An active metal in chemistry is characterized by its high degree of chemical instability and strong tendency to react quickly and vigorously with other substances. This high reactivity differentiates active metals, such as sodium and potassium, from stable metals like iron or gold. Understanding this characteristic is fundamental to safely working with these unique elements.
Defining Active Metals and Their Location
An active metal is formally defined as an element that is easily oxidized, meaning it readily gives up one or more electrons during a chemical reaction. Losing electrons transforms the neutral metal atom into a positively charged ion, or cation. The ease of this conversion dictates the metal’s place on the scale of chemical activity.
The most active metals are consistently found on the far left side of the periodic table, specifically within the first two main columns: Group 1 (Alkali Metals) and Group 2 (Alkaline Earth Metals). Elements in these groups, including lithium, sodium, and potassium, share physical properties that hint at their reactivity, such as being relatively soft and having low densities compared to most other metals.
The Chemical Mechanism of High Reactivity
The high activity of these metals stems from their electron configuration. Atoms seek a stable, full outer electron shell. Metals in Group 1 and Group 2 possess only one or two valence electrons, respectively, making it energetically favorable to lose these few electrons to reveal a full, stable inner shell.
The energy required to remove an electron is measured as ionization energy. Active metals exhibit significantly low ionization energy, meaning only a small amount of energy is needed to strip away their valence electrons. This low energy requirement translates directly into high chemical reactivity because the atom is eager to shed the electron.
Because they readily donate electrons to other elements, active metals function as powerful reducing agents in chemical reactions. A reducing agent causes another substance to gain electrons, while the reducing agent itself is oxidized.
Comparing Activity: The Reactivity Series
To quantify and compare the relative activity of different metallic elements, chemists use the Metal Reactivity Series, also known as the Activity Series. This ranked list places metals in descending order of their chemical reactivity; the most active metals, which lose electrons most easily, are placed at the top.
The primary function of the Reactivity Series is to predict the outcome of a displacement reaction. A metal ranking higher on the list will displace a less active metal from its compound. For instance, a more active metal, like zinc, can displace copper from an aqueous solution of a copper salt.
In this reaction, the zinc metal readily gives up its electrons to the copper ions, causing the copper to precipitate as a solid metal. This observable chemical change confirms the relative positions of the two metals on the series. The Reactivity Series is therefore an experimental tool that translates the theoretical concept of electron-loss tendency into a practical, predictive ranking.
Common Active Metals and Safe Handling
Specific examples of highly active metals include lithium, sodium, and potassium. These elements react so vigorously that they must be handled with extreme caution, as their high activity means they react instantly with oxygen and moisture in the atmosphere.
The reaction with water is particularly violent, often releasing hydrogen gas and generating substantial heat that can cause a fire or explosion. Due to this extreme reactivity, these elements cannot be stored in the open air. They are commonly stored immersed completely under an inert substance, such as mineral oil or kerosene, to prevent contact with air or moisture.
In controlled environments, highly pure active metals are often handled within a glove box, a sealed enclosure filled with an inert gas like argon. This practice ensures the metal remains in an oxygen-free and moisture-free atmosphere.