Final answer:
Sulfur (S) is most likely to attract electrons in a bond due to its higher electronegativity, being further to the right and up on the periodic table compared to Al, Si, and P, and due to stronger d-orbital interactions in multiple bonds.
Step-by-step explanation:
The atom most likely to attract electrons in a bond among Al (aluminum), Si (silicon), P (phosphorus), or S (sulfur) is sulfur (S). This is because of the concept known as electronegativity, which is the measure of how strongly atoms attract bonding electrons to themselves. The periodic trends indicate that electronegativity increases as you move from left to right across a period and decreases as you go down a group. Consequently, within the same period, sulfur, being further to the right, is more electronegative compared to aluminum and silicon. While phosphorus is also to the right of aluminum and silicon, sulfur is in group 16 and has a more negative electron affinity compared to phosphorus in group 15, which generally has electron affinities that are less negative than expected.
Additionally, multiple bonds between atoms like carbon, oxygen, or nitrogen and a period 3 element such as sulfur can be unusually strong and dominating in chemistry, particularly due to possible d-orbital interactions. It also implies a greater tendency to form stronger bonds due to a higher electronegativity that aids in attracting electrons more effectively.