Final answer:
Understanding atomic structure allows us to explain the trend in electronegativity across the periodic table; it increases from left to right due to a decreasing atomic size and stronger nuclear pull and decreases from top to bottom due to an increasing atomic size and a reduced nuclear pull on bonding electrons.
Step-by-step explanation:
Understanding atomic structure is crucial in explaining the periodic trend in electronegativity. Electronegativity is the tendency of an atom to attract bonding electrons, and varies across the periodic table. The atomic size, or covalent radius, and the effective nuclear charge play significant roles in this trend.
As you go from top to bottom in a column on the periodic table, electronegativity tends to decrease. This is because the atomic size increases, resulting in a greater distance between the bonding electrons and the nucleus. This increased distance diminishes the nucleus's pull on the bonding electrons.
Conversely, as you move from left to right across a row, electronegativity generally increases. This is because the atomic size decreases due to a higher effective nuclear charge. As more protons are added to the nucleus without a significant increase in shielding electrons, the pull on the bonding electrons becomes stronger, thereby increasing electronegativity.
Therefore, the electron configurations provide insight into the attractive force an atom can exert on electrons. This helps us understand why atoms with higher electronegativities, such as fluorine, can attract electrons more effectively, compared to elements like cesium, which have lower electronegativities.