Identification of ions and gases- CIE iGCSE Chemistry Notes - New Syllabus
Diffusion for iGCSE Chemistry Notes
Core Syllabus
- Describe tests to identify the anions:
(a) carbonate, CO3²–, by reaction with dilute acid and then testing for carbon dioxide gas
(b) chloride, Cl–, bromide, Br–, and iodide, I–, by acidifying with dilute nitric acid then adding aqueous silver nitrate
(c) nitrate, NO3–, reduction with aluminium foil and aqueous sodium hydroxide and then testing for ammonia gas
(d) sulfate, SO4²–, by acidifying with dilute nitric acid and then adding aqueous barium nitrate
(e) sulfite, SO3²–, by reaction with acidified aqueous potassium manganate(VII) - Describe tests using aqueous sodium hydroxide and aqueous ammonia to identify the aqueous cations:
(a) aluminium, Al³+
(b) ammonium, NH4+
(c) calcium, Ca²+
(d) chromium(III), Cr³+
(e) copper(II), Cu²+
(f) iron(II), Fe²+
(g) iron(III), Fe³+
(h) zinc, Zn²+ - Describe tests to identify the gases:
(a) ammonia, NH3, using damp red litmus paper
(b) carbon dioxide, CO2, using limewater
(c) chlorine, Cl2, using damp litmus paper
(d) hydrogen, H2, using a lighted splint
(e) oxygen, O2, using a glowing splint
(f) sulfur dioxide, SO2, using acidified aqueous potassium manganate(VII) - Describe the use of a flame test to identify the cations:
(a) lithium, Li+
(b) sodium, Na+
(c) potassium, K+
(d) calcium, Ca²+
(e) barium, Ba²+
(f) copper(II), Cu²+
Tests to identify anions
Tests to identify anions
Different anions can be identified using characteristic chemical tests that produce distinctive observations such as effervescence, precipitates, colour changes, or gas evolution. Below are the required tests:
(a) Carbonate ions, \( \text{CO}_3^{2-} \)
Test: Add dilute acid (such as hydrochloric acid or nitric acid) to the solid or solution suspected to contain carbonate ions.
Observation: Effervescence occurs as carbon dioxide gas is released.
Confirmatory step: Pass the gas through limewater. If it turns milky, carbon dioxide is present, confirming carbonate ions.
Equation:
\( \text{CaCO}_3 + 2\text{HCl} \rightarrow \text{CaCl}_2 + \text{H}_2\text{O} + \text{CO}_2 \) \( \text{CO}_2 + \text{Ca(OH)}_2 \rightarrow \text{CaCO}_3 + \text{H}_2\text{O} \)
(b) Halide ions: chloride (\( \text{Cl}^- \)), bromide (\( \text{Br}^- \)), iodide (\( \text{I}^- \))
Test: Acidify the sample with dilute nitric acid (to remove interfering ions such as carbonates or sulfites), then add aqueous silver nitrate.
Observations:
- Chloride ions: white precipitate of silver chloride, \( \text{AgCl} \)
- Bromide ions: cream precipitate of silver bromide, \( \text{AgBr} \)
- Iodide ions: yellow precipitate of silver iodide, \( \text{AgI} \)
Example equation:
\( \text{Ag}^+ + \text{Cl}^- \rightarrow \text{AgCl (s)} \)
(c) Nitrate ions, \( \text{NO}_3^- \)
Test: Add aluminium foil and aqueous sodium hydroxide to the test solution, then warm gently.
Observation: Ammonia gas is released if nitrate ions are present. Ammonia can be detected by its characteristic pungent smell or by turning damp red litmus paper blue. Aluminium reduces nitrate ions to ammonia under alkaline conditions.
Ionic equation:
\( 3\text{NO}_3^- + 8\text{Al} + 5\text{OH}^- + 18\text{H}_2\text{O} \rightarrow 3\text{NH}_3 + 8[\text{Al(OH)}_4]^- \)
(d) Sulfate ions, \( \text{SO}_4^{2-} \)
Test: Acidify the solution with dilute nitric acid, then add aqueous barium nitrate.
Observation: A white precipitate of barium sulfate forms if sulfate ions are present.
Equation:
\( \text{Ba}^{2+} + \text{SO}_4^{2-} \rightarrow \text{BaSO}_4 (s) \)
The acidification step removes carbonate or sulfite ions that might also give precipitates.
(e) Sulfite ions, \( \text{SO}_3^{2-} \)
Test: Add acidified aqueous potassium manganate(VII).
