Precipitation Reactions (VCE SSCE Chemistry): Revision Notes

Precipitation Reactions

Introduction

A precipitation reaction happens when dissolved ions in a solution join together to create a compound that cannot dissolve in water. The solid compound that forms and settles out of solution is known as a precipitate.

Precipitation reactions have important practical applications, including:

• Removing minerals from drinking water
• Eliminating heavy metals from wastewater
• Purifying water in reservoir treatment plants

Ions in aqueous solution

When ionic compounds dissolve in water, they form aqueous solutions where the ions separate from each other and move freely. This process can be represented by equations showing the dissolving process.

For example, when sodium chloride dissolves:

$\text{NaCl(s)} \xrightarrow{\text{H}_2\text{O}} \text{Na}^+\text{(aq)} + \text{Cl}^-\text{(aq)}$

When copper(II) sulfate dissolves:

$\text{CuSO}_4\text{(s)} \xrightarrow{\text{H}_2\text{O}} \text{Cu}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)}$

If you mixed solutions of sodium chloride and copper(II) sulfate together, the mixture would contain separate $\text{Na}^+$, $\text{Cl}^-$, $\text{Cu}^{2+}$, and $\text{SO}_4^{2-}$ ions all moving about independently without combining with each other.

Reactions between ionic compounds in solution

In some cases, mixing two ionic solutions produces a visible reaction. For example, when a colourless solution of silver nitrate ($\text{AgNO}_3$) is mixed with a colourless solution of sodium chloride ($\text{NaCl}$), the solution turns cloudy because a fine white solid has formed.

To understand what happens in this reaction, you need to identify the ions present:

In silver nitrate solution:

• Silver ions ($\text{Ag}^+$)
• Nitrate ions ($\text{NO}_3^-$)

In sodium chloride solution:

• Sodium ions ($\text{Na}^+$)
• Chloride ions ($\text{Cl}^-$)

When the solutions are mixed, all four types of ions are present in the mixture. As the ions move around in solution, they collide with each other. When positive and negative ions collide, they may join together to form a new, insoluble precipitate.

Two new combinations of positive and negative ions are possible:

• Sodium and nitrate ions (forming sodium nitrate)
• Silver and chloride ions (forming silver chloride)

To work out which compound forms the precipitate, you need to check the solubility tables.

Solubility rules

Solubility tables help you predict whether mixing two ionic solutions will produce a precipitate. The tables classify compounds as soluble, insoluble, or slightly soluble based on how much dissolves in water at 25°C.

Soluble compounds in water

Compound type	Soluble (>0.1 mol/L at 25°C)	Insoluble exceptions (<0.01 mol/L at 25°C)	Slightly soluble exceptions (0.01-0.1 mol/L at 25°C)
Chlorides ($\text{Cl}^-$), bromides ($\text{Br}^-$), iodides ($\text{I}^-$)	Most chlorides, bromides, and iodides	$\text{AgCl}$, $\text{AgBr}$, $\text{AgI}$, $\text{PbI}_2$	$\text{PbCl}_2$, $\text{PbBr}_2$
Nitrates ($\text{NO}_3^-$)	All nitrates	No exceptions	No exceptions
Ammonium salts ($\text{NH}_4^+$)	All ammonium salts	No exceptions	No exceptions
Sodium ($\text{Na}^+$) and potassium ($\text{K}^+$) salts	All sodium and potassium salts	No exceptions	No exceptions
Ethanoates ($\text{CH}_3\text{COO}^-$)	All ethanoates	No exceptions	No exceptions
Sulfates ($\text{SO}_4^{2-}$)	Most sulfates	$\text{SrSO}_4$, $\text{BaSO}_4$, $\text{PbSO}_4$	$\text{CaSO}_4$, $\text{Ag}_2\text{SO}_4$

Insoluble compounds in water

Compound type	Insoluble (<0.01 mol/L at 25°C)	Soluble exceptions	Slightly soluble exceptions
Hydroxides ($\text{OH}^-$)	Most hydroxides	$\text{NaOH}$, $\text{KOH}$, $\text{Ba(OH)}_2$	$\text{Ca(OH)}_2$, $\text{Sr(OH)}_2$
Carbonates ($\text{CO}_3^{2-}$)	Most carbonates	$\text{Na}_2\text{CO}_3$, $\text{K}_2\text{CO}_3$, $(\text{NH}_4)_2\text{CO}_3$	No exceptions
Phosphates ($\text{PO}_4^{3-}$)	Most phosphates	$\text{Na}_3\text{PO}_4$, $\text{K}_3\text{PO}_4$, $(\text{NH}_4)_3\text{PO}_4$	No exceptions
Sulfides ($\text{S}^{2-}$)	Most sulfides	$\text{Na}_2\text{S}$, $\text{K}_2\text{S}$, $(\text{NH}_4)_2\text{S}$	No exceptions

Therefore, in the reaction between silver nitrate and sodium chloride, the precipitate must be silver chloride ($\text{AgCl}$).

