Qualitative Inorganic Analysis

Goals and Scope of Qualitative Inorganic Analysis

Qualitative inorganic analysis is concerned with which ions or elements are present in a (usually unknown) inorganic sample, not how much of them there is. It answers questions like:

• “Does this solution contain chloride?”
• “Which metal cations are present in this mixture?”
• “Is there sulfate, nitrate, or carbonate in this salt?”

In contrast to quantitative methods, the result is usually a list of detected ions or elements (the ion spectrum of the sample). Classical qualitative analysis uses simple laboratory equipment and relies mainly on wet‑chemical reactions and visual observation (precipitate, color change, gas formation, etc.).

This chapter focuses on inorganic qualitative analysis of ions in aqueous solution, not on organic functional group tests or instrumental methods, which are treated elsewhere.

General Strategy and Typical Workflow

Qualitative inorganic analysis follows a fairly systematic approach. A typical workflow is:

1. Sample reception and preliminary observations
• State: solid, liquid, mixture?
• Color, odor, texture.
• Solubility in water, pH of the solution.
• Behavior on gentle heating (decomposition, melting, gas evolution).
2. Dissolution and preparation of a test solution
• Dissolve the solid in water or, if necessary, in a suitable acid (e.g. dilute HCl or HNO₃).
• Remove insoluble residues by filtration.
• Obtain a clear solution containing the ions to be identified.
3. Separation of cation groups (systematic group analysis)
• Add specific group reagents that selectively precipitate certain categories of cations (e.g. sulfide‑forming, hydroxide‑forming, etc.).
• Separate the precipitates from the remaining solution.
• Work up each group separately.
4. Confirmation of individual cations within each group
• Subject each group’s precipitate (or its solution) to specific confirmatory tests.
• Perform appropriate blank and comparison tests when possible.
5. Detection of common anions
• Use dedicated tests for specific anions (e.g. carbonate, sulfate, halides).
• If necessary, destroy interfering anions before testing for others.
6. Evaluation, plausibility checks, and documentation
• Check whether all observations are consistent with the proposed ions.
• Rule out incompatible combinations.
• Summarize results in a clear and structured report.

The precise sequence and choice of tests can vary depending on the sample type and educational or laboratory tradition, but the underlying logic—gradual restriction of possibilities through selective reactions—is common to most schemes.

Basic Testing Techniques and Observations

Qualitative analysis utilizes a limited set of basic operations that generate observable changes:

• Precipitation reactions
  Formation of a solid when two solutions are mixed.

Example:
$$\text{Ag}^+_{(aq)} + \text{Cl}^-_{(aq)} \rightarrow \text{AgCl}_{(s)} \downarrow$$

Observations: color, texture, and behavior of the precipitate under further treatments.

• Complex formation
  Color change or increased solubility upon addition of a ligand (e.g. ammonia, EDTA, thiocyanate).

Example: soluble deep blue complex with copper(II) and ammonia:
$$\text{Cu}^{2+} + 4\,\text{NH}_3 \rightleftharpoons [\text{Cu}(\text{NH}_3)_4]^{2+}$$

• Acid–base reactions
  Changes in solubility or color with variation of pH (e.g. dissolution of hydroxide precipitates in excess base or acid).
• Redox reactions
  Change in oxidation state accompanied by color change, gas evolution, or precipitation. Often used in specific tests for particular ions.
• Gas evolution
  Detection of certain anions or reactions by observing formation of gases and their properties (odor, effect on indicators, extinguishing or supporting combustion).
• Flame tests
  Coloration of a flame by certain metal cations (e.g. Na, K, Ba, Sr, Cu) used as a quick screening tool.

Correct observation and interpretation of these phenomena are central skills in classical qualitative analysis.

Systematic Cation Group Analysis

In many classical schemes, cations are first divided into groups based on their behavior toward certain group reagents. The exact definitions of the groups can differ by tradition; the following is a common student‑laboratory scheme.

Group Reagents and Precipitation Order

Group reagents are selected so that:

• Only certain cations form a precipitate under the specified conditions.
• Other cations remain in solution and are carried forward to the next step.
• Later groups do not interfere with earlier ones.

