Predicting and identifying reactions and products Study Guide

The reactivity of a metal refers to how readily it undergoes chemical reactions, particularly with other elements or compounds like oxygen, water, or acids. More reactive metals tend to lose electrons more easily to form positive ions. We can deduce an order of reactivity for different metals by observing their reactions under controlled conditions. For instance, comparing reactions with water: very reactive metals like potassium and sodium react vigorously with cold water, producing hydrogen gas and a metal hydroxide. Less reactive metals like magnesium will only react with steam, while even less reactive metals like copper and silver do not react with water at all. Another common method involves observing displacement reactions. A more reactive metal will displace a less reactive metal from a solution of its salt. For example, if zinc is placed in a copper sulfate solution, the more reactive zinc will displace copper, forming zinc sulfate and solid copper. No reaction would occur if copper were placed in a zinc sulfate solution, as copper is less reactive than zinc. By systematically testing metals against these reagents or against each other in displacement reactions, we can construct a reactivity series, typically listing metals from most reactive to least reactive. This series helps predict the outcomes of various chemical reactions and is fundamental to understanding extraction processes and corrosion.

The Periodic Table is a powerful tool for predicting the chemical behaviour of elements. Elements are arranged in order of increasing atomic number (number of protons) and in groups and periods based on their electron configurations. Elements in the same group (vertical column) have similar chemical properties because they have the same number of outer shell electrons. These outer electrons are involved in chemical reactions.

For metals, reactivity generally increases down a group and decreases across a period from left to right. This is because metals react by losing electrons. Down a group, the outer electrons are further from the nucleus and are shielded by more inner electron shells, making them easier to lose. Across a period, the nuclear charge increases, pulling the outer electrons more strongly, making them harder to lose. For example, Group 1 metals (alkali metals) are highly reactive, and reactivity increases from lithium to francium.

For non-metals, reactivity generally decreases down a group and increases across a period from left to right. Non-metals typically react by gaining or sharing electrons. Down a group, the atomic radius increases, and the outer electrons are further from the positively charged nucleus, reducing the attraction for incoming electrons. Across a period, the nuclear charge increases, and the atomic radius decreases, increasing the attraction for incoming electrons. For example, Group 17 non-metals (halogens) are highly reactive, and reactivity decreases from fluorine to iodine. Observing the positions of elements in the Periodic Table allows us to make informed predictions about how they will react with other substances, such as which element will displace another in a reaction.