Salt analysis involves chemical procedures that help identify an ion based on its precipitates (insoluble/hardly soluble salts with different solubility, color, texture etc.).
There are several aspects that differentiate ions, making them easier to identify, such as:
Some salts can be solubilized using excess reagent or other solutions, such as alkaline hydroxides or ammonia solution, often resulting in complex salts (coordination compounds), where the central atom is usually the ion to identify, bonded to surrounding molecules (\(\ce{NH3}\)) or ions (\(\ce{OH^-}\)) called ligands.
Changes of precipitate color can occur in redox reactions caused by the oxygen in the air
These precipitates have specific aspects for each ion and are formed when two soluble compounds react.
Below are presented particularities for common ions:
Silver Ion (\(\ce{Ag^+}\))
Ion
Reagent
Precipitate
Color, Aspect
Soluble in / Observations
\(\ce{Ag^+}\)
\(\ce{NaI}\), \(\ce{NaBr}\), \(\ce{NaCl}\) or \(\ce{HI}\), \(\ce{HBr}\), \(\ce{HCl}\)
\(\ce{AgI}\), \(\ce{AgBr}\), \(\ce{AgCl}\)
\(\ce{AgI}\) - Yellow \(\ce{AgBr}\) - Pale Yellow \(\ce{AgCl}\) - Milky White
\(\ce{AgCl}\) and \(\ce{AgBr^-}\) insoluble in excess reagent but soluble in \(\ce{NH3}\) and \(\ce{KCN}\) sol. => \(\ce{[Ag(NH3)2]^+}\) and \(\ce{[Ag(CN)2]^-}\), colorless; AgBr - insoluble in excess reagent but soluble in \(\ce{KCN}\) => \(\ce{[Ag(CN)2]^-}\), colorless
\(\ce{(NH4)2S}\) or \(\ce{H2S}\)
\(\ce{Ag2S}\)
Black
Insoluble in excess reagent but soluble in \(\ce{KCN}\) => \(\ce{[Ag(CN)2]^-}\), also in hot \(\ce{HNO3}\)
\(\ce{NaOH}\)
\(\ce{Ag2O}\) (hydroxide decomposes in \(\ce{Ag2O}\) and water)
Dark Brown
Soluble in excess reagent => \(\ce{[Ag(NH3)2]^+}\) colorless
\(\ce{NH3}\) sol.
\(\ce{Ag2O}\) (\(\ce{NH3}\) solution is basic, \(\ce{Ag2O}\) forms for the same reason as above)
Dark Brown
Soluble in excess reagent => \(\ce{[Ag(NH3)2]^+}\) colorless
\(\ce{(NH4)2CO3}\)
\(\ce{Ag2CO3}\)
Pale Yellow
Soluble in excess reagent => \(\ce{[Ag(NH3)2]^+}\) colorless
Aluminium Ion (\(\ce{Al^{3+}}\))
Ion
Reagent
Precipitate
Color, Aspect
Soluble in / Observations
\(\ce{Al^{3+}}\)
\(\ce{(NH4)2S}\)
\(\ce{Al(OH)3}\) (\(\ce{Al2S3}\) and \(\ce{Al2(CO3)3}\) hydrolyze, with the formation of \(\ce{H2S}\) and \(\ce{CO2}\) respectively)
White, Gelatinous
Insoluble in excess reagent
\(\ce{NaOH}\)
Soluble in excess reagent => \(\ce{[Al(OH)4]^-}\) colorless
\(\ce{NH3}\) sol.
partially soluble in excess reagent => \(\ce{[Al(OH)4]^-}\) colorless
\(\ce{Na2CO3}\)
Barium (\(\ce{Ba^{2+}}\)), Strontium (\(\ce{Sr^{2+}}\)) and Calcium (\(\ce{Ca^{2+}}\)) ions
These three ions have been grouped due to their similar properties.
Generally, they can be differentiated through their solubilities and the speed they form precipitates, with Barium being the most insoluble, followed by Strontium that forms slower even with similar concentrations and lastly Calcium, that forms partially soluble precipitates for almost every reagent except alkaline carbonates.
Insoluble in excess reagent but soluble in mineral acids => colorless sol.
\(\ce{(NH4)2SO4}\)
\(\ce{MSO4}\)
Crystaline White
Highly insoluble, except for \(\ce{CaSO4}\) - partially soluble
\(\ce{NaOH}\)
\(\ce{M(OH)2}\)
White
The hydroxides of these metals are partly soluble and usually form in small amounts, hardly noticeable at first.
Lead Ion (\(\ce{Pb^{2+}}\))
Ion
Reagent
Precipitate
Color, Aspect
Soluble in / Observations
\(\ce{Pb^{2+}}\)
\(\ce{NaOH}\)
\(\ce{Pb(OH)2}\)
White
Soluble in excess reagent => \(\ce{[Pb(OH)4]}\) colorless
\(\ce{KI}\)
\(\ce{PbI2}\)
Bright yellow, powdery, easy to recognize
Soluble in boiling water, recrystalizes into golden spangles after cooling ("golden rain")
\(\ce{NaCl}\), \(\ce{HCl}\) dil.
\(\ce{PbCl2}\)
White
Insoluble in excess reagent but soluble in hot water => colorless sol.
