The basic testing procedure for identifying a salt is as follows.
- Appearance of compound
The compound will most likely be in solid form. Note the color and
shape of the crystals. Ionic compounds formed from the representative
elements tend to be white or colorless, while ions of transition
elements tend to be colored. The following is a table of the colors of
metal ions in solution with NO3-.
Ion Color Co2+ rose Co3+ violet Cr3+ violet Cu2+ blue Fe2+ pale green, pale violet Fe3+ yellow-brown - Heating effect
Heating a compound can cause a liquid to condense on the inside of
the test tube. This is probably water, indicating that the compound is
a hydrate. If a gas is given off, note the color and odor of the gas.
The nitrate, carbonate, and sulfite ions may decompose, as illustrated
by the reactions:
2 Pb(NO3)2(s) + heat → 2 PbO(s) + O2(g) + 4 NO2(g, brown)Some bromides and iodides decompose to give Br2(g, orange-brown) and I2(g, purple).
CaCO3(s) + heat → CaO(s) + CO2(g, colorless, odorless)
CaSO3(s) + heat → CaO(s) + SO2(g, colorless, pungent)
- Flame test
Solutions of ions, when mixed with concentrated HCl and heated on a
nickel/chromium wire in a flame, cause the flame to change to a color
characteristic of the atom.Visible colors occur with the following ions:
Sodium Bright yellow (intense, peristent) Potassium Pale violet (slight, fleeting) Calcium Brick red (medium, fleeting) Strontium Crimson (medium) Barium Light green (slight) Lead Pale bluish (slight, fleeting) Copper Green or blue (medium, persistent) - Solubility in water
Place one small spatula of the compound in 1 mL of water. If the
compound is soluble this amount will dissolve after considerable
stirring. If the compound is moderately soluble, some of this amount
will dissolve. If the compound is insoluble, even a very small amount
will not dissolve.
General solubility rules:
- All nitrates are soluble.
- Practically all sodium, potassium, and ammonium salts are soluble.
- All chlorides, bromides, and iodides are soluble except those of silver, mercury(I), and lead(II).
- All sulfates are soluble except those of strontium, barium, and lead(II), which are insoluble, and those of calcium and silver which are moderately soluble.
- All carbonates, sulfites, and phosphates are insoluble except those of sodium, potassium, and ammonium.
- All sulfides are insoluble except those of the alkali metals, the alkaline earth metals, and ammonium.
- All hydroxides are insoluble except those of the alkali metals. The hydroxides of calcium, strontium, and barium are moderately soluble. Ammonium hydroxide does not exist; ammonium hydroxide is a misnomer for aqueous ammonia, NH3(aq).
- Reaction with nitric acid
Add nitric acid to the compound and observe any reaction that
occurs. If the compound dissolved in water, it should dissolve in
nitric acid. If it did not dissolve in water, but appears to be
dissolving in nitric acid, it is undergoing a chemical reaction. In
general, compounds that contain anions that are the conjugate bases of
weak acids will react (unless the compounds are very insoluble). For
example:
CaCO3(s) + 2 H+(aq) → Ca2+(aq) + H2O(l) + CO2(g, colorless, odorless)
NiS(s) + 2 H+(aq) → Ni2+(aq) + H2S(g, colorless, rotten egg smell)
Ca3(PO4)2(s) + 6 H+(aq) → 3 Ca2+(aq) + 2 H3PO4(aq)The remaining tests must be perfomed on a solution of the compound.
If the compound is insoluble in water, dissolve it in nitric acid. Otherwise, dissolve in water.
- Reaction with sodium hydroxide
Add NaOH dropwise to the solution, stir or shake the solution, and
observe any reaction (if the compound was dissolved in nitric acid, the
first several drops will neutralize the acid so be sure to check the pH
with litmus paper). Look for a precipitate (refer to the solubility
rules for hydroxides). If a precipitate forms, continue adding NaOH.
Some metal hydroxides are amphoteric and will form a complex ion and
redissolve. See Figures 1, 2, and 3 for an example of this reaction.