A typical grouping scheme (for illustrative purposes) may look like this:

1. Group I cations (e.g. Ag⁺, Pb²⁺, Hg₂²⁺)
• Group reagent: dilute HCl
• Principle: Formation of poorly soluble chlorides.
  $$\text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl}\downarrow$$
2. Group II cations (e.g. Cu²⁺, Cd²⁺, Bi³⁺, Hg²⁺, Pb²⁺ (re‑appearing), etc.)
• Group reagent: H₂S in acidic solution
• Principle: Formation of sparingly soluble metal sulfides under acidic conditions.
3. Group III cations (e.g. Fe³⁺, Al³⁺, Cr³⁺, sometimes others)
• Group reagent: NH₃ or NH₄OH with NH₄Cl buffer
• Principle: Precipitation as hydroxides (e.g. Fe(OH)₃, Al(OH)₃) at controlled pH, while other cations remain soluble.
4. Group IV cations (e.g. Ni²⁺, Co²⁺, Zn²⁺, Mn²⁺)
• Group reagent: H₂S or (NH₄)₂S in neutral or slightly basic solution
• Principle: Precipitation as metal sulfides under conditions where earlier groups have been removed.
5. Group V cations (e.g. Ca²⁺, Sr²⁺, Ba²⁺)
• Group reagent: (NH₄)₂CO₃ (with NH₃/NH₄Cl buffer)
• Principle: Precipitation as carbonates.
6. Group VI cations (e.g. Mg²⁺, Na⁺, K⁺, NH₄⁺ and others that stay in solution)
• No common group reagent; these are identified by specific tests.

This scheme is a working tool for practical analysis rather than a fundamental chemical classification. Its main purpose is to structure the analysis so that complex mixtures become manageable.

Working Up Individual Groups

After the group separation, each precipitate is examined separately with selective confirmatory tests:

• Dissolve the group precipitate under chosen conditions (e.g. with acid or complexing agent).
• Carry out a succession of specific reactions that:
• produce characteristic colors or precipitates,
• use selective complex formation or redox changes,
• confirm or exclude particular cations.

The sequence within each group aims to:

• Minimize loss of analytes.
• Limit mutual interference.
• Maintain clarity and logical progression.

Examples of Important Cation Tests

Only a few representative tests are outlined here; full schemes are typically much more extensive and are presented stepwise in laboratory manuals.

Silver, Lead, and Mercury(I) (Chloride Group)

After precipitation with HCl and washing:

• Solubility behavior of the chlorides
• $\text{PbCl}_2$ is more soluble in hot water than $\text{AgCl}$ and $\text{Hg}_2\text{Cl}_2$.
• Separation: extract with hot water and test the filtrate for Pb²⁺ (e.g. by forming $\text{PbCrO}_4$ with chromate).
• Ammonia treatment of the residue
• $\text{AgCl}$ dissolves in dilute NH₃ forming a colorless complex $[\text{Ag}(\text{NH}_3)_2]^+$.
• $\text{Hg}_2\text{Cl}_2$ with NH₃ often yields a black mixture (formation of finely divided Hg and basic mercury(II) amidochloride), a characteristic observation.

These behaviors allow logical discrimination among Ag⁺, Pb²⁺, and Hg₂²⁺.

Iron(III), Aluminum, and Chromium (Hydroxide Group)

These cations are commonly precipitated as hydroxides at moderate pH (in presence of ammonium salts), forming gelatinous solids:

• $\text{Fe(OH)}_3$: reddish‑brown
• $\text{Al(OH)}_3$: white
• $\text{Cr(OH)}_3$: greenish

Typical distinguishing behaviors:

• Fe³⁺
• Reacts with thiocyanate $\text{SCN}^-$ to give a blood‑red complex:
  $$\text{Fe}^{3+} + 3\,\text{SCN}^- \rightleftharpoons [\text{Fe}(\text{SCN})_3]$$
• Al³⁺
• Hydroxide dissolves in excess strong base (e.g. NaOH) forming aluminate:
  $$\text{Al(OH)}_3 + \text{OH}^- \rightarrow [\text{Al(OH)}_4]^-$$
• Cr³⁺
• In basic oxidizing medium (e.g. with hydrogen peroxide), converts to chromate $\text{CrO}_4^{2-}$, producing yellow solution.

These signature reactions, combined with controlled pH manipulation, enable selective identification.

Alkali and Alkaline Earth Metals

These are more difficult to separate by simple precipitation reactions (or do not precipitate at all under typical conditions). Common tests include:

• Flame tests
• Na⁺: intense yellow flame.
• K⁺: violet flame (often observed through cobalt glass).
• Ca²⁺: orange‑red.
• Sr²⁺: crimson.
• Ba²⁺: yellow‑green.
• Characteristic precipitates
• Ba²⁺: formation of insoluble barium sulfate with sulfate solution ($\text{BaSO}_4$).
• Ca²⁺ and Sr²⁺: different solubilities of their salts, e.g. with carbonate or sulfate, used comparatively.

Such tests are often used after more complex cations have already been removed.

