How does electronegativity generally change across a period in the periodic table?

Electronegativity is a measure of an atom's ability to attract and hold onto electrons when it is part of a compound. As you move from left to right across a period in the periodic table, the electronegativity generally increases. This increase is primarily due to the rise in the number of protons in the nucleus of the atoms within that period. With a higher proton count, the positive charge in the nucleus becomes stronger, which enhances its ability to attract the negatively charged electrons in a bond.

Additionally, as you move across a period, the atomic radius typically decreases due to increased nuclear charge without a corresponding increase in electron shielding. A smaller radius means that the electrons being attracted are closer to the nucleus, further increasing electronegativity. As a result, elements on the right side of the periodic table, such as fluorine and oxygen, tend to have much higher electronegativities than those on the left, such as sodium and magnesium.

This trend is fundamental in understanding chemical bonding and reactivity, as elements with high electronegativity are more likely to attract electrons and form polar covalent or ionic bonds.