The d-block contains elements whose highest energy electrons occupy d sub-shells. However, not all d-block elements are classified as transition metals. In this lesson, we will establish the precise definition of a transition metal, analyze electron configurations, and outline their characteristic chemical properties.
A transition metal is a d-block element that forms at least one stable ion with an incomplete (partially filled) d sub-shell.
A complex ion is a species consisting of a central metal ion surrounded by ligands bonded via coordinate (dative covalent) bonds.
Electron Configurations of the d-Block Elements
When writing the electron configurations of transition metals from Period 4 (Titanium to Copper), we fill orbitals in order of increasing energy. The 4s orbital is filled before the 3d orbital because the 4s is at a lower energy level in neutral atoms. We can use the noble gas core shorthand [Ar] to represent the inner 18 electrons (\(1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6\)).
🔑 Key Principle: Anomalous Configurations of Cr and Cu
Chromium and copper show anomalies in their ground-state electron configurations due to the extra stability associated with half-filled and fully-filled 3d sub-shells:
- Chromium (Z = 24): \([\text{Ar}]3\text{d}^5 4\text{s}^1\) (rather than \([\text{Ar}]3\text{d}^4 4\text{s}^2\))
- Copper (Z = 29): \([\text{Ar}]3\text{d}^{10} 4\text{s}^1\) (rather than \([\text{Ar}]3\text{d}^9 4\text{s}^2\))
🔑 Key Principle: Formation of Ions
When transition metal atoms lose electrons to form positive ions (cations), the 4s electrons are lost before the 3d electrons. For example:
- Iron atom: \(\text{Fe} = [\text{Ar}]3\text{d}^6 4\text{s}^2\)
- Iron(II) ion: \(\text{Fe}^{2+} = [\text{Ar}]3\text{d}^6\) (4s electrons lost)
- Iron(III) ion: \(\text{Fe}^{3+} = [\text{Ar}]3\text{d}^5\) (4s and one 3d electron lost)
Writing configurations of ions is a common trap. If you write the configuration of \(\text{Fe}^{2+}\) as \([\text{Ar}]3\text{d}^4 4\text{s}^2\), you will lose the mark. You must always remove the 4s electrons first: \(\text{Fe}^{2+}\) is \([\text{Ar}]3\text{d}^6\).
Why Scandium and Zinc are NOT Transition Metals
Scandium and zinc sit at the edges of the d-block, and neither meets the formal definition of a transition metal:
- Scandium (Sc): Scandium only forms one stable ion, \(\text{Sc}^{3+}\). To form this ion, the scandium atom (\([\text{Ar}]3\text{d}^1 4\text{s}^2\)) loses all three outer electrons, resulting in the electron configuration \([\text{Ar}]\) or \([\text{Ne}]3\text{s}^2 3\text{p}^6 3\text{d}^0\). Because its only stable ion has an empty d sub-shell (\(3\text{d}^0\)), scandium is not a transition metal.
- Zinc (Zn): Zinc only forms one stable ion, \(\text{Zn}^{2+}\). The zinc atom (\([\text{Ar}]3\text{d}^{10} 4\text{s}^2\)) loses its two 4s electrons to form \(\text{Zn}^{2+} = [\text{Ar}]3\text{d}^{10}\). Because its only stable ion has a completely filled d sub-shell (\(3\text{d}^{10}\)), zinc is not a transition metal.
If asked why zinc is not a transition metal, you must state: Zinc forms only the stable ion \(\text{Zn}^{2+}\), which has a full d sub-shell (\(3\text{d}^{10}\)). You must state the formula of the stable ion and specify its electron configuration to secure the marks.
General Properties of Transition Metals
Transition metals exhibit four characteristic physical and chemical properties that distinguish them from s-block metals like calcium:
1. Variable Oxidation States
Unlike Group 2 metals which only form +2 ions, transition metals form stable ions in a variety of oxidation states. For example, manganese can exist in oxidation states ranging from +2 in \(\text{Mn}^{2+}\) to +7 in the manganate(VII) ion, \(\text{MnO}_4^-\).
This is because the energy levels of the 3d and 4s orbitals are very close. The energy required to remove successive electrons is relatively small and gradual, allowing varying numbers of d electrons to participate in bonding without causing chemical instability.
2. Formation of Coloured Ions
Solutions containing transition metal ions are often intensely coloured. In contrast, solutions of s-block compounds (e.g. \(\text{NaCl}\), \(\text{CaCl}_2\)) and zinc compounds are colourless.