\(\ce{NH3}\)
\(\ce{Pb(OH)2}\)
Insoluble in excess reagent but soluble in \(\ce{NaOH}\) => \(\ce{[Pb(OH)4]^{2-}}\) colorless sol. and mineral acids (\(\ce{HNO3}\))
\(\ce{(NH4)2SO4}\)
\(\ce{PbSO4}\)
\(\ce{Na2CO3}\)
\(\ce{PbCO3}\)
Iron Ions: \(\ce{Fe^{2+}}\) (Ferrous) , \(\ce{Fe^{3+}}\) (Ferric)
Ion
Reagent
Precipitate
Color, Aspect
Soluble in / Observations
\(\ce{Fe^{2+}}\) \(\ce{Fe^{3+}}\)
\(\ce{(NH4)2S}\)
\(\ce{FeS}\) \(\ce{Fe2S3}\)
Black
Insoluble in excess reagent but soluble in \(\ce{HCl}\) (highly corrosive)
\(\ce{NaOH}\)
\(\ce{Fe(OH)2}\) \(\ce{->}\) \(\ce{Fe(OH)3}\)
Dark Green (algae-like) \(\ce{->}\) Red Brown
Insoluble in excess reagent or \(\ce{NH3}\) sol.
\(\ce{NH3}\) sol.
\(\ce{Na2CO3}\)
\(\ce{FeCO3}\) \(\ce{->}\) \(\ce{Fe(OH)3}\)
White \(\ce{->}\) Red Brown
Redox reaction occuring with a change in color
\(\ce{FeCO3 + 2 H2O -> Fe(OH)2 + H2O + CO2}\)
\(\ce{4 Fe(OH)2 + O2 + 2 H2O -> 4 Fe(OH)3}\)
Copper Ion (\(\ce{Cu^{2+}}\))
Ion
Reagent
Precipitate
Color, Aspect
Soluble in / Observations
\(\ce{Cu^{2+}}\)
\(\ce{H2S}\) \(\ce{(NH4)2S}\)
\(\ce{CuS}\)
Black
Insoluble in excess reagent or \(\ce{NH3}\) sol., soluble in mineral acids (\(\ce{HNO3}\))
\(\ce{KI}\)
\(\ce{CuI2 -> CuI + I2}\)
White ppt. in Brown sol.
Insoluble in excess reagent or \(\ce{NH3}\) sol.
\(\ce{NaOH}\)
\(\ce{Cu(OH)2}\)
Light Blue
Soluble in \(\ce{NH3}\) => \(\ce{[Cu(NH3)4]^{2+}}\), blue
\(\ce{NH3}\) sol.
Blue
The decomposition of \(\ce{CuI2}\) gives a characteristic brown coloration to the solution, indicating the presence of iodine
\(\ce{2 CuI2 -> 2 CuI + I2}\)
Experiment Overview
Salt analysis is a qualitative inorganic analysis used to identify the cations or anions present in a salt / mixture.
What this Experiment Demonstrates
How to identify the ions present in an inorganic salt using chemical tests
Apparatus Required
Pipettes
Measuring cylinder (optional)
Test tubes
Test tube stand
Wash bottle with distilled water
Method
Pipette a small amount (1-2 drops) of unknown ion solution into a clean tube
Pipette a similar amount of reagent into the same tube
For the reaction to occur efficiently, gently spray some distilled water on the tube's walls from the wash bottle
Write down any observations regarding color and texture of the salt formed; try to notice effervescence or characteristic ammonia gas smell
Carefully add excess reagent / (or solutions such as \(\ce{NH3}\), \(\ce{NaOH}\), etc) and observe the changes
Observations
Depending on the ion, color changes or certain aspects can be observed in the inorganic salt formed, with textures ranging from gelatinous to powdery, and colors ranging from white or dark colors to vibrant yellow or blue.
Calculations
Due to the method being of qualitative type and not quantitative, solution / compound measurements are generally not required.
Result
By knowing the certain properties of ions (color, aspect), they can be easily identified using the right methods
Video Demonstration
Video courtesy of ChemXpert (YouTube). Video used for educational purposes under YouTube’s embedding policy.
Practice
Imagine you're in the IJSO experimental exam and you recieve 5 solutions, numbered 1-5, that contain \(\ce{Pb(NO3)2}\), \(\ce{CuSO4}\), \(\ce{KI}\), \(\ce{NaOH}\), \(\ce{BaCl2}\). You perform all the cross reactions, and you get the following results:
1
2
3
4
5
2
And a white precipitate
X
X
X
X
3
No visible reaction
X
X
X
4
X
X
5
No visible reaction
X
Using the tables above, identify the substance contained in each solution.
You'll often find that white precipitates are quite annoying and not the most helpful.
Before starting to do the actual reactions, you should make a table, like the one above, with what you expect to see for each of the cross reactions.
For example, you know \(\ce{KI}\) will form a yellow precipitate (\(\ce{PbI2}\)) and a brown solution (\(\ce{I2}\)). Therefore you can identify 1 = \(\ce{KI}\), 2 = \(\ce{CuSO4}\), 4 = \(\ce{Pb(NO3)2}\).
Looking at the substances, \(\ce{CuSO4}\), in reaction with \(\ce{NaOH}\), will form blue \(\ce{Cu(OH)2}\), so 3 = \(\ce{NaOH}\). We’re left with 5 = \(\ce{BaCl2}\). You can check the other reactions and you'll see they match.
