20.1
Properties of Transition Metals
Transition metals are defined as those elements that have partially filled d orbitals. As shown in Figure 1, the d-block elements in groups 3–12 are t…
Transition metals are a group of elements located in the d-block. They are placed between the main group elements and exhibit unique characteristics, such as electrical conductivity, colors, hardness, high melting points, and magnetism.
Many of these characteristics are attributed to the filling of d orbitals. Moving across the period, electrons are added following the Aufbau principle. However, the outermost s orbital is filled first, before the electron enters the (n − 1) d subshell, with some exceptions. Chromium and copper have a half-filled s orbital because it is energetically more favored to have a half-filled or full d subshell.
The atomic size of transition metals decreases slightly across a period. As the electrons enter the d subshell, the number of valence electrons in the outermost s orbital remains steady, leading to an overall constant nuclear charge. Down the column, the trend increases from the fourth to the fifth period, but not below.
The 6th period contains 14 additional elements of the lanthanide series, where electrons enter the (n − 2)f subshell before entering the (n − 1)d subshell. Electrons located in the f orbitals do not sufficiently shield the valence electrons in the outermost s orbital, resulting in an increased effective nuclear charge. Thus, the valence electrons are attracted more strongly, leading to smaller than expected atomic radii. This is known as the Lanthanide contraction.
Similarly, the transition elements’ electronegativity is influenced by its atomic size and nuclear charge. Moving across a period, it increases only slightly. Down the column, the electronegativity increases from the fourth to the fifth period but remains constant in the sixth period. This is attributed to a decrease in atomic size and an increase in the nuclear charge.
The ionization energy of transition metals increases slightly across a period. Interestingly, down the column, it is higher for the 6th period than the 4th and 5th periods. Here, the valence electrons are held more tightly as the nuclear charge increases with a negligible increase in atomic size.
Lastly, transition metals can have a range of stable oxidation states by losing outermost s orbital electrons before the unpaired d orbital electrons. The partially filled d orbitals enable oxidation states ranging from +7 to +1, characteristic colors, and magnetic properties.
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Q1: Why do transition metals have unique properties like electrical conductivity and high melting points?
Transition metals' unique properties stem from their partially filled d orbitals. The filling of d orbitals contributes to their electrical conductivity, hardness, high melting points, and magnetism. These characteristics distinguish transition metals from main group elements and enable their diverse chemical behaviors and applications in various materials and compounds.
Q2: How does the lanthanide contraction affect transition metal properties?
The lanthanide contraction occurs because f-orbital electrons poorly shield valence electrons from nuclear charge in the sixth period. This increased effective nuclear charge pulls valence electrons closer, resulting in smaller-than-expected atomic radii. Consequently, transition metals in the sixth period have higher ionization energies and increased electronegativity compared to lighter periods, despite larger atomic sizes.
Q3: Why can transition metals form multiple stable oxidation states?
Transition metals form multiple stable oxidation states because they lose outermost s orbital electrons before unpaired d orbital electrons. The partially filled d orbitals enable oxidation states ranging from +1 to +7, allowing these metals to exist in several different stable forms. This flexibility in oxidation states is central to their rich and fascinating chemistry.
Q4: What causes the color of transition metal compounds?
Electron transitions in transition metals can absorb photons in the visible electromagnetic spectrum, producing colored compounds. These colored hydrated ions, such as Cr3+, Fe3+, and Co2+, result from d-orbital electron transitions. The variety of colors observed in transition metal compounds reflects their complex valence shell structure and electronic properties.
Q5: How does atomic size change across and down the transition metal series?
Atomic size decreases slightly across a period as d electrons fill, because the outermost s orbital remains constant while nuclear charge increases. Down a group, atomic size increases from the fourth to fifth period but not below. The sixth period shows smaller-than-expected atomic radii due to the lanthanide contraction, where f-orbital electrons inadequately shield valence electrons.
Q6: What is the electron configuration pattern for transition metals?
Transition metals follow the Aufbau principle, with the outermost s orbital filling before electrons enter the (n − 1) d subshell. Exceptions like chromium and copper have half-filled or full d subshells because this configuration is energetically favored. The sixth period contains 14 lanthanide elements where electrons enter the (n − 2) f subshell before the (n − 1) d subshell.
Q7: How do transition metals differ in reactivity and oxidation state stability across the series?
Early transition metals like scandium readily form stable cations and are strong reducing agents, while later elements like platinum and gold resist oxidation due to higher reduction potentials. Elements in the second and third transition series generally stabilize higher oxidation states than first-series metals. Heavier d-block elements often form oxyanions rather than simple aqueous cations, reflecting their increased stability in higher oxidation states.