3.5: Precipitation Reactions

Learning Objectives 

• Predict the products of two or mor reactant when various aqueous solutions are mixed 
• Predict the states of the products and write balanced molecular equations
• Write the complete ionic equations for reactions 
• Write the net ionic equations for molecular compounds 
• Differentiate between molecular equations, complete ionic equations, and net ionic equations

 

Introduction 

 

Now that we have learned the solubility rules, we can predict what will happen for single and double replacement reactions that occur in aqueous solutions. Note, if in a double displacement reaction two aqeuous solutions combine and form a solid, you have a precipitation reaction. Many texts treat this as a type of reaction, but we will treat it as a subset of the double displacement reactions.

 

Writing Molecular Equations: Predicting Products for Double and Single Displacement Reactions

 

Determine what the products will be for the following reactions. Initially we will ignore phases, but by the time we finish this exercise you will need to be able to identify phases as you write the equation. This process requires that you identify the ions in the reactants and for double displacement reactions swap partners and use the principle of charge neutrality to determine what the product are. For single replacement reaction you need to figure which species (anion or cation) is gaining or losing charge and the other species is a spectator ion. Once that is done you need to balance the equations

Students frequently swap ions without thinking about the product formula and this leads to mistakes. It is suggested you follow these steps:

1. Identify reactant ions and their charge
2. Swap Ions
3. Determine Product Formula based on principle of charge neutrality
4. Balance Equation
5. Predict the state of the products. 

Example:

\[Pb\left ( ClO_{4} \right )_{2}\left ( aq \right )+Na_{2}SO_{4}\left ( aq \right )\rightarrow\]

1. Identify reactant ions and their charge

   Pb⁺² ClO₄^-

   Na⁺ SO₄^-2

2. Swap Ions
   (Forms lead(II) sulfate and sodium perchlorate)
3. Determine Product Formula based on principle of charge neutrality
   \[Pb\left ( ClO_{4} \right )_{2}\left ( aq \right )+Na_{2}SO_{4}\left ( aq \right )\rightarrow PbSO_{4}\left ( ? \right )+NaClO_{4}\left ( ? \right )\]
   If you need help determining formulas review section 2.6
4. Balance Equation

\[Pb\left ( ClO_{4} \right )_{2}\left ( aq \right )+Na_{2}SO_{4}\left ( aq \right )\rightarrow PbSO_{4}\left ( ? \right )+2NaClO_{4}\left ( ? \right )\]

 

5. Predict the states of the products - Remember, (aq) = aqueous, (s) = solid, (g) = gas and (l) = liquid.

Pb(ClO₄)₂(aq) + Na₂SO₄(aq) → PbSO₄(s) + 2 NaClO₄(aq) 

 

You have just written a balanced complete (molecular) equation, an equation that shows all the products and reactants with their states and is balanced. Let's see another example in the following video. 

 

Video Tutor: Writing Molecular Equations

Al₂(SO₄)₃ + BaCl₂ ==> ?

Part 1 Practice Worksheet - write a balanced molecular equation for the following reactants, being sure to identify the states of the products and balance the reaction. 

Part 1 Key- Check your work. 

 

Complete Ionic Equation:

Molecular equations identify the compounds in a reaction,  but they are not entirely accurate. If you think about it, ions of soluble ionic compounds (aq) do not stay together, but ionize. For example, writing aqueous sodium perchlorate as NaClO₄ (aq) is not accurate because the compound is no longer intact, but dissociated into Na⁺(aq) and ClO₄^- (aq) ions. The complete ionic equation shows aqueous compounds as separate ions, and compounds that exist in the solid, liquid, or gas form as compound. For example, write NaCl(aq) as Na⁺(aq) and Cl^-(aq), but write AgCl(s) as AgCl(s).

