Introduction

Zinc is a chemical element with the symbol Zn, and its atomic number is 30. It is a slightly brittle metal at room temperature and has a shiny greyish appearance when oxidation is removed. It is the periodic table’s first element in group 12 (2nd B). In some respects, zinc is chemically similar to magnesium; both elements have the same oxidation number (+2).

Zinc plays a vital role chemically and biologically. Zinc is an important mineral for humans, plants, animals, and microorganisms. Zinc is also an essential nutrient; it plays an enzymatic action. Zinc is an essential trace element in the human body, where it is found in higher concentrations in red blood cells as a necessary part of the enzyme carbonic anhydrase.

A little more abundant zinc makes up an average of 65 grams (2.3 ounces) of every ton of earth’s crust. The chief zinc mineral is sulfide sphalerite (zinc blend).

The early 21st-century zinc producers are Australia, China, and Peru.

A Brief History of Zinc Metal

Humans have worked with zinc for at least 2500 years because it’s a critical alloying element in BrassBrass. Zinc itself was not recognized as a distinct element until much later. In the 1400s, calamine and wool were heated to obtain metallic zinc in India. Andreas Sigismund Marggraf discovered zinc again in 1746 when he heated calamine with charcoal. This traditional method differs from the electrolysis of aqueous zinc sulfate (ZnS04) used in modern zinc production.

Historically, “spelter” was used interchangeably with “zinc,” though technically, spelter is a zinc-lead alloy. It’s also known by other terms like “galvanizing metal,” “blende,” and “calamine.”

Properties Of Zinc

Physical Properties

Appearance: 

Zinc is a bluish-white metal with a shiny finish when freshly cut.

Density:

 The density of zinc is approximately 7.14 g/cm³, making it relatively light compared to other metals. This makes zinc relatively dense compared to other metals like aluminium (2.7 g/cm³) but less dense than metals like lead (11.34 g/cm³) and iron (7.87 g/cm³). The density of zinc can vary slightly depending on its specific alloy form, but pure zinc generally exhibits this density.

Melting and Boiling Points:

 Zinc has a melting point of 419.5 °C (787.1 °F) and changes from solid to liquid, and the temperature where zinc changes from liquid to gas at a boiling point is 907 °C (1,665 °F).

Zinc’s boiling point is much higher than its melting point, indicating that it can remain in a liquid state at high temperatures before it vaporizes.

Malleability and Ductility:

Zinc is moderately malleable at room temperature. This means it can be hammered or rolled into thin sheets without breaking but becomes brittle at temperatures below 100°C.

Zinc’s malleability increases when heated, particularly in the 100–150°C range, where it becomes more easily shaped and worked. This property is useful in manufacturing and forming zinc sheets and products.

2. Chemical Properties

Reactivity

Zinc readily reacts with acids to produce hydrogen gas and zinc salts. It also reacts with bases and can oxidize in air and moisture, forming a protective zinc oxide layer. Zinc is the most stable metal because its electrons are paired; no electron is unpaired. It has a filled d-shell, and it is diamagnetic and mostly colourless.

 Alloys

There are different types of zinc alloys known the following,

1. Brass
2. Zinc-aluminum alloys
3. Zinc-copper alloys
4. Zinc-nickel alloys
5. Zinc-steel alloys

Zinc can easily form alloys with other metals, such as copper, aluminium, and magnesium, enhancing their properties.

For example, BrassBrass is an alloy of zinc and copper. It contains 60-70% copper and 30-40% zinc. It is used in plumbing, musical instruments, coins, and decorative items.

Corrosive resistant

Zinc forms a protective layer of zinc oxide on its surface, which protects it from further corrosion when exposed to air. This property is used in galvanizing to coat other metals. This layer is a barrier, preventing further oxidation and protecting the metal beneath. In addition, zinc is often used to coat other metals, a galvanization process that significantly extends the life of steel and iron structures. Even when the protective layer is scratched, zinc corrodes sacrificially, offering continued protection to the underlying metal. As a result, zinc is widely used in outdoor and marine applications for its durability against corrosion.

Solubility

Pure zinc metal itself is not soluble in Water. However, when zinc reacts with acids or alkalis, it forms soluble compounds. For example, zinc reacts with hydrochloric acid (HCl) to form zinc chloride (ZnCl₂), which is highly soluble in Water. Inc does not dissolve in most organic solvents, but certain organic compounds like zinc acetate or zinc chloride can dissolve in some.

Thus, while metallic zinc is not soluble in Water, many of its salts, such as zinc sulfate and zinc chloride, are highly soluble. 

