Welcome to Chemical Detective Work!

Welcome to the exciting world of Chemical Tests! In this chapter, you are learning how to become a chemical detective. When you have an unknown substance, you need reliable tests—or clues—to figure out exactly what it is.
This knowledge is crucial not only for exams but also for understanding quality control, environmental testing, and laboratory safety. Don't worry if this seems like a lot of information; we will break down each test into simple, memorable steps!

Section 1: Tests for Common Gases

Identifying gases is often the easiest part of chemical testing, as the results are usually dramatic and unique.

1. Testing for Oxygen (\(O_2\))

  • The Test: Place a glowing splint into the gas sample.
  • The Result (The Clue): The splint relights (bursts back into flame).
  • Explanation: Oxygen is highly flammable and supports combustion, allowing the smoldering carbon (the glow) on the splint to reignite.

2. Testing for Hydrogen (\(H_2\))

  • The Test: Place a lit splint into the gas sample.
  • The Result (The Clue): You hear a distinctive high-pitched squeaky pop.
  • Explanation: Hydrogen is highly flammable. The 'pop' is caused by the rapid, small explosion as the hydrogen gas combines instantly with oxygen in the air due to the heat of the flame, forming water (\(2H_2 + O_2 \rightarrow 2H_2O\)).

3. Testing for Carbon Dioxide (\(CO_2\))

  • Prerequisite Concept: The reagent used is limewater, which is an aqueous solution of Calcium Hydroxide, \(Ca(OH)_2\).
  • The Test: Bubble the gas through the limewater solution.
  • The Result (The Clue): The limewater turns cloudy or milky white (a white precipitate forms).
  • Explanation: Carbon dioxide reacts with calcium hydroxide to form insoluble Calcium Carbonate (\(CaCO_3\)), which is the white solid responsible for the cloudiness.
    \(Ca(OH)_2 (aq) + CO_2 (g) \rightarrow CaCO_3 (s) + H_2O (l)\)

4. Testing for Ammonia (\(NH_3\))

Ammonia is an alkaline gas. Think of its properties when choosing the test.

  • The Test: Place a piece of damp red litmus paper into the gas.
  • The Result (The Clue): The litmus paper turns blue.
  • Mnemonic Tip: Ammonia has an unpleasant smell, but it is a base (alkaline), so it turns red litmus blue.

5. Testing for Chlorine (\(Cl_2\))

  • The Test: Place a piece of damp blue litmus paper into the gas.
  • The Result (The Clue): The litmus paper turns red and then is bleached white.
  • Did You Know? Chlorine is a powerful bleaching agent and is used in disinfectants and water purification.
Quick Review: Gases

O₂ → Relights glowing splint
H₂ → Squeaky pop with lit splint
CO₂ → Turns limewater milky

Section 2: Tests for Anions (The Negative Ions)

Anions are usually tested using precipitation reactions—mixing two clear solutions to form an insoluble solid (a precipitate).

1. Testing for Carbonates (\(CO_3^{2-}\))

Since carbonates contain carbon, we test for the release of \(CO_2\).

  • The Test: Add a few drops of dilute acid (e.g., dilute nitric acid) to the unknown substance.
  • The Observation: Effervescence (fizzing) is seen, indicating gas is released.
  • The Confirmation: Bubble the gas produced through limewater. If the limewater turns milky, the substance contained carbonate ions.

2. Testing for Halides (Chloride, Bromide, Iodide)

These tests rely on forming insoluble silver halides using Silver Nitrate solution.

Step 1 (Crucial Preparation): Always add dilute nitric acid (\(HNO_3\)) first. This removes any unwanted carbonates or sulfites that could also form a white precipitate with silver nitrate, preventing a false positive result.

Step 2 (The Reaction): Add a few drops of aqueous Silver Nitrate (\(AgNO_3\)) solution.

The Results:

  • Chloride Ion (\(Cl^-\)): Forms a white precipitate (Silver Chloride, \(AgCl\)).
  • Bromide Ion (\(Br^-\)): Forms a cream precipitate (Silver Bromide, \(AgBr\)).
  • Iodide Ion (\(I^-\)): Forms a yellow precipitate (Silver Iodide, \(AgI\)).

Memory Aid (Colour Order): Think of traffic lights going from light to dark:
White (Cl⁻) $\rightarrow$ Cream (Br⁻) $\rightarrow$ Yellow (I⁻)

3. Testing for Sulfates (\(SO_4^{2-}\))

Sulfates are tested using Barium Chloride or Barium Nitrate solution.

Step 1 (Crucial Preparation): Add dilute hydrochloric acid (\(HCl\)) first. (Similar to the halide test, this removes any unwanted carbonates).

Step 2 (The Reaction): Add a few drops of aqueous Barium Chloride (\(BaCl_2\)) solution.

  • The Result (The Clue): A thick white precipitate forms (Barium Sulfate, \(BaSO_4\)).
  • Crucial Fact: This white precipitate must be insoluble in excess dilute acid to confirm it is a sulfate.
Common Mistake to Avoid!

