Solubility Rules Revision Notes for HSC SSCE Chemistry

Solubility Rules

Introduction to aqueous solutions

In chemistry, solutions are homogeneous mixtures where one substance (the solute) is dissolved in another (the solvent). When we talk about aqueous solutions, we mean solutions where water is the solvent. The word 'aqueous' comes from the Latin word aqua, meaning water.

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Aqueous solutions are everywhere in nature. Oceans are aqueous solutions containing many dissolved compounds, including ionic substances. All the chemical reactions that keep living things alive happen in aqueous solutions within our cells and bodies.

What happens when ionic substances dissolve

When ionic substances dissolve in water, they undergo a process called dissociation. This means the ionic compound separates into its individual ions, which then move freely and independently through the solution.

For example, when sodium chloride (common table salt) dissolves in water:

$\text{NaCl(s)} \rightarrow \text{Na}^+(aq) + \text{Cl}^-(aq)$

The solid sodium chloride breaks apart into sodium ions ( $\text{Na}^+$ ) and chloride ions ( $\text{Cl}^-$ ), both of which are surrounded by water molecules and can move throughout the solution.

Understanding solubility categories

Not all ionic compounds dissolve to the same extent in water. Chemists classify the solubility of compounds into three categories:

Soluble: The compound dissolves to more than $10 \text{ g L}^{-1}$ (or $1 \text{ g}/100 \text{ mL}$ )
Insoluble: The compound dissolves to less than $1 \text{ g L}^{-1}$
Sparingly soluble: The compound dissolves in the range $1 \text{ g L}^{-1}$ to $10 \text{ g L}^{-1}$

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Exam tip: Units can be written as either $\text{g L}^{-1}$ or $\text{g/L}$ - both mean exactly the same thing!

Precipitation reactions

A precipitation reaction occurs when solutions of certain ionic compounds are mixed together and produce a solid substance called a precipitate.

How precipitation happens

Let's consider what happens when sodium chloride solution is mixed with silver nitrate solution. Initially, we have:

$\text{Na}^+$ and $\text{Cl}^-$ ions from the sodium chloride solution
$\text{Ag}^+$ and $\text{NO}_3^-$ ions from the silver nitrate solution

When these solutions are mixed, all four types of ions are moving randomly throughout the combined solution. Each ion is surrounded by water molecules. However, the $\text{Ag}^+$ and $\text{Cl}^-$ ions are more strongly attracted to each other than they are to the water molecules. This causes them to cluster together and form solid silver chloride, which precipitates out of solution.

Spectator ions

In the precipitation reaction above, the $\text{Na}^+$ and $\text{NO}_3^-$ ions don't participate in the reaction. They remain dissolved in the solution before, during, and after the precipitation occurs. We call these non-participating ions spectator ions - they just 'watch' the reaction happen without getting involved.

Writing equations for precipitation reactions

Precipitation reactions can be represented using three different types of equations:

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1. Overall equation

This shows all reactants and products as complete, neutral compounds:

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

2. Complete ionic equation

This shows all the ions present in the solution:

$\text{Na}^+\text{(aq)} + \text{Cl}^-\text{(aq)} + \text{Ag}^+\text{(aq)} + \text{NO}_3^-\text{(aq)} \rightarrow \text{AgCl(s)} + \text{Na}^+\text{(aq)} + \text{NO}_3^-\text{(aq)}$

3. Net ionic equation

This shows only the reacting ions (spectator ions removed):

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

chatImportant

Exam tip: The net ionic equation most accurately represents what's actually happening in the precipitation reaction, so it's the most commonly used form.

Predicting precipitation using solubility rules

Chemists have systematically mixed many different solutions to determine experimentally which compounds are soluble and which form precipitates. These results have been organised into solubility tables that we can use to predict whether a precipitation reaction will occur when two solutions are mixed.

Table of solubility rules

The following table summarises the solubility of common ionic compounds:

Soluble Anions	Exceptions	Insoluble Anions	Exceptions
$\text{NO}_3^-$	None	$\text{OH}^-$	Group 1, $\text{NH}_4^+$ , $\text{Ba}^{2+}$ , $\text{Sr}^{2+}$ soluble; $\text{Ca}^{2+}$ slightly soluble
$\text{CH}_3\text{COO}^-$	$\text{Ag}^+$ slightly soluble	$\text{O}^{2-}$	Group 1, $\text{NH}_4^+$ , $\text{Ba}^{2+}$ , $\text{Sr}^{2+}$ , $\text{Ca}^{2+}$ soluble
$\text{Cl}^-$	$\text{Ag}^+$ insoluble; $\text{Pb}^{2+}$ slightly soluble	$\text{S}^{2-}$	Groups 1 and 2, $\text{NH}_4^+$ soluble
$\text{Br}^-$	$\text{Ag}^+$ insoluble; $\text{Pb}^{2+}$ slightly soluble	$\text{CO}_3^{2-}$	Group 1, $\text{NH}_4^+$ soluble
$\text{I}^-$	$\text{Ag}^+$ , $\text{Pb}^{2+}$ insoluble	$\text{SO}_3^{2-}$	Group 1, $\text{NH}_4^+$ soluble
$\text{SO}_4^{2-}$	$\text{Ba}^{2+}$ , $\text{Pb}^{2+}$ , $\text{Sr}^{2+}$ insoluble; $\text{Ag}^+$ , $\text{Ca}^{2+}$ slightly soluble	$\text{PO}_4^{3-}$	Group 1, $\text{NH}_4^+$ soluble

chatImportant

Key generalisation: All Group 1 and $\text{NH}_4^+$ salts are soluble.

