Qualitative Analysis

Lab 10
Page 393-411
 Pre-Lab
• Page 402
• Complete all pre-lab questions

• Quiz 6 –Next week

– Factors Affecting Reaction Rates
 Introduction to Qualitative Analysis

• The purpose of this experiment is to study the reactions of

6 cations. You will use your observations to identify the
ions in samples of unknown composition.
• Qualitative analysis is used to separate and detect cations
and anions in a sample substance.
• Qualitative analysis is the procedure by which one can
determine the nature, but not the amount of species in a
mixture. This experiment deals with tests isolated for
specific ions in solution.
 Description of Lab
• Knowledge of chemical reactivity to
identify 6 metal cations
• Written substantiation of the method used to
identify each solid such as a qualitative
analysis flow chart
 Experimental Procedure
• You will be provided with 2 test-tubes.
• One will contain all six of the cations to be
detected. This is your reference solution
• The other test-tube will contain a number of
unknown cations that you will need to
identify by experimental observations.
• Make careful notes of your observations.
• Look closely for gases, and note colors of
precipitates
• Types of reactions encountered in qualitative
analysis include
– precipitation
– complex ion formation
– redox reactions
–acid-base reactions

• Some means of identifying ions by qualitative

analysis are:
– color changes,
–evolution of gas,
–change in pH (acidity or basicity)
–or ability to redissolve a precipitate by addition of a
complexing ligand.
 Common grouping of cations

• Group I: Ag+, Hg22+, Pb2+

Precipitated in 1 M HCl

• Group II: Bi3+, Cd2+, Cu2+, Hg2+, (Pb2+), Sb3+ and

Sb5+, Sn2+ and Sn4+
Precipitated in 0.1 M H2S solution at pH 0.5
 Common grouping of cations

• Group III: Al3+, (Cd2+), Co2+, Cr3+, Fe2+ and Fe3+,

Mn2+, Ni2+, Zn2+
Precipitated in 0.1 M H2S solution at pH 9

• Group IV: Ba2+, Ca2+, K+, Mg2+, Na+, NH4+

Ba2+, Ca2+, and Mg2+ are precipitated in 0.2 M
(NH4)2CO3 solution at pH 10; the other ions are
soluble
Common Qualitative Analysis Reagents

• Many reagents are used in qualitative analysis, but

only a few are involved in nearly every group
procedure.
• The four most commonly used reagents are HCl,
HNO3, NaOH and NH3.
• Understanding the uses of the reagents is helpful
when planning an analysis.
Reagent Effect
HCl Increases [H+] Increases [Cl-]
Decreases [OH-]
Dissolves insoluble carbonates,
chromates, hydroxides, some
sulfates
Destroys hydroxo and NH3
complexes
Precipitates insoluble chlorides

HNO3 Increases [H+] Decreases [OH-]

Dissolves insoluble carbonates,
chromates, and hydroxides
Dissolves insoluble sulfides by
oxidizing sulfide ion
Destroys hydroxo and ammonia
complexes
Good oxidizing agent.
 Increases [OH-] Decreases
[H+]
Forms hydroxo complexes
NaOH Precipitates insoluble
hydroxides

NH3 Increases [NH3] Increases

[OH-] Decreases [H+]
Precipitates insoluble
hydroxides
Forms NH3 complexes
Forms a basic buffer with
NH4+
 Process for Salt Identification

• Appearance of compound
• Heating effect
• Flame test
• Solubility in water
• Reaction with nitric acid
• The remaining tests must be performed on a
solution of the compound.
Reactions to be performed on a solution of the
solid:

– Reaction with sodium hydroxide

– Reaction with ammonia
– Reaction with hydrochloric acid
– Reaction with sulfuric acid
– Reaction with silver nitrate
– Reaction with barium nitrate
 Appearance of Compounds

• 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
• Ions of transition elements tend to be
colored.
Ion Color

Co2+ rose

Co3+ violet

Cr3+ violet

Cu2+ blue

Fe2+ pale green, pale violet

Fe3+ yellow-brown

Mn2+ pale pink

Ni2+ blue-green
 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.
• Hydrated salts tend to be sparkly, and have well
defined crystalline facets.
• CuSO4•xH2O blue → CuSO4 (white)
• 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)
• CaCO3(s) + heat --> CaO(s) + CO2(g, colorless,
odorless)
• CaSO3(s) + heat --> CaO(s) + SO2(g, colorless,
pungent)
• Some bromides and iodides decompose to give
Br2(g, orange-brown) and I2(g, purple).
 Flame Tests
• 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.

• A flame test can be used as a confirmatory test.

• One problem, is that sodium is often an impurity so will

almost always see a yellow flame. Therefore careful
observations need to be recorded.
 Sodium Bright yellow (intense, persistant)

Potassium Pale violet (slight, fleeting)

Calcium Brick red (medium, fleeting)

Strontium Crimson (medium)

Barium
Light green

Lead Pale bluish (slight, fleeting)

Copper Green or blue (medium, persistant)

Solubility
• 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
 Solubility Guidelines
1. All nitrates are soluble.
2. Practically all sodium, potassium, and ammonium salts
are soluble.
3. All chlorides, bromides, and iodides are soluble except
those of silver, mercury(I), and lead(II).
4. 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.
5. 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).
• The remaining tests must be performed on a solution of the
compound, usually in water. If the compound is not water
soluble, use nitric acid.

