Question

Activity series are prepared by comparing single-displacement reactions between metals. Based on observations, the metals can be ranked by their ability to react. However, reactivity can be explained by the ease with which atoms of metals lose electrons. Using information from the activity series, identify the locations in the periodic table of the most reactive metals and the least reactive metals. Using your knowledge of electron configurations and periodic trends, infer possible explanations for the metals’ reactivity and position in the periodic table.

Short answer

Most reactive metals are in Groups 1 and 2 at the bottom; least reactive are transition metals like Au, Pt, Ag. Reactivity is due to ease of losing valence electrons.

Step-by-step solution

Introduction to Activity Series

Understand that the activity series of metals ranks them by their ability to undergo single-displacement reactions. Metals that lose electrons more easily are more reactive.

Identifying Reactive Metals

The most reactive metals are typically found in the alkali metals group (Group 1) and alkaline earth metals group (Group 2) of the periodic table. Examples include Lithium (Li), Sodium (Na), and Potassium (K) for Group 1, and Magnesium (Mg) and Calcium (Ca) for Group 2.

Identifying Less Reactive Metals

The least reactive metals, also known as noble metals, are found in the transition metals and near the bottom of the activity series. Examples include Gold (Au), Platinum (Pt), and Silver (Ag).

Electron Configuration and Reactivity

Reactive metals have electron configurations where the outermost electrons are easily lost. For alkali metals, the electron configuration ends in a single electron in the outermost s-orbital (e.g., Na: [Ne] 3s^1). Losing this electron requires low ionization energy, making the metal highly reactive.

Periodic Trends and Reactivity

As you move down a group in the periodic table, atomic radii increase and ionization energy decreases. This makes it easier for atoms to lose their valence electrons, increasing their reactivity. Thus, elements like Cesium (Cs) and Francium (Fr) are extremely reactive.

Summary of Positions and Reactivity

The most reactive metals are located at the bottom of Groups 1 and 2, where the ionization energy is the lowest. The least reactive metals are located among the transition metals, particularly towards the right side and bottom of the periodic table.

Key concepts

Reactivity of Metals

The reactivity of metals is their ability to participate in chemical reactions, often by losing an electron to form positive ions. Highly reactive metals easily lose their outermost electrons. For instance, alkali metals like Sodium (Na) and Potassium (K) readily lose their single valence electron. This makes them highly reactive.
Metals are ranked by their reactivity in what is known as the activity series. Metals at the top of the activity series, like Lithium (Li) and Calcium (Ca), are more reactive because they lose electrons effortlessly. Conversely, less reactive metals, such as Gold (Au) and Platinum (Pt), hold onto their electrons more firmly.
Reactive metals usually engage in single-displacement reactions, displacing less reactive metals from compounds. For example, Zinc (Zn) can displace Iron (Fe) from iron sulfate, showcasing higher reactivity.

Periodic Table Trends

Understanding periodic table trends is key to predicting metal reactivity. As you move from left to right across a period, elements have increasing nuclear charge, pulling electrons more tightly. This makes metals on the left side, like those in Group 1 and Group 2, more reactive because they lose electrons more easily.
As you go down a group, atomic size increases and ionization energy decreases. This makes the outer electrons easier to lose. For example, Cesium (Cs) is more reactive than Sodium (Na) due to its larger atomic radius and lower ionization energy.
Metals in the top right corner of the periodic table, particularly noble metals like Gold (Au) and Silver (Ag), are least reactive. This is because their high ionization energies make electron loss more difficult.
Hence, the periodic table not only organizes elements but also helps us predict and compare their reactivity based on their position.

Electron Configurations

Electron configurations describe the arrangement of electrons in an atom's orbitals. For reactive metals, the configuration shows outer electrons that are easy to remove. For example, Sodium (Na) has the electron configuration [Ne] 3s^1. The single s-electron in the outer shell is easily lost, making it highly reactive.
In contrast, less reactive metals have configurations that stabilize their outer electrons. For example, Platinum (Pt) has the configuration [Xe] 4f^14 5d^9 6s^1. The complex arrangement makes it harder for Platinum to lose electrons, resulting in lower reactivity.
Alkali metals (Group 1) and Alkaline Earth Metals (Group 2) have simpler electron configurations, ending in s-orbitals, which makes electron loss easier. For instance, Calcium (Ca), with the configuration [Ar] 4s^2, readily loses two electrons.
Through their electron configurations, we can infer the reactivity of metals—simple configurations with fewer electrons to lose indicate higher reactivity, while more complex configurations indicate lower reactivity.