d-block elements?

The d-block elements occupy a central region of the periodic table, positioned between Group 2 and Group 13. This section covers the first row of d-block elements, from scandium (Sc) to zinc (Zn) in Period 4.

These elements are characterized by having their highest energy electrons entering the 3d sub-shell as you progress across the period. This makes them distinct from the s-block elements on the left and the p-block elements on the right of the periodic table.

Physical properties of d-block elements

All d-block elements are metals and display typical metallic characteristics:

• High melting and boiling points - They require significant energy to overcome metallic bonding
• Shiny appearance - They have a lustrous metallic surface
• Good electrical conductivity - Free electrons can move through the metal structure
• Good thermal conductivity - They efficiently transfer heat energy

Electron configuration of d-block elements

The arrangement of electrons in atoms follows a specific order based on energy levels. Understanding this helps explain the chemical behavior of d-block elements.

Electron configurations from scandium to zinc

As we move across the first row of d-block elements from scandium to zinc, the 3d sub-shell gradually fills with electrons. Here is how this progresses:

Most elements follow a regular pattern where electrons fill the 4s sub-shell first (giving $4s^2$), and then the 3d sub-shell fills progressively. However, chromium and copper show exceptions to this pattern.

Special cases: chromium and copper

Electron configuration of d-block ions

When d-block elements form positive ions, they lose electrons. A crucial point is that 4s electrons are always removed before any 3d electrons, even though the 4s sub-shell fills first when forming atoms.

Transition elements

Not all d-block elements are classified as transition elements. There is a specific definition that must be satisfied.

Why scandium and zinc are not transition elements

Although scandium (Sc) and zinc (Zn) are both d-block elements, neither qualifies as a transition element:

Scandium:

• Electron configuration of Sc: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^1 4s^2$
• Scandium only forms Sc³⁺ ions by losing two 4s electrons and one 3d electron
• Electron configuration of Sc³⁺: $1s^2 2s^2 2p^6 3s^2 3p^6$
• The Sc³⁺ ion has empty d-orbitals (no 3d electrons)

Zinc:

• Electron configuration of Zn: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10} 4s^2$
• Zinc only forms Zn²⁺ ions by losing its two 4s electrons
• Electron configuration of Zn²⁺: $1s^2 2s^2 2p^6 3s^2 3p^6 3d^{10}$
• The Zn²⁺ ion has full d-orbitals (ten 3d electrons)

Since neither scandium nor zinc forms ions with partially filled d-orbitals, they do not meet the definition of transition elements.

Properties of transition metals

Transition elements exhibit three characteristic properties that distinguish them from other metals:

1. They form compounds with variable oxidation states
2. They form colored compounds
3. They can act as catalysts (both the elements and their compounds)

Variable oxidation states

Transition elements can form compounds where the metal exists in different oxidation states. This means they can lose different numbers of electrons when forming compounds.

For example, iron commonly forms two types of chloride:

• Iron(II) chloride (FeCl₂) - iron in +2 oxidation state
• Iron(III) chloride (FeCl₃) - iron in +3 oxidation state

All transition elements form compounds with an oxidation state of +2, which results from the loss of their two 4s electrons. As you move across the series from scandium to manganese, the number of possible oxidation states increases, reaching a maximum at manganese with seven possible states (+2 to +7). After manganese, the number of oxidation states decreases.

Common examples of oxidation states and their colors:

• Iron(II) compounds (Fe²⁺): pale green solutions (electron configuration $1s^2 2s^2 2p^6 3s^2 3p^6 3d^6$)
• Iron(III) compounds (Fe³⁺): yellow solutions (electron configuration $1s^2 2s^2 2p^6 3s^2 3p^6 3d^5$)
• Chromium(III) compounds (Cr³⁺): green solutions
• Chromium(VI) compounds (Cr⁶⁺): yellow or orange solutions

Determining oxidation states

You can calculate the oxidation state of a transition metal in a compound or ion using the rule that the sum of all oxidation numbers equals the overall charge.

Formation of colored compounds

Transition metal compounds and their aqueous solutions are frequently colored. This is one of the most visually distinctive properties of transition elements.

Different oxidation states of the same element produce different colors because they have different numbers of d-electrons:

The image shows the distinct color difference between iron in +2 oxidation state (pale green/colorless solution on the left) and iron in +3 oxidation state (yellow solution on the right).

Catalytic behavior

Transition metals and their compounds serve as catalysts in numerous chemical reactions. A catalyst is a substance that increases the rate of a reaction without being permanently changed or consumed itself. It works by providing an alternative reaction pathway with lower activation energy.

Catalysts can be classified into two types:

Another example of homogeneous catalysis is the reaction of zinc metal with sulfuric acid, which is catalyzed by Cu²⁺(aq) ions:

$\ce{Zn(s) + H2SO4(aq) -> ZnSO4(aq) + H2(g)}$