Reaction with NaOH

• 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.
 Amphoteric Ions
• Having the characteristics of an acid and a
base and capable of reacting chemically
either as an acid or a base.
• Whether an amphoteric chemical acts as an
acid or a base depends on what other
chemicals happen to be around.
• If a base ( like NH3) is present, water can act as
an acid and react by donating a proton to that
base. In doing so, water is changed into its
conjugate base, hydroxide ion.
H2O + NH3 → NH4+ + OH-

• If an acid (like HCl) is present, water can act as a

base and react by accepting a proton from that
acid. In doing so, water is changed into its
conjugate acid, hydronium ion.
H2O + HCl → Cl- + H3O+
 Amphoteric Ions
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+(s) 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)

 Excess NaOH is added Al(OH)3 precipitates
Al(NO3)3 in with the addition of
solution the precipitate redissolves
NaOH.
as the Al(OH)4- complex ion
is formed.
 The remaining tests must be performed on a
solution of the compound, usually in water. If
compound is not water soluble, use nitric acid.

Reaction with NaOH

• 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.
Reaction with ammonia

• Add NH3 dropwise to the solution, stir or

shake the solution.
• Observe any reaction.
• If a metal hydroxide precipitate forms,
continue adding ammonia.
• Some metal hydroxides form a complex ion
and redissolve.
 Complexes with Ammonia
Acid Basic Solution with Excess Color of
Solution Solution NH3 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)42+(aq) colorless

Ag+(aq) Ag2O(s) Ag(NH3)2+(aq) colorless

Cd2+(aq) Cd(OH)2(s) Cd(NH3)42+(aq) colorless

 Reaction with HCl or H2SO4

• 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).
Pb2+ + 2Cl- --> PbCl2(s)

• Add H2SO4 dropwise until solution is acidic and observe

any reaction. A precipitate will form with any cation that
forms an insoluble sulfate
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
Ag+ + Cl- --> AgCl(s)
• SO42- forms a moderately insoluble compound.
 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 (e.g. CH3COOH and H2S
are acidic, conjugate bases are CH3COO- and HS- . These
will generally react (unless the compounds are very
insoluble).
 Recap: Conjugate pairs of acids and bases.
• When an acid gives up its proton, what remains is called
the conjugate base of that acid. When a base accepts a
proton, the resulting chemical is called the conjugate acid
of that original base. HF and F- are a conjugate acid-base
pair. H2O and H3O+ are a conjugate pair, where H3O+ is the
acid and H2O is the base.

• 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)
 Na+, K+, NH4+, Ag+, Cu2+, Bi3+

xxxxxx 2
xxxxxx 1

NH3(g)
[CaO, (NH4)2CO3] AgCl
Cu2+, Bi3+
NH3
∆,H2O
xxxxx [CH3CSNH2, heat) xxxx] 4
3

[flame] CuS and Bi2S3

xxxxxx 6
5 xxxx
Na+ K+ HNO3
AgCl
S Cu2+, Bi3+
NH3

[Sn(OH)42-, OH- [CH3COOH, K4[Fe(CN)6]

8 9
• From the ions on the top line, you need to
perform chemical tests, that will allow you
the identification of each metal ion.

• You will not need to memorize the flow

chart but in a quiz or exam will need to be
able to put reagents and products into an
incomplete flow-chart.
 Completing the flow chart:
1. The first product isolated is NH3 (g).
Looking at the cations, what could we add
that would give us NH3(g)
2. Calcium oxide (basic anhydride) and
ammonium chloride are reacted with the
cations to precipitate all cations as the
hydroxide except Na+, K+ and NH4+.
• One more reagent needs to be added
to the initial solution (2). This needs to
form AgCl. What would react with Ag+ to
give AgCl. Fill in space 2.

• Product 3 = The reaction of AgCl + NH3

• Product 4: 2 reagents are added. H2S is an acidic
anhydride and will form and acidic solution. Metal
II Sulfides: (Metal has oxidation state of +2; for
example, M = Cu2+, Fe2+)
• MS(s) → M2+ (aq) + S2–(aq)
S2–(s) + H2O(l) → HS-(aq) + OH–(aq)
• The S2– ion is a strong base and will react
immediately to form HS– and a hydroxide ion. The
concentration of S2– in solution is negligible
• MS(s) + H2O(l) → M2+ (aq) + HS–(aq) + OH-.
• Ksp for the dissolution of a metal II sulphide is :
» Ksp = [M2+][HS–][OH–]
• The addition of acid will use up OH– and hence,
shift the summed equilibrium to the right thus
dissolving more of the salt (MS).

• Since the Solubility is higher in acid solution and

quite low in base solution, it is often more
convenient (and conventional) to rewrite the
equation for the dissolution in an acidic solution.

• As the pH is lowered (higher H3O+ concentration)

the solubility of the metal sulfide increases.
• Product 5: Potassium has been identified by a
positive flame test. What technique could be
used to confirm the Na+ presence?

• Reactant 6: Needs to be something that will

precipitate out the product AgCl after the earlier
reaction of silver + (aq) ammonia. The initial
reaction was performed to separate out Ag from
the other heavier cations.
• Product 7: Bi3+ / Cu2+ + NH3(aq) → ? Consider, will
one of the products be soluble or insoluble . (This
reaction was performed in the metathesis lab, look back
at the extra sheet of questions)

• Product 8: The Bi3+ in the bismuth hydroxide is reduced,

and the tin hydroxide compound is the reducing agent,
therefore oxidized itself to [Sn(OH)6]2-

• Product 9: Potassium hexacyanoferrate = K4[Fe(CN) 6] is

used a confirmatory test for Cu, forming a reddish brown
ppt. (see page 396 Beran).