Reactions of zinc

• Reactions with acids

Reaction with Hydrochloric Acid (HCl): Zinc reacts with hydrochloric acid to produce hydrogen gas and zinc chloride.

                          Zn(s)+2HCl(aq)→ZnCl2​(aq)+H2​(g)

• Reaction with sulfuric acids:

Zinc reacts with sulfuric acid to produce hydrogen gas and zinc sulfate.

                              Zn(s)+H2SO4​(aq)→ZnSO4​(aq)+H2​(g)

• Reaction with Oxygen

Reaction with Oxygen (Oxidation): Zinc reacts slowly with oxygen in the air, forming zinc oxide (ZnO), a white powder that protects the zinc from further oxidation.

                               2Zn(s)+O2​(g)→2ZnO(s)

Zinc oxide is a protective layer, making zinc resistant to further corrosion.

• Reaction With Water (at high temperature)

Zinc does not react significantly with cold or warm Water but reacts with steam at high temperatures to produce zinc oxide and hydrogen gas.

                                Zn(s)+H2​O(g)→ZnO(s)+H2​(g)

• Reaction with Carbon

Zinc reacts with carbon at high temperatures to form zinc vapour and carbon monoxide.

                                Zn(s)+C(s)→Zn(g)+CO(g)

• Zinc as a Reducing Agent

As a reducing agent, zinc loses electrons in a reaction. This process is known as oxidation, where zinc is oxidized (i.e., its oxidation state increases from 0 to +2. When zinc loses electrons, they are transferred to a reduced substance (i.e., its oxidation state decreases). The substance being reduced gains electrons.

Zinc is a relatively strong reducing agent, especially in aqueous solutions where it readily loses electrons to form zinc ions (Zn²⁺).

It is widely used in electrochemical reactions (e.g., galvanic cells, batteries) and in industrial processes (e.g., reducing other metals from their ores).

Zinc’s ability to donate electrons makes it a valuable material in electroplating, catalytic reactions, and battery technologies.

Example:

• In Displacement Reactions:

Zinc can displace less reactive metals from their compounds. For instance, in the reaction with copper sulfate:

                               Zn(s)+CuSO4​(aq)→ZnSO4​(aq)+Cu(s)

Electrochemical properties

Standard electrode potential

Zinc has a standard electrode potential of -0.76 V (relative to the standard hydrogen electrode, SHE).

This negative potential indicates that zinc is a good reducing agent and is readily oxidized, meaning it can easily lose electrons to other substances.

In a typical electrochemical cell, zinc tends to act as the anode, where it undergoes oxidation:

                               Zn(s)→Zn2+(aq)+2e−

Corrosion resistance (Galvanic protection)

Galvanization is when zinc coats iron or steel to prevent rusting. Zinc acts as a sacrificial anode in the galvanic cell. Because zinc is more reactive (has a more negative electrode potential), it corrodes first, protecting the underlying metal from oxidation.

The zinc metal undergoes oxidation and releases electrons:

                                   Zn(s)→Zn2+(aq)+2e−

Zinc in Batteries (Electrochemical cells)

Zinc-based batteries use zinc as the anode to generate electrical power. A common example is the zinc-carbon battery or the alkaline battery.

The zinc anode undergoes oxidation in these batteries, releasing electrons that flow through an external circuit to the cathode, where reduction occurs. The typical reaction at the anode in a zinc-carbon battery is:

                                    Zn(s)+2e−→Zn2+(aq)

The reaction occurs at the cathode, typically involving compounds such as manganese dioxide:

                                     MnO2​(s)+H2​O(l)+2e−→Mn(OH)2​(s)

Zinc as anode

Zinc is commonly used as an anode in galvanic (voltaic) cells because it tends to lose electrons and undergo oxidation. This makes it a good material for generating electrical energy.

In a simple galvanic cell:

Anode (Zinc): Zinc metal oxidizes to zinc ions (Zn²⁺) and releases electrons.

                                       Zn(s)→Zn2+(aq)+2e−

Cathode:

 The electrons from the zinc anode flow through an external circuit to the cathode, where reduction occurs (usually with a metal ion or compound, such as copper ions in a copper-zinc cell).

Zinc in Electrochemical Reactions

Zinc’s electrochemical reactivity makes it valuable in numerous industrial electrochemical processes, including:

Electrorefining: Zinc is used in electrorefining processes to purify other metals.

Electrochemical reduction: Zinc is used in certain chemical syntheses as a reducing agent.