Periodic trends increase more gradually across the d-block than the s- or p-blocks
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refutedsupported
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Reference sources report that atomic radii and other properties decrease or change only slightly across the d-block due to poor d-electron shielding, leading to more gradual periodic trends compared to main-group blocks.
This is the cause of introducing 4f electrons in Row 3. In Row 3, we would expect the elements to carry on the same trend as was witnessed between Rows 1 and 2 (large increase in atomic radii) but we do not. This is because the 4f orbitals are not doing a great job of shielding. D Block Contraction (Scandide Contraction)
The d block contraction, also known as the Scandide Contraction, describes the atomic radius trend that the d block elements (Transition metals) experience. Normally the trend for atomic radius, moving across the periodic table is that the atomic radius decreases significantly. In the transition metals with D electrons as we move from left to right across the periodic table, the element’s atomic radius only decreases slightly. This is because they have the same amount of s electrons, but are only differing in d electrons. These d electrons are in an inner shell (penultimate shell) and electrons are getting added to this shell, another shell is not created. The d electrons are not good at shielding the nuclear charge, so the atomic radius does not change much as electrons are added. Almost like disregarding the D electrons being added.
To understand why the d-metals are roughly the same size, we need to examine two competing effects: nuclear charge and electron-electron repulsions. As one moves across the periodic table, the size (that is, the atomic radius) of elements generally decreases due to increasing nuclear charge. The increasing nuclear charge is usually more significant than the electron-electron repulsions resulting from the addition of electrons as one moves across a period. In keeping with this trend, transition metals also generally decrease in size as one moves from left to right across the d-block. Here, the exceptionally weak shielding offered by the d-electrons causes these electrons to experience a greater nuclear charge and be pulled more strongly toward the nucleus. The last few elements in each row of the d-block are slightly larger than those preceding them; here, electron-electron repulsions outweigh increasing nuclear charge as the d-orbitals are filled. The trend of decreasing size is most pronounced in row 6 of the periodic table (the third row of the d-block, in which the 5d-orbitals are being filled).
The s-block includes Groups 1 and 2, plus hydrogen and helium. In these elements, the s orbitals are being filled with electrons. The p-block includes Groups 13 to 18. These elements are filling up their p orbitals. The s-block and p-blocks together make up most of the main group elements, which include both metals and nonmetals. The d-block is in the middle of the table and includes Groups 3 to 12. These are the transition metals. They are special because they can form different types of charged ions and complex compounds. The f-block is at the bottom of the table and includes the lanthanides and actinides. These elements are filling their f orbitals. Even though they are placed separately to keep the table neat, they actually belong in the 6th and 7th rows (periods). The periodic table also shows trends, or patterns, in the properties of elements. Atomic radius (the size of an atom) gets smaller across a period from left to right because the nucleus pulls electrons in tighter. But it gets larger going down a group because more electron shells are added. Electronegativity (how strongly an atom attracts electrons) usually increases across a period and decreases down a group.
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