This table applies to common cations including: Groups 1 and 2, $\text{NH}_4^+$ , $\text{Al}^{3+}$ , $\text{Cr}^{3+}$ , $\text{Mn}^{2+}$ , $\text{Fe}^{2+}$ , $\text{Fe}^{3+}$ , $\text{Cu}^{2+}$ , $\text{Co}^{2+}$ , $\text{Ni}^{2+}$ , $\text{Ag}^+$ , $\text{Zn}^{2+}$ , $\text{Sn}^{2+}$ , $\text{Cd}^{2+}$ and $\text{Pb}^{2+}$ .

How to use the solubility table

To predict whether a precipitate will form when two solutions are mixed:

Write the formulae of the compounds present
Determine what products would form by swapping the partners of the reacting compounds
Check the formulae are correct
Use the solubility table to identify whether any products are insoluble
If an insoluble product forms, that's your precipitate!

Applying solubility data

Precipitation reactions have important practical applications in chemistry. Chemists often need to identify substances such as cations and anions that may have contaminated food, soil, or water supplies.

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For example, precipitation reactions can be used to test for dangerous metals like lead or cadmium in drinking water. They can also be used to remove unwanted substances from water through selective precipitation.

Worked example: Predicting precipitation

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Worked Example: Predicting Precipitation

Question 1: What precipitate, if any, will form when aqueous solutions of magnesium iodide and silver nitrate are mixed? If a precipitate forms, write the neutral species and net ionic equations for the reaction.

Solution:

Step 1: Write the formulae of the compounds present

Magnesium iodide: $\text{MgI}_2$
Silver nitrate: $\text{AgNO}_3$

Step 2: Predict the products by swapping partners

$\text{MgI}_2 + \text{AgNO}_3 \rightarrow \text{AgI} + \text{Mg(NO}_3)_2$

Step 3: Use the solubility table to identify if either product is insoluble

All nitrates are soluble, so $\text{Mg(NO}_3)_2$ is soluble
All iodides are soluble except $\text{Ag}^+$ and $\text{Pb}^{2+}$ , so $\text{AgI}$ will precipitate

Answer: The precipitate is $\text{AgI}$ .

Neutral species equation: $\text{MgI}_2\text{(aq)} + 2\text{AgNO}_3\text{(aq)} \rightarrow 2\text{AgI(s)} + \text{Mg(NO}_3)_2\text{(aq)}$

Net ionic equation: $\text{Ag}^+\text{(aq)} + \text{I}^-\text{(aq)} \rightarrow \text{AgI(s)}$

Worked example: Selecting a reagent

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Worked Example: Selecting a Reagent

Question 2: Select a reagent that could be used to precipitate the cation in $\text{BaBr}_2$ , explaining your selection.

Solution:

Step 1: Identify the ions present in solution

$\text{Ba}^{2+}$ and $\text{Br}^-$ are the ions present in $\text{BaBr}_2$ solution

Step 2: Identify which ion needs to be precipitated

$\text{Ba}^{2+}$ is the cation to be precipitated

Step 3: Use the solubility table to find suitable anions

Looking at insoluble anions that form precipitates with $\text{Ba}^{2+}$ :

$\text{SO}_4^{2-}$ (barium sulfate is insoluble)
$\text{S}^{2-}$ (barium sulfide is insoluble)
$\text{CO}_3^{2-}$ (barium carbonate is insoluble)
$\text{SO}_3^{2-}$ (barium sulfite is insoluble)
$\text{PO}_4^{3-}$ (barium phosphate is insoluble)

Step 4: Choose an appropriate anion and cation

Let's select $\text{SO}_4^{2-}$ as the anion to precipitate $\text{Ba}^{2+}$
We need a cation that won't precipitate with $\text{Br}^-$
$\text{Na}^+$ is a good choice because all Group 1 salts are soluble

Answer: $\text{Na}_2\text{SO}_4$ can be used to precipitate $\text{Ba}^{2+}$ from a solution of $\text{BaBr}_2$ .

chatImportant

Exam tip: When choosing a reagent, it's usually best to select a cation from Group 1 or $\text{NH}_4^+$ because compounds containing these cations are always soluble. Avoid choosing $\text{Ag}^+$ or $\text{Pb}^{2+}$ as they form precipitates with many anions.

Remember!

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Key Points to Remember:

Aqueous solutions are substances dissolved in water. When ionic compounds dissolve, they dissociate into freely moving ions.
Precipitation reactions occur when two ionic solutions are mixed and form an insoluble solid (precipitate).
Three types of equations: Overall equations show complete compounds, complete ionic equations show all ions present, and net ionic equations show only the reacting ions (most useful for precipitation reactions).
Spectator ions are ions that don't participate in the reaction - they remain dissolved in solution throughout.
Solubility rules allow us to predict whether a precipitate will form when two solutions are mixed. All Group 1 and $\text{NH}_4^+$ salts are soluble!

Solubility Rules (HSC SSCE Chemistry): Revision Notes