Observation: The purple solution of manganate(VII) is decolourised to colourless if sulfite ions are present. Sulfite ions act as a reducing agent, reducing manganate(VII) ions to manganese(II).
Example ionic equation:
\( 2\text{MnO}_4^- + 5\text{SO}_3^{2-} + 6\text{H}^+ \rightarrow 2\text{Mn}^{2+} + 5\text{SO}_4^{2-} + 3\text{H}_2\text{O} \)
Example
A student adds dilute nitric acid followed by aqueous silver nitrate to an unknown solution. A cream precipitate is observed. Identify the anion present.
▶️Answer/Explanation
A cream precipitate with silver nitrate indicates the presence of bromide ions, \( \text{Br}^- \). Reaction: \( \text{Ag}^+ + \text{Br}^- \rightarrow \text{AgBr (s)} \).
Tests to identify aqueous cations
Tests to identify aqueous cations
- Aqueous cations can be identified using characteristic reactions with sodium hydroxide (NaOH) solution and aqueous ammonia (NH3).
- These reagents cause precipitates to form, which may or may not dissolve in excess.
- Some ions also release gases on heating.
(a) Aluminium ions, \( \text{Al}^{3+} \)
With NaOH: White precipitate of aluminium hydroxide forms, soluble in excess to give a colourless solution.
With NH3: White precipitate forms, insoluble in excess.
Equation:
\( \text{Al}^{3+} + 3\text{OH}^- \rightarrow \text{Al(OH)}_3 (s) \)
(b) Ammonium ions, \( \text{NH}_4^+ \)
With NaOH: On warming, ammonia gas is released. Ammonia has a pungent smell and turns damp red litmus paper blue.
With NH3: No visible reaction (since it already contains ammonia).
Equation:
\( \text{NH}_4^+ + \text{OH}^- \rightarrow \text{NH}_3 (g) + \text{H}_2\text{O} \)
(c) Calcium ions, \( \text{Ca}^{2+} \)
With NaOH: White precipitate forms, insoluble in excess.
With NH3: No precipitate or only very slight.
Equation:
\( \text{Ca}^{2+} + 2\text{OH}^- \rightarrow \text{Ca(OH)}_2 (s) \)
(d) Chromium(III) ions, \( \text{Cr}^{3+} \)
With NaOH: Grey-green precipitate forms, soluble in excess to form a green solution.
With NH3: Grey-green precipitate forms, insoluble in excess.
Equation:
\( \text{Cr}^{3+} + 3\text{OH}^- \rightarrow \text{Cr(OH)}_3 (s) \)
(e) Copper(II) ions, \( \text{Cu}^{2+} \)
With NaOH: Light blue precipitate forms, insoluble in excess.
With NH3: Light blue precipitate forms, soluble in excess to form a deep blue solution (due to the tetraammine copper(II) complex).
Equations:
\( \text{Cu}^{2+} + 2\text{OH}^- \rightarrow \text{Cu(OH)}_2 (s) \)
\( \text{Cu(OH)}_2 + 4\text{NH}_3 \rightarrow [\text{Cu(NH}_3)_4]^{2+} + 2\text{OH}^- \)
(f) Iron(II) ions, \( \text{Fe}^{2+} \)
With NaOH: Green precipitate forms, slowly turns brown on standing (due to oxidation to iron(III)).
With NH3: Green precipitate forms, insoluble in excess.
Equation:
\( \text{Fe}^{2+} + 2\text{OH}^- \rightarrow \text{Fe(OH)}_2 (s) \)
(g) Iron(III) ions, \( \text{Fe}^{3+} \)
With NaOH: Red-brown precipitate forms, insoluble in excess.
With NH3: Red-brown precipitate forms, insoluble in excess.
Equation:
\( \text{Fe}^{3+} + 3\text{OH}^- \rightarrow \text{Fe(OH)}_3 (s) \)
(h) Zinc ions, \( \text{Zn}^{2+} \)
With NaOH: White precipitate forms, soluble in excess to give a colourless solution.
With NH3: White precipitate forms, soluble in excess to form a colourless solution (different from Al\(^{3+}\), which is insoluble in excess NH3).
Equation:
\( \text{Zn}^{2+} + 2\text{OH}^- \rightarrow \text{Zn(OH)}_2 (s) \)
Example
A student adds aqueous sodium hydroxide to an unknown solution. A white precipitate forms, which dissolves in excess NaOH to give a colourless solution. The same test with aqueous ammonia gives a white precipitate that remains insoluble in excess. Identify the cation present.