The following ions are amphoteric:
Species Acidic Solution Slightly Basic Solution Basic Solution Al3+ Al3+(aq) Al(OH)3(s) Al(OH)4-(aq) Cr3+ Cr3+(aq) Cr(OH)3(s) Cr(OH)4-(aq) Pb2+ Pb2+(aq) Pb(OH)2(s) Pb(OH)42-(aq) Zn2+ Zn2+(aq) Zn(OH)2(s) Zn(OH)42-(aq) Sn4+ Sn4+(aq) Sn(OH)4(s) Sn(OH)62-(aq) - Reaction with ammonia
Add NH3 dropwise to the solution, stir or shake the
solution, and observe any reaction. If a metal hydroxide precipitate
forms, continue adding ammonia. Some metal hydroxides fom a complex ion
and redissolve. See Figures 4, 5, and 6 for an example of this
reaction. The following ions form ammonia complexes:
Acid Solution Basic Solution Solution with Excess NH3 Color of Complex Ni2+(aq) Ni(OH)2(s) Ni(NH3)62+(aq) violet Cu2+(aq) Cu(OH)2(s) Cu(NH3)42+(aq) blue Zn2+(aq) Zn(OH)2(s) Zn(NH3)62+(aq) colorless Ag+(aq) Ag2O(s) Ag(NH3)2+(aq) colorless Cd2+(aq) Cd(OH)2(s) Cd(NH3)42+(aq) colorless - Reaction with hydrochloric acid
Add HCl dropwise until solution tests acidic to litmus paper and
observe any reaction. A precipitate will form with any cation that
forms an insoluble chloride (refer to the solubility rules). For
example:
Pb2+ + 2Cl- → PbCl2(s) - Reaction with sulfuric acid
Add H2SO4 dropwise until solution is acidic
and observe any reaction. A precipitate will form with any cation that
forms an insoluble sulfate (refer to the solubility rules). For example:
Ba2+ + SO42- → BaSO4(s) - Reaction with silver nitrate
Add HNO3 dropwise until solution is acidic (unless of course it was dissolved in nitric acid), then add a few drops of AgNO3
and observe any reaction. A precipitate will form with certain cations
that form insoluble silver compounds, but because of the acidic
environment, some insoluble silver salts (e.g. salts containing CO32-, S2-, and PO43- ions) are "destroyed." Cl-, Br-, and I- form insoluble compounds, while SO42- forms a moderately insoluble compound.
Ag+ + Cl- → AgCl(s) - Reaction with barium nitrate
Add HNO3 dropwise until solution is acidic, boil the
solution for two minutes, then test with litmus paper. Continue adding
and boiling until solution remains acidic after boiling. Cool the
solution and add a few drops of Ba(NO3)2 and
observe any reaction. A precipitate will form with anions that form an
insoluble barium compound (except the ones destroyed by acid as in the
above test).
- Specific Tests
Sometimes the above tests can not definitively confirm the presence
of a specific ion. In these cases, it is necessary to do specific tests
for a particular ion.
Example Unknown Salts
Sample 1 had the following characteristics:
- Visual test: white crystalline powder
- Heat test: brown gas given off
- Flame test: no color identified
- Solubility in water: soluble
- Nitric acid: soluble
- Hydroxide: formed an insoluble white precipitate, then dissolved with excess NaOH
- Ammonia: formed an insoluble white precipitate which did not redissolve in excess NH3
- Hydrochloric acid: formed an insoluble white precipitate
Analysis of observations:
- The brown gas given off during the heat test indicates presence of the NO3- ion, since the NO3- ion reacts to form brown NO2 gas as shown below:
2 Pb(NO3)2(s) + heat → 2 PbO(s) + O2(g) + 4 NO2(g)
- The white precipitate which formed and then redissolved with the
addition of sodium hydroxide indicates the presence of an amphoteric
cation. The possibilities are Al, Cr, Pb, Zn, and Sn. The reaction
occurs as follows:
Cr3+(aq) + 3 OH-(aq) → Cr(OH)3(s)
Cr(OH)3(s) + OH-(aq) → Cr(OH)4-(aq) - Since the cation does not form an ammonia complex, it eliminates Zn from the list of possible cations established above.
- There are only three cations which form precipitates with hydrochloric acid, Pb2+, Ag+, and Hg22+. Of these three, only Pb2+ is amphoteric. The reaction with chloride occurs as follows:
Pb2+ + 2 Cl- → PbCl2(s)
- The brown gas given off during the heat test indicates presence of the NO3- ion, since the NO3- ion reacts to form brown NO2 gas as shown below:
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