Let's do a (theoretical) experiment. We pour 250 mL of a colorless, silver nitrate solution (AgNO₃) into a beaker. We then submerge a copper wire into the beaker. After waiting several minutes, we check back in on our solution, and it looks like this:

Fig.1 Copper wire in silver nitrate solution

So, what happened? Why is our copper wire suddenly silver? And why is our solution suddenly blue? Well, a single replacement reaction occurred. This means that silver crystals are forming on the wire (hence why it is silvery) and copper nitrate (Cu(NO₃)₂) is in solution, which is blue.

In this article, we will be learning all about both single and double replacement reactions, so we can better understand why reactions like this occur.

• This article is about Single and Double Replacement Reactions.
• First, we will define what a single replacement reaction is.
• Then, we will look at some examples of single-replacement reactions.
• Next, we will learn about activity series and see how they are used in relation to single replacement reactions.
• Thereafter, we will define double replacement reactions and look at some examples.
• Lastly, we will learn about two special types of double-replacement reaction: precipitation reactions and neutralization reactions.

Single Replacement Reaction Definition

Let's start by looking at the definition of a single replacement reaction.

This reaction is sometimes called single displacement, since one element is "pushing out" (displacing) another.

The general formula for this reaction is:

$$A + BC \rightarrow AC + B$$

Essentially, element A is swapping with element B, which forms the new compound AC. This leaves element B as a lone element.

It's important to remember that the lone element will always swap with a species of like charge. Basically, only cation-forming (positively charged) species can displace another cation-forming species. For example, Iron (Fe), which can have a charge of +2 or +3, cannot displace oxygen (O), which can have a charge of -2.

Also, the element that does the replacing is always more reactive than the species it is replacing. We will talk about this more in the section "Activity Series".

Single Replacement Reactions

Now that we've covered what the general reaction looks like, let's look at some specific examples.

Cation replacement reactions

$$Cu + AgNO_3 \rightarrow CuNO_3 + Ag$$

$$Zn + CuCl_2 \rightarrow ZnCl_2 + Cu$$

$$Zn + H_2S \rightarrow H_2 + ZnS$$

In the above reactions, the lone (neutral) element can form a cation, which will replace the cation in the compound.

The first and second reactions are actually the reactions we talked about in our introduction. These elements don't need to have the same charge, just have the same type of charge (i.e. positive or negative). In the third reaction, zinc (Zn) has a +2 charge in zinc sulfide (ZnS), while hydrogen (H) has a +1 charge when bonded to sulfur (S).

Anion replacement reactions

$$Br_2 + 2KI \rightarrow 2KBr + I_2$$

$$ F_2 + NaCl \rightarrow NaF + Cl_2$$

Anion replacement reactions are virtually the exact same as cation replacement reactions, except the negatively charged species is the one being replaced.

When an element is by itself (with no charged noted) it is neutral. It only has a charged when bonded. Because of this, you need to remember which species are cation forming and which are anion forming. For example, the following reaction isn't possible:

$$Zn + HCl \rightarrow H_2Zn$$

This is because zinc only forms a cation (+2 charge), so it can't replace chlorine (Cl), which has a -1 charge when bonded to hydrogen.

Activity Series

Single-replacement reactions can only happen when the element replacing another is more reactive. To help us remember what species is more reactive, we use an activity series.

First up, here is the activity series for metals.

Fig.2- Activity Series for metals

From this chart, we can see some general patterns:

• Group 1 and 2 metals are more reactive than transition metals

• The reactivity of transition metals decreases as you go down the rows

Let's work on an example:

And now for the non-metals:

Fig.3-Activity Series for nonmetals

Here are some general rules:

• Group 17 is generally more reactive than group 16 (except for iodine)
• Moving down group 17, the elements get less reactive
• Moving down group 16, the elements get less reactive

While these are a good rule of thumb, it's best to keep an activity series handy!

Let's do another example, this time using the nonmetal series:

Double Replacement Reaction Definition

Now let's move on to double replacement reactions. Its definition is as follows:

The general formula for double replacement reactions is:

$$AB + CD \rightarrow CB + AD$$

Where A and C are both cations and B and D are both anions.

Double Replacement Reactions

In a double replacement reaction, either both cations are swapped or both anions. Unlike with single replacement reactions, you can't really tell which ones have been swapped, because you would end up with the same product regardless.

Below are some examples of double replacement reactions:

$$(NH_4)_2CO_3 + Ba(NO_3)_2 + 2NH_4NO_3 + BaCO_3$$

$$2KI + Pb(NO_3)_2 \rightarrow PbI_2 + 2KNO_3$$

$$Ca(OH)_2 + HCl \rightarrow H_2O + CaCl_2$$

Precipitation Reactions

Below are some examples of a precipitation reaction:

$$AgNO_{3\,(aq)} + NaCl_{(aq)} \rightarrow NaNO_{3\,(aq)} + AgCl_{(s)}$$

$$CuSO_{4\,(aq)} + 2NaOH_{(aq)} \rightarrow Na_2SO_{4\,(aq)} + Cu(OH)_{2\,(s)}$$

Neutralization Reactions

Another special type of double replacement reaction is the neutralization reaction.

The general formula for this type of reaction is:

$$HA + BOH \rightarrow H_2O + BA$$

Where A is the anion in the acid and B is the cation in the base.

The reaction is called a neutralization reaction because both salts (here shown as BA) and water (H₂O) have a neutral pH (pH=7)

Below are some examples:

$$HCl + NaOH \rightarrow H_2O + NaCl$$

$$HBr + KOH \rightarrow H_2O + KBr$$