When zinc metal is submerged into a quantity of aqueous $$\ce{HCl}$$, the following reaction occurs (Figure $$\PageIndex{1}$$):

\<\ce{Zn (s) + 2HCl (aq) → H2 (g) + ZnCl2(aq)} \label{eq:1}\>

This is one example of what is sometimes called a single replacement reaction because $$\ce{Zn}$$ replaces $$\ce{H}$$ in combination with $$\ce{Cl}$$.

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Example $$\PageIndex{1}$$: Reducing Silver Ions

Write and balance the redox reaction that has silver ions and aluminum metal as reactants and silver metal and aluminum ions as products. Identify the substance oxidized, substance reduced, reducing agent and reducing agent.

Solution

We start by using symbols of the elements and ions to represent the reaction:

\<\ce{Ag^{+} + Al → Ag + Al^{3+}} \nonumber\>

The equation looks balanced as it is written. However, when we compare the overall charges on each side of the equation, we find a charge of +1 on the left but a charge of +3 on the right. This equation is not properly balanced. To balance it, let us write the two half reactions. Silver ions are reduced, and it takes one electron to change Ag+ to Ag:

Reduction half-reaction: \<\ce{Ag^{+} + e^{−} → Ag} \nonumber\>

Aluminum is oxidized, losing three electrons to change from Al to Al3+:

Oxidation half-reaction: \<\ce{Al → Al^{3+} + 3e^{−}} \nonumber\>

To combine these two half reactions and cancel out all the electrons, we need to multiply the silver reduction reaction by 3:

Exercise $$\PageIndex{1}$$

Write and balance the redox reaction that has calcium ions and potassium metal as reactants and calcium metal and potassium ions as products. Identify the substance oxidized, substance reduced, reducing agent and reducing agent.

Reduction: Ca2+ + 2e− → Ca

Oxidation: 2 (K → K+ + e−)

Combined: Ca2+ + 2K → Ca + 2K+

The substance oxidized is the reactant that had undergone oxidation: K The substance reduced is the reactant that had undergone reduction: Ca2+ The reducing agent is the same as the substance oxidized: K The oxidizing agent is the same as the substance reduced: Ca2+

Potassium has been used as a reducing agent to obtain various metals in their elemental form.

To Your Health: Redox Reactions and Pacemaker Batteries

All batteries use redox reactions to supply electricity because electricity is basically a stream of electrons being transferred from one substance to another. Pacemakers—surgically implanted devices for regulating a person’s heartbeat—are powered by tiny batteries, so the proper operation of a pacemaker depends on a redox reaction.

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Pacemakers used to be powered by NiCad batteries, in which nickel and cadmium (hence the name of the battery) react with water according to this redox reaction:

\<\ce{Cd(s) + 2NiOOH(s) + 2H2O(ℓ) → Cd(OH)2(s) + 2Ni(OH) 2(s)} \nonumber\>

The cadmium is oxidized, while the nickel atoms in NiOOH are reduced. Except for the water, all the substances in this reaction are solids, allowing NiCad batteries to be recharged hundreds of times before they stop operating. Unfortunately, NiCad batteries are fairly heavy batteries to be carrying around in a pacemaker. Today, the lighter lithium/iodine battery is used instead. The iodine is dissolved in a solid polymer support, and the overall redox reaction is as follows:

\<\ce{2Li(s) + I2(s) → 2LiI (s)} \nonumber\>

Lithium is oxidized, and iodine is reduced. Although the lithium/iodine battery cannot be recharged, one of its advantages is that it lasts up to 10 years. Thus, a person with a pacemaker does not have to worry about periodic recharging; about once per decade a person requires minor surgery to replace the pacemaker/battery unit. Lithium/iodine batteries are also used to power calculators and watches.

Figure $$\PageIndex{1}$$: A small button battery like this is used to power a watch, pacemaker, or calculator. (CC BY-SA; Gerhard H Wrodnigg via Wikipedia)