How Colour Arises
When ligands coordinate to a transition metal ion, their lone pairs repel the metal d electrons. This splits the five degenerate (equal-energy) d orbitals into two subsets of different energy levels separated by an energy gap, \(\Delta E\).
When visible light passes through the solution, d electrons in the lower energy level absorb a specific frequency of light. The energy of the absorbed photon corresponds to the energy gap: \[\Delta E = h\nu = \frac{hc}{\lambda}\] where \(h\) is Planck's constant, \(\nu\) is frequency, \(c\) is the speed of light, and \(\lambda\) is wavelength. This absorption promotes an electron to a higher d orbital (a d-to-d transition). The colour we observe is the complementary colour: the wavelengths of visible light that were not absorbed.
🔑 Key Principle: Why Sc³⁺ and Zn²⁺ are Colourless
To exhibit colour, there must be space in the upper d subset to accept an electron, and there must be electrons in the lower d subset to be promoted:
- \(\text{Sc}^{3+}\) has a \(3\text{d}^0\) configuration. There are no d electrons to promote, so it is colourless.
- \(\text{Zn}^{2+}\) has a \(3\text{d}^{10}\) configuration. The d sub-shell is full. There is no empty space in the upper orbitals to receive a promoted electron, so it is colourless.
3. Catalytic Activity
Transition metals and their compounds are extremely important industrial catalysts. They can act as heterogeneous or homogeneous catalysts:
- Heterogeneous catalysts: The catalyst is in a different physical phase from the reactants. Transition metals have partially vacant d orbitals that allow reactant molecules to adsorb onto their surface, weakening their covalent bonds and lowering the activation energy. Examples include:
- Iron (\(\text{Fe}\)): Catalyses the production of ammonia from nitrogen and hydrogen in the Haber process.
- Nickel (\(\text{Ni}\)): Catalyses the hydrogenation of vegetable oils to make margarine.
- Vanadium(V) oxide (\(\text{V}_2\text{O}_5\)): Catalyses the oxidation of sulfur dioxide to sulfur trioxide in the Contact process.
- Homogeneous catalysts: The catalyst is in the same phase as the reactants. Transition metals can act as homogeneous catalysts because they readily switch between oxidation states to provide alternative reaction pathways. Examples include:
- Manganese(IV) oxide (\(\text{MnO}_2\)): Catalyses the decomposition of hydrogen peroxide (\(\text{H}_2\text{O}_2\)).
4. Complex Ion Formation
Transition metal ions have a high charge density and empty, low-lying d orbitals. This makes them highly attractive to ligands (molecules or ions with lone pairs of electrons). The ligands donate their lone pairs into the empty d orbitals of the transition metal ion, forming coordinate bonds and creating complex ions.
1. State the full ground-state electron configuration of an iron atom (\(\text{Fe}\)), an iron(II) ion (\(\text{Fe}^{2+}\)), and an iron(III) ion (\(\text{Fe}^{3+}\)).
2. Use these configurations to explain why iron is classified as a transition metal.
Step 1: Write the configurations
- Iron atom (Z = 26): \(1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^6 4\text{s}^2\)
- \(\text{Fe}^{2+}\): \(1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^6\) (the two 4s electrons are lost first)
- \(\text{Fe}^{3+}\): \(1\text{s}^2 2\text{s}^2 2\text{p}^6 3\text{s}^2 3\text{p}^6 3\text{d}^5\) (a 3d electron is lost after the 4s electrons)
Step 2: Justify transition metal classification
By definition, a transition metal must form at least one stable ion with a partially filled d sub-shell. Both \(\text{Fe}^{2+}\) (which contains six 3d electrons) and \(\text{Fe}^{3+}\) (which contains five 3d electrons) have incomplete 3d sub-shells (i.e. between 1 and 9 electrons). Therefore, iron is classified as a transition metal.
Step 1: Write the configurations
- Zinc atom (Z = 30): \([\text{Ar}]3\text{d}^{10} 4\text{s}^2\)
- Zinc ion (\(\text{Zn}^{2+}\)): \([\text{Ar}]3\text{d}^{10}\)
Step 2: Apply the definition
The definition of a transition metal requires the formation of at least one stable ion with a partially filled d sub-shell. Zinc only forms one stable ion, \(\text{Zn}^{2+}\), which has a completely filled d sub-shell containing 10 electrons (\(3\text{d}^{10}\)). Because it does not form any stable ion with a partially filled d sub-shell, zinc is not classified as a transition metal.
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