How precipitation occurs

When the hydrated $\text{Ag}^+$ and $\text{Cl}^-$ ions come into contact in solution, the strong attraction between them causes an ionic lattice of $\text{AgCl}$ to form. The sodium and nitrate ions remain dissolved in solution.

This diagram shows the complete process:

1. Initially, the two solutions contain their respective ions surrounded by water molecules
2. When mixed, the ions move freely in the combined solution
3. The strong attraction between $\text{Ag}^+$ and $\text{Cl}^-$ ions causes them to form an ionic lattice
4. The $\text{AgCl}$ precipitate settles to the bottom
5. The $\text{Na}^+$ and $\text{NO}_3^-$ ions remain soluble and don't participate in the reaction

Predicting which compound precipitates

There's a simple visual method to work out which compound will form the precipitate:

1. Write down the positive ion from the first compound, followed by its negative ion
2. Write down the positive ion from the second compound, followed by its negative ion
3. Draw two lines: one connecting the positive ion of the first compound to the negative ion of the second compound, and another connecting the negative ion of the first compound to the positive ion of the second compound
4. Use solubility tables to determine which of the two new combinations will be insoluble

For example, with silver nitrate and sodium chloride:

• $\text{Ag}^+$, $\text{NO}_3^-$, $\text{Na}^+$, $\text{Cl}^-$
• New combinations: $\text{Ag}^+ + \text{Cl}^-$ and $\text{Na}^+ + \text{NO}_3^-$
• From solubility tables: AgCl is insoluble (precipitate), $\text{NaNO}_3$ is soluble

Writing full equations for precipitation reactions

Once you've identified the precipitate, you can write a chemical equation to represent the complete reaction.

The reaction between silver nitrate and sodium chloride can be written in words as:

silver nitrate solution + sodium chloride solution → silver chloride solid + sodium nitrate solution

Using formulas and state symbols, this becomes a full equation:

$\text{AgNO}_3\text{(aq)} + \text{NaCl(aq)} \rightarrow \text{AgCl(s)} + \text{NaNO}_3\text{(aq)}$

Spectator ions

Although sodium nitrate ($\text{NaNO}_3$) is written as a product in the equation, the sodium and nitrate ions are not actually combined with each other. They move freely through the solution and are present as separate ions at both the start and end of the reaction.

Writing ionic equations

An ionic equation shows only the species that directly participate in forming the precipitate. Spectator ions are not included because they don't change during the reaction.

You can think of an ionic equation as a simplified version of the full equation with the spectator ions removed. For the reaction between silver nitrate and sodium chloride:

Full equation: $\text{AgNO}_3\text{(aq)} + \text{NaCl(aq)} \rightarrow \text{AgCl(s)} + \text{NaNO}_3\text{(aq)}$

Ionic equation: $\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}$

The ionic equation focuses on the essential chemistry – the combination of silver and chloride ions to form the precipitate.

The chemistry of colour

Precipitation reactions have played an important role in the history of art. Many traditional paint pigments were created using precipitation reactions, with the coloured precipitate collected and used as the pigment.

Historical pigments from precipitation reactions

Ancient civilizations ground natural minerals to make pigments, but in the nineteenth century, chemists learned to manufacture pigments more reliably through precipitation reactions. This created a much wider range of colours for artists to use.

Pigment	Chemical formula
Cadmium yellow	$\text{CdS}$
Chrome yellow	$\text{PbCrO}_4$
Mosaic gold	$\text{SnS}_2$
Titanium white	$\text{TiO}_2$
Vermilion	$\text{HgS}$
Zinc yellow	$\text{ZnCrO}_4$

The availability of pigments based on ionic precipitates provided artists with a much wider palette. This helped enable the more vibrant use of colour seen in the works of the French Impressionists. For example:

• Cadmium sulfide was used to make red pigments, and cadmium red is still widely used today
• Viridian, a deep green pigment, is manufactured from chromium(III) oxide dihydrate