Atomic radius decreases across a period due to increasing nuclear charge from "summary" of The Periodic Table: A Very Short Introduction by Eric R. Scerri
The atomic radius of an element refers to the distance from the nucleus to the outermost electron shell. Understanding the trend of atomic radius across a period on the periodic table can provide insights into the behavior of elements. As one moves from left to right across a period, the atomic radius generally decreases. This phenomenon can be explained by the concept of increasing nuclear charge. The nuclear charge refers to the positive charge in the nucleus of an atom, which is determined by the number of protons.
When moving across a period, the number of protons in the nucleus increases, leading to a stronger positive charge. This increase in nuclear charge exerts a greater force of attraction on the electrons in the outermost shell. As a result, the electrons are pulled closer to the nucleus, causing the atomic radius to decrease.
The decreasing atomic radius across a period can be visualized as the electrons being more tightly held by the nucleus. This concept is essential in understanding the trends in chemical reactivity and bonding behavior of elements. Elements with smaller atomic radii tend to form stronger bonds due to the closer proximity of electrons to the nucleus.
In summary, the trend of decreasing atomic radius across a period is a result of the increasing nuclear charge as one moves from left to right on the periodic table. This concept provides a foundational understanding of the interactions between electrons and the nucleus in different elements.
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