 

Example:

\[Pb_{+2}\left ( aq \right )+2ClO_{4}^{-}\left ( aq \right )+2Na^{+}\left ( aq \right )+SO_{4}^{-2}\left ( aq \right )\rightarrow PbSO_{4}\left ( s \right )+2Na^{+}\left ( aq \right )+2ClO_{4}^{-}\left ( aq \right )\]

 

 

 

Part 2 Practice Worksheet - write the complete ionic equation 

Note: this is the same worksheet as in part 1. You should take your answers from the worksheet in part 1, and now rewrite them as complete ionic equations. This means you need to redo the reactants if they are soluble, and so you may want to write this on a separate sheet of paper.

Part 2 Key - Check you work .

Net Ionic Equation

If you think about it, if an ion is a reactant and a product, it does not do anything. It is simply floating around in solution and not contributing to the formation of a solid, liquid, or gas to form a reaction.  We can call it a spectator ion, and it can be ignored. The net ionic equation ignores the spectator ions, and shows what reactions are really happening, that is, what bonds are being broken, and what bonds are being made.

If a product is a soluble ionic compound, it is probably a spectator ion. At this point that works all the time, but in the next section when we introduce acids and some amine bases, which are not ionic compounds, we may have instances where that statement is not true.

Net Ionic Equation. Under each equation write the net ionic equation. You do this by cancelling out spectator ions from the total ionic equation. Note, the stoichimetric coefficent of the net ionic equation may be different than from the molecular or total ionic.

Example:

\[Pb^{+2}\left ( aq \right )+SO_{4}^{-2}\left ( aq \right )\rightarrow PbSO_{4}\left ( s \right )\]

 

Watch this video for one more example on how to writ complete ion and net ionic equations: 

Part 3 Practice Worksheet - write the net ionic equation 

Note: You should use practice worksheet 2 as a starting point to this exercise. The key is at the bottom of the page.

Part 3 Key

 

Mixtures of many solutions

Mixing Multiple Solutions

Consider 3 aqueous salt solutions each with different cations and anions. For example, consider NaCl, AgNO₃ and (NH₄)₂CO₃. There are 3 cations and 3 anions which results in 3² or 9 combinations (of which 3 are the reactants). This seems like a very complicated problem (4 salts would result in 16 potential combinations and 5 salts would result in 25) and so we need to develop a technique to see the problem. This can be done through a matrix, where the rows represent the cations, the columns the anions and the cells the potential combinations:

	Cl^-	NO3^-	CO3^-2
Na^+
Ag^+
NH4^+

Now think about it for a minute. By writing the reactants out this way the diagonal represents your reactants and anything in the first or third row must be soluble (Rules 1A), so you can strike them out. Likewise anything in the second column must be soluble, so you have reduced this problem to two questions. Are silver chloride and silver carbonate precipitates or do they form aqueous solutions? 

	Cl^-	NO3^-	CO3^-2
Na^+	xxxx	xxxx	xxxx
Ag^+	AgCl	xxxx	Ag2CO3
NH4^+	xxxx	xxxx	xxxx

Note, all sodium, nitrate and ammonium salts are soluble and so we can ignore them in identifying potential precipitates.

So the answer is AgCl(s) and Ag₂(CO₃) (s)

Video: Predicting Multiple Products

NaCl + Na₃PO₄ + AgNO₃ + Pb(NO₃)₂ + Na₂SO₄

When trying to predict if any (s) form from multiple solutions mixed together, draw a matrix and use the solubility rules to cancel out all the (aq).

 

Exercise \(\PageIndex{1}\)

Identify if any precipitates in the following solutions are mixed:

Pb(ClO₄)₂(aq), K₂SO₄(aq), AgCH₃CO₂ (aq), KCl(aq)

Answer

 

Exercise \(\PageIndex{2}\)

Identify if any precipitates in the following solutions are mixed:

Ba(ClO₄)₂(aq), Li₂SO₄(aq), NH₄CH₃CO₂ (aq), BaCl₂(aq) 

Answer

 

Contributors: 

• Robert Belford (UA of Little Rock) 
• Ronia Kattoum (UA of Little Rock)