What is Precipitate?
The definition of precipitate is a solid that precipitates (comes out of) solution. In chemistry, the solid usually forms due to a precipitation reaction taking place. Solid can also form due to a change in temperature or any other environmental change that affects the solubility of the compound. The solid compound may remain suspended in solution or fall to the bottom of the container.
Precipitate can also be used as a verb in chemistry. To precipitate is the act of a compound going from being aqueous in a solution to forming a solid product. These can also be called precipitation reactions.
What is a Precipitation Reaction?
The definition of a precipitation reaction is when two (or more) soluble salts react to form an insoluble product. The reactants are ions in the solution. The insoluble product is referred to as precipitate. A salt is an ionic compound.
Often, a precipitation reaction is a ‘double replacement reaction’. That is when there are two salts that are soluble, and when the cation of one binds with the anion of the other it forms a solid product that is not soluble. The other pair of cation and anion may or may not be soluble. This type of reaction takes the following form.
AB(aq) + CD(aq) –> AD(s) + BC(s or aq)
To be a precipitate reaction, either AD or CB will be an insoluble solid. AB and CD are both ionic compounds aqueous in solution.
In precipitation reactions, the formed precipitate can remain suspended in solution or may sink to the bottom. The solid particles can then also be removed from the solution by various means such as filtration, decantation, centrifuging. The liquid left behind is referred to as the supernatant.
These reactions are commonly used to help determine what ions are in the solution.
How to Identify a Precipitation Reaction
A precipitation reaction will always have a solid product. The reactants are usually two or more ionic aqueous molecules. The product must include a solid product.
The most general form of a precipitation reaction then is:
A+(aq) + B–(aq) –> AB(s)
Ag+(aq) + Cl–(aq) –> AgCl(s)
The reactants must be ionic compounds in solution. So the reaction including all components would be:
AC(aq) + BD(aq) –> AD(s) + B+(aq) + C–(aq)
AgNO3(aq) + NaCl (aq) –> AgCl(s) + Na+(aq) + NO3–(aq)
Because the reactants are aqueous and we want to know the ions in solution, it is common to write the reaction in terms of the ions. This format is called the complete ionic equation:
Ag+(aq) + NO3–(aq) + Na+(aq) + Cl–(aq) –> AgCl(s) + Na+(aq) + NO3–(aq)
To simplify this equation further, get rid of any ion that appears on both the reactant and product side as that indicates they are not part of the reaction. These ions are also called spectator ions. What is left is called the net ionic equation.
Net ionic equations must be balanced to be accurate. Both the charge and the number of atoms of each element must be balanced.
What are solubility rules?
These are general guidelines or rules on what compounds will form a precipitate. A great resource is finding a good solubility table or solubility chart. There are also some general rules you can learn for the solubility of different compounds.
- Alkali metals (Group I) are soluble.
- Nitrates, acetates, chlorates, and perchlorates are generally soluble.
- Chlorides, bromides, and iodides are soluble except when with Ag+, Pb2+, and Hg22+.
- Most silver salts are insoluble. The two main exceptions are silver acetate and silver nitrate.
- Sulfates are soluble, except with Ca2+, Sr2+, Ba2+, Pb2+, and Ag2+.
- Hydroxides are insoluble except for with alkali metals. Hydroxides are slightly soluble with the alkaline earth metals.
- Sulfides are insoluble except with the alkali and alkaline earth metals.
- Carbonates, chromates, phosphates, and fluorides are all insoluble except for with alkali metals and ammonium.
A full solubility chart and description of what compounds are soluble can be found on the Solubility Rules tutorial.
To learn how to name these compounds, read the Naming Ionic Compounds tutorial!
Uses of Precipitation Reactions and Real-Life Examples
Precipitation reactions are commonly used to identify if certain ions are present in a solution. For example, to determine if lead (Pb2+) is present in the solution, a solution containing chlorides or hydroxides could be added. The lead would precipitate out as either PbCl2 or Pb(OH)2 and indicate that lead is present. Additionally, these tests are also commonly used in chemistry labs. For those tests, a series of compounds can be added to deduce what ions are present.
Nature also makes some cool precipitate structures. Near deep-sea hydrothermal vents, many minerals precipitate, particularly sulfides, leaving behind huge chimneys on the ocean floor.
Precipitation Reaction Example Problems
Predict the precipitate(s) in the following precipitation reaction examples.
- AgF(aq) + CaCl2(aq)
- AgClO3(aq) + CaI2 (aq)
- LiNO3(aq) + Na2CO3(aq)
- BaCl2(aq) + K2SO3(aq)
- ZnF2 and MgSO3
The possible products are listed below. By looking at our solubility rules or a solubility table we can check which ones will form a precipitate.
- AgCl is the precipitate. The other possible product is CaF2 which is soluble.
- AgI is the precipitate. The other possible product is Ca(ClO3)2 which is soluble.
- No precipitate. The two possible products are LiCO3 and NaNO3 neither of which is insoluble. Both are soluble. Therefore this is NOT actually a precipitation reaction.
- BaSO3 is the precipitate. The other possible product is KCl which is soluble.
- MgF2 and ZnSO3 are both precipitates in this reaction.
What would you add to a solution to determine if Mg2+ is present?
To easily determine if Mg2+ is present we want to add an ion that will precipitate when it binds with Mg2+. Knowing our solubility rules we can see that F–, OH–, CO32-, and PO43- would all cause precipitates. We can’t directly add these ions and instead need to combine them with a cation they are soluble with to make a salt. There are many options for this requirement. Some possibilities are NaF, NH4F, Ba(OH)2, LiOH, Li2CO3, K3PO4, and Na3PO4.