The atomic radii of elements in Period 3 are generally larger than those in Period 2. This difference is primarily due to the addition of a new electron shell in the elements of Period 3. For instance, elements in Period 2, such as carbon or oxygen, have their valence electrons in the second shell, while elements in Period 3, like silicon or sulphur, have their valence electrons in the third shell. As the principal quantum number (n) increases, the distance of the valence shell from the nucleus also increases, resulting in a larger atomic radius. Furthermore, although the nuclear charge increases across each period, the increased distance and the added inner electron shells in Period 3 elements provide a greater shielding effect, reducing the effective nuclear charge experienced by the outermost electrons. This reduction in effective nuclear charge further contributes to the larger atomic radii observed in Period 3 compared to Period 2.

The ionic radii of elements in Period 3 differ from their atomic radii primarily due to the gain or loss of electrons during ion formation. When an atom forms a cation by losing electrons, its radius decreases. This is because the loss of one or more valence electrons reduces electron-electron repulsion and often results in the loss of an entire electron shell, making the ion smaller than its neutral counterpart. Conversely, when an atom gains electrons to form an anion, its radius increases due to increased electron-electron repulsion within the added electron shell. These changes in ionic radii significantly affect the chemical properties of the ions. Smaller cations, for example, have a stronger attraction to nearby anions, leading to the formation of more compact and often stronger ionic bonds. In contrast, larger anions can more easily accommodate additional electrons, which can influence their reactivity and the nature of the compounds they form. Thus, the changes in ionic radii upon ion formation play a crucial role in determining the chemical behavior of elements.

Electronegativity, the ability of an atom to attract shared electrons in a covalent bond, is closely related to the trend in atomic radii observed across Period 3. As we move from left to right across Period 3, the atomic radii decrease due to the increasing effective nuclear charge. This increasing nuclear charge results in a stronger attraction of the nucleus for the valence electrons, both in the atom itself and in covalent bonds it forms. Consequently, as the atomic radius decreases, the electronegativity generally increases. For example, elements such as sodium and magnesium at the beginning of Period 3 have larger atomic radii and lower electronegativity, while elements like sulphur and chlorine towards the end of the period have smaller atomic radii and higher electronegativity. This relationship is crucial in understanding chemical reactivity and the nature of chemical bonds formed by elements, as elements with higher electronegativity tend to form more polar bonds, influencing the chemical and physical properties of the compounds they form.

Argon's atomic radius is significantly smaller than that of potassium, despite both elements residing in Period 3 of the periodic table. The key to understanding this difference lies in the electron configuration and the state of the elements. Argon is a noble gas with a complete outer electron shell, making its atomic structure highly stable and compact. This complete valence shell results in a strong effective nuclear charge, as the shielding effect is minimal. In contrast, potassium, which is the first element of the next period, has an additional electron shell compared to argon. This extra shell means that the outermost electrons are further away from the nucleus, leading to a larger atomic radius. Furthermore, potassium's additional electron shell offers more shielding, which reduces the effective nuclear charge experienced by the outermost electrons, contributing to its larger size.

The trend of decreasing atomic radius across Period 3 is not a universal trend applicable to all periods in the periodic table. While this trend holds true generally for elements within a given period, particularly for periods 2 and 3 where the added electrons are in the same principal energy level, it does not necessarily apply to periods involving transition metals or inner transition metals. In periods containing transition metals, for instance, the addition of electrons occurs in d-orbitals, which have a different shielding effect compared to s- and p-orbitals. This can lead to anomalies in the general trend of decreasing atomic radius. Additionally, in periods with inner transition metals (lanthanides and actinides), the filling of f-orbitals adds complexity to the trend due to the poor shielding effect of f-electrons. Therefore, while the trend of decreasing atomic radius from left to right is a useful guideline, it must be applied with an understanding of the specific electronic configurations and shielding effects of each period.