If you forget to add the acid before testing for Sulfates or Halides, and you happen to have a Carbonate present, you will get a white precipitate anyway (a false positive!). Always acidify first!

Section 3: Tests for Cations (The Positive Metal Ions)

Cations are usually metal ions. We test for them by adding a fixed amount of Sodium Hydroxide solution (\(NaOH\)) and aqueous Ammonia solution (\(NH_3\)) and observing the colour and the solubility of the resulting precipitate.

The precipitate formed is the metal hydroxide (e.g., \(Fe^{3+} + 3OH^- \rightarrow Fe(OH)_3\)).

1. Testing for Iron Ions (\(Fe^{2+}\) and \(Fe^{3+}\))

Iron ions are the easiest to identify because they have distinct, strong colours.

  • Iron(II) Ion (\(Fe^{2+}\)):
    • With NaOH: Forms a dirty green precipitate.
    • In Excess NaOH: Precipitate remains insoluble.
  • Iron(III) Ion (\(Fe^{3+}\)):
    • With NaOH: Forms a characteristic red-brown precipitate (like rust).
    • In Excess NaOH: Precipitate remains insoluble.

2. Testing for Copper Ion (\(Cu^{2+}\))

  • With NaOH: Forms a light blue precipitate.
  • In Excess NaOH: Precipitate remains insoluble.
  • With Aqueous Ammonia (\(NH_3\)): Forms the light blue precipitate, but if excess ammonia is added, the precipitate dissolves to form a deep, clear dark blue solution.

3. Testing for Calcium Ion (\(Ca^{2+}\))

  • With NaOH: Forms a white precipitate.
  • In Excess NaOH: Precipitate remains insoluble.
  • Note: Calcium hydroxide is slightly soluble, so the precipitate may appear sparse or faint.

4. The "Amphoteric" Cations (Aluminium and Zinc)

Some hydroxides are amphoteric, meaning they can react with both acids (like normal bases) AND strong bases (like sodium hydroxide). When you add excess strong base, the hydroxide precipitate dissolves.

Aluminium Ion (\(Al^{3+}\))
  • With NaOH: Forms a white precipitate.
  • In Excess NaOH: The precipitate dissolves, forming a clear, colourless solution.
  • With Aqueous Ammonia (\(NH_3\)): Forms a white precipitate.
  • In Excess \(NH_3\): Precipitate remains insoluble.

Memory Trick: Aluminium Only Dissolves in NaOH.

Zinc Ion (\(Zn^{2+}\))

Zinc is the 'social butterfly' of the cations—it dissolves everywhere!

  • With NaOH: Forms a white precipitate.
  • In Excess NaOH: The precipitate dissolves, forming a clear, colourless solution.
  • With Aqueous Ammonia (\(NH_3\)): Forms a white precipitate.
  • In Excess \(NH_3\): The precipitate dissolves, forming a clear, colourless solution.

Summary Table for Cations (MUST KNOW!)

This table is the most important part of this section. Practice drawing it until you know it by heart.

Ion Result with NaOH (Few Drops) Result with Excess NaOH Result with Aqueous Ammonia (Few Drops) Result with Excess Aqueous Ammonia
\(Fe^{2+}\) (Iron II) Dirty Green ppt Insoluble Dirty Green ppt Insoluble
\(Fe^{3+}\) (Iron III) Red-Brown ppt Insoluble Red-Brown ppt Insoluble
\(Cu^{2+}\) (Copper II) Light Blue ppt Insoluble Light Blue ppt Dissolves to deep blue solution
\(Al^{3+}\) (Aluminium) White ppt Dissolves White ppt Insoluble
\(Zn^{2+}\) (Zinc) White ppt Dissolves White ppt Dissolves
\(Ca^{2+}\) (Calcium) White ppt (sparse) Insoluble No ppt / very slight No ppt / very slight
Key Takeaway: Distinguishing Cations

If the precipitate is white, you must check solubility in excess reagents to determine if it is \(Ca^{2+}\), \(Al^{3+}\), or \(Zn^{2+}\). The colours (Green, Red-Brown, Blue) immediately identify the Iron and Copper ions.

Section 4: Testing for Water (\(H_2O\))

Finally, we need tests to prove the presence of water, especially to distinguish between anhydrous (water-free) and hydrated (containing water) substances.

1. Using Anhydrous Copper(II) Sulfate

  • The Test: Place the suspected water sample onto white anhydrous Copper(II) Sulfate (\(CuSO_4\)).
  • The Result (The Clue): The white solid turns blue.
  • Explanation: The anhydrous salt absorbs the water to become the hydrated salt (\(CuSO_4 \cdot 5H_2O\)), which has a blue colour. This is a reversible process (heating blue copper sulfate drives off water, returning it to white).

2. Using Cobalt(II) Chloride Paper

  • The Test: Place the suspected water sample onto blue Cobalt(II) Chloride paper.
  • The Result (The Clue): The paper turns pink.
  • Did You Know? This test is sometimes used in damp indicators to show if an area has high humidity!

Keep practicing these tests until the colours and reagents become second nature. You’ve got this!