Detection of Common Anions

Anion analysis is often carried out separately from cation analysis. The presence of acids in the cation analysis and the necessity of avoiding precipitation of undesired salts mean that special conditions are required. Typical examples:

Carbonate and Hydrogen Carbonate

• Add dilute acid (e.g. HCl) to the solid or solution.
• Observation: vigorous effervescence due to liberation of $\text{CO}_2$ gas:
  $$\text{CO}_3^{2-} + 2\,\text{H}^+ \rightarrow \text{H}_2\text{O} + \text{CO}_2 \uparrow$$
• $\text{CO}_2$ can be confirmed by turning limewater (saturated $\text{Ca(OH)}_2$ solution) milky from $\text{CaCO}_3$ formation.

Sulfate

• First, ensure absence of interfering anions (especially carbonate or sulfite that also precipitate with barium).
• Add $\text{Ba}^{2+}$ solution (e.g. $\text{BaCl}_2$) under acidic conditions.
• Observation: formation of a fine white precipitate of barium sulfate:
  $$\text{Ba}^{2+} + \text{SO}_4^{2-} \rightarrow \text{BaSO}_4 \downarrow$$
  which is very insoluble in dilute acids and difficult to dissolve.

Chloride, Bromide, Iodide (Halides)

• Typically, add $\text{AgNO}_3$ solution to a slightly acidified sample.
• $\text{AgCl}$: white precipitate, soluble in dilute ammonia.
• $\text{AgBr}$: cream precipitate, only partly soluble in ammonia.
• $\text{AgI}$: yellow precipitate, essentially insoluble in ammonia.
• Under controlled conditions, further specific tests (often involving oxidation or reduction) can differentiate between halides more clearly.

Nitrate

Nitrate is often tested after other interfering anions are removed or destroyed. One classical test is the brown ring test:

• Mix the solution with ferrous sulfate (Fe²⁺) solution.
• Carefully add concentrated $\text{H}_2\text{SO}_4$ along the side of the tube so that it forms a layer below the aqueous phase.
• Observation: formation of a brown ring at the phase boundary, due to formation of a nitrosyl iron complex.
  (This test must be carried out with caution due to concentrated acid.)

Numerous other classical tests exist (e.g. for sulfide, nitrite, phosphate, etc.), each with distinctive visual signals.

Interferences and Sample Pretreatment

Real samples often contain multiple ions that may interfere with one another’s tests. Common strategies to manage interferences include:

• Masking agents
• Add ligands that form stable complexes with interfering ions, keeping them in solution while the ion of interest reacts.
• Example: Using complexing agents to keep Fe³⁺ in solution when precipitating certain other metals.
• Separation by precipitation or extraction
• Remove problematic components before specific tests.
• Example: precipitating carbonate as $\text{CaCO}_3$ before testing for sulfate with $\text{Ba}^{2+}$.
• Destruction of interfering anions
• Oxidation or reduction to convert them into forms that do not interfere with subsequent tests (for instance, oxidizing sulfite to sulfate under conditions where one is testing for carbonate).
• pH control and buffering
• Adjust pH so that only desired ions precipitate or react.
• Use buffering systems (e.g. $\text{NH}_4^+/\text{NH}_3$) to maintain nearly constant pH.

Careful planning of the order of tests is crucial to avoid “consuming” analytes or generating misleading precipitates.

Safety and Good Laboratory Practice in Qualitative Analysis

Classical qualitative inorganic analysis often uses:

• Strong acids and bases.
• Sulfide reagents (which can release toxic $\text{H}_2\text{S}$ gas).
• Heavy metal salts (e.g. lead, mercury, cadmium, barium).
• Concentrated oxidizing or reducing agents.

This calls for rigorous safety measures:

• Work in a fume hood when evolving or using toxic gases (e.g. $\text{H}_2\text{S}$).
• Wear appropriate personal protective equipment (lab coat, goggles, gloves).
• Handle heavy metal solutions with care and collect waste separately for appropriate disposal.
• Avoid ingestion or inhalation of powders and aerosols.
• Clean glassware thoroughly to prevent cross‑contamination, which can lead to false positives or negatives.

Additionally, documentation of every step—reagent used, conditions, observations—is essential for reproducibility and for tracing the logic of the analysis.

Role and Limitations of Classical Qualitative Inorganic Analysis

Despite widespread use of instrumental techniques, classical qualitative analysis remains important because it:

• Teaches core chemical concepts such as solubility, complex formation, redox behavior.
• Trains observational skills and logical reasoning.
• Requires only basic equipment and inexpensive reagents.

However, limitations must be recognized:

• Detection limits are typically higher than in modern instrumental methods.
• Complex mixtures can lead to ambiguous or inconclusive results.
• Subjective evaluation of color and intensity may introduce error.

In modern practice, classical qualitative tests are often used as preliminary screening or educational tools, while detailed and highly sensitive analyses are carried out with instrumental techniques discussed elsewhere.