▶️Answer/Explanation
These results match the behaviour of aluminium ions, \( \text{Al}^{3+} \). With NaOH: \( \text{Al(OH)}_3 \) dissolves in excess to give a colourless solution. With NH3: \( \text{Al(OH)}_3 \) precipitate remains insoluble.
Tests for gases
Tests for gases
Several gases can be identified by their characteristic tests, observations, and reactions. These simple laboratory tests allow chemists to distinguish between different gases quickly.
(a) Ammonia, \( \text{NH}_3 \)
Test: Place damp red litmus paper in the gas.
Observation: Turns blue (because ammonia is alkaline).
Additional test: Produces dense white fumes with hydrogen chloride gas due to the formation of ammonium chloride.
Equation:
\( \text{NH}_3 + \text{HCl} \rightarrow \text{NH}_4\text{Cl} \)
(b) Carbon dioxide, \( \text{CO}_2 \)
Test: Bubble the gas through limewater (aqueous calcium hydroxide).
Observation: Limewater turns milky (white precipitate of calcium carbonate forms).
Equation:
\( \text{Ca(OH)}_2 + \text{CO}_2 \rightarrow \text{CaCO}_3 (s) + \text{H}_2\text{O} \)
(c) Chlorine, \( \text{Cl}_2 \)
Test: Place damp blue litmus paper in the gas.
Observation: Turns red (acidic), then bleaches white. Chlorine dissolves in water to form hydrochloric acid and hypochlorous acid, which bleaches dyes.
(d) Hydrogen, \( \text{H}_2 \)
Test: Insert a lit splint into the gas.
Observation: Gas burns with a squeaky pop sound.
Equation:
\( 2\text{H}_2 + \text{O}_2 \rightarrow 2\text{H}_2\text{O} \)
(e) Oxygen, \( \text{O}_2 \)
Test: Insert a glowing splint into the gas.
Observation: Splint relights brightly. Oxygen supports combustion strongly.
(f) Sulfur dioxide, \( \text{SO}_2 \)
Test: Bubble the gas through acidified potassium manganate(VII) solution (purple).
Observation: Solution decolourises (turns colourless).
Equation:
\( 2\text{H}^+ + 2\text{MnO}_4^- + 5\text{SO}_2 \rightarrow 2\text{Mn}^{2+} + 5\text{SO}_4^{2-} + 2\text{H}_2\text{O} \)
Example
A student collects a colourless gas from a reaction and bubbles it through limewater. The limewater turns milky. Identify the gas and write the equation for the reaction taking place in limewater.
▶️Answer/Explanation
The gas is carbon dioxide, \( \text{CO}_2 \). Limewater turns milky because of the formation of insoluble calcium carbonate:
\( \text{Ca(OH)}_2 + \text{CO}_2 \rightarrow \text{CaCO}_3 (s) + \text{H}_2\text{O} \).
Flame tests for cations
Flame tests for cations
- Flame tests are used to identify certain metal cations by the characteristic colour they produce when heated in a flame.
- The test works because the heat of the flame excites electrons in the metal ion to higher energy levels.
- When the electrons fall back to their original levels, they release energy as visible light of a specific colour.
Method:
- Dip a clean platinum or nichrome wire loop into concentrated hydrochloric acid to clean it.
- Dip the loop into a solid sample or a solution of the salt to be tested.
- Hold the loop in a non-luminous (blue) Bunsen flame and observe the flame colour.
Flame colours of common cations:
- Lithium, \( \text{Li}^+ \) → red flame (crimson red)
- Sodium, \( \text{Na}^+ \) → yellow flame (bright yellow)
- Potassium, \( \text{K}^+ \) → lilac flame (pale violet)
- Calcium, \( \text{Ca}^{2+} \) → orange-red flame (brick red)
- Barium, \( \text{Ba}^{2+} \) → green flame (apple green)
- Copper(II), \( \text{Cu}^{2+} \) → blue-green flame
Explanation of flame colours:
- Each metal ion has electrons that can absorb heat energy and move to a higher energy level.
- When they return to their lower energy level, the difference in energy is released as visible light.
- The wavelength (colour) depends on the energy gap, which is unique to each ion.
Example
A student performs a flame test on an unknown salt. The flame colour observed is bright yellow. Identify the cation present in the salt.
▶️Answer/Explanation
The bright yellow flame indicates the presence of sodium ions, \( \text{Na}^+ \).
Therefore, the salt contains sodium.