Chemical analysis

Explanation

• In chemistry, a pure substance contains one element or one compound and is not mixed with another substance.
• Test purity by measuring a melting point or boiling point and comparing it with reliable reference data.
• For example, a sample that melts sharply at the accepted melting point is consistent with a pure substance; a mixture usually changes the value and melts over a range.
• A common error is to use the everyday meaning of pure, such as natural or unadulterated, instead of the chemical meaning.

Explanation

• A formulation is a mixture designed to be a useful product, with every component included for a particular purpose.
• Make a formulation by measuring and mixing its components carefully so the final product has the required properties.
• For example, a paint may contain a pigment for colour, a solvent to control flow and a binder that leaves a solid coating.
• A common error is to call every mixture a formulation; a formulation must have a designed use and controlled composition.

Explanation

• Chromatography separates a mixture because its substances distribute differently between a stationary phase and a mobile phase; a mixture may give several spots, while a pure compound gives one spot in every solvent.
• For paper chromatography, draw a pencil origin line, add small sample spots, keep the solvent below the line, allow the solvent to rise, then mark the solvent front.
• Calculate $R_f=\frac{\text{distance moved by the centre of the spot}}{\text{distance moved by the solvent front}}$; for example, $3.6/6.0=0.60$.
• A common error is to measure from the paper edge or spot boundary instead of from the origin to the centre of the spot.
• Separation also depends on solubility in the mobile phase and attraction to the stationary phase, so changing the solvent can change the pattern.

Explanation

• Hydrogen is identified by the characteristic pop produced when it burns rapidly. Hold a burning splint at the open end of the test tube containing the collected gas.
• For example, a gas that gives a pop with a lit splint has a positive hydrogen-test result.
• A common error is to insert a glowing splint; that is the test for oxygen, not hydrogen.
• Only a small collected sample should be tested because hydrogen burns rapidly in oxygen.
• The splint stays outside the test tube opening.

Explanation

• Oxygen supports combustion and is identified when a glowing splint relights. Insert a glowing wooden splint into the test tube containing the gas.
• For example, a splint with no flame that bursts back into flame gives a positive oxygen result.
• A common error is to use a burning splint and look for a pop, which tests for hydrogen.
• The test uses oxygen's ability to support combustion and should be performed on a collected gas sample.
• Relighting is the required positive observation.

Explanation

• Carbon dioxide is identified because it forms a fine white precipitate in limewater, making the liquid look milky or cloudy.
• Bubble the gas through, or shake it with, aqueous calcium hydroxide solution.
• For example, clear limewater becoming cloudy is a positive result for carbon dioxide under this test.
• A common error is to report that limewater becomes colourless; the required observation is that it turns milky or cloudy.
• The cloudiness is caused by insoluble calcium carbonate, so ‘white precipitate forms’ is also a valid observation.

Explanation

• Chlorine gas bleaches damp litmus paper white. Expose damp litmus paper to the gas while using appropriate small-scale safety precautions because chlorine is toxic.
• For example, litmus losing all its colour and becoming white is the positive chlorine result.
• A common error is to use dry litmus paper; moisture is required for the specified test.
• A positive result is bleaching rather than a permanent acid-base indicator colour, so the final white appearance is essential.
• The gas must not be inhaled.

Explanation

• Flame tests identify some metal ions: Li⁺ crimson, Na⁺ yellow, K⁺ lilac, Ca²⁺ orange-red and Cu²⁺ green.
• Place a small amount of a clean sample into a non-luminous flame and compare the observed colour with the known colours.
• For example, a lilac flame indicates potassium ions, whereas an orange-red flame indicates calcium ions.
• A common error is to treat a flame test as reliable for every mixture; a strong colour, especially sodium yellow, can mask another ion.
• Contamination must be controlled by using clean apparatus because traces of sodium can produce an intense yellow flame.

Explanation

• With sodium hydroxide, Cu²⁺ gives a blue precipitate, Fe²⁺ green and Fe³⁺ brown.
• Add sodium hydroxide solution dropwise, record any precipitate, then add excess when distinguishing aluminium from other white precipitates.
• Al³⁺, Ca²⁺ and Mg²⁺ form white hydroxide precipitates, but only aluminium hydroxide dissolves in excess sodium hydroxide.
• A common error is to claim every white precipitate dissolves in excess; calcium and magnesium hydroxides remain.
• The coloured precipitates identify copper(II), iron(II) and iron(III) directly, whereas the three white precipitates require careful discrimination.

Explanation

• Carbonate ions are tested by adding a dilute acid; a carbonate reacts to release carbon dioxide gas.
• Pass the gas produced into limewater to confirm that it is carbon dioxide.
• For example, effervescence followed by limewater turning cloudy is a positive sequence for carbonate ions.
• A common error is to identify a carbonate from fizzing alone; the gas should be confirmed with limewater.
• The confirmatory gas test matters because effervescence alone can be produced by reactions involving gases other than carbon dioxide.

Explanation

• Halide ions are tested by acidifying the solution with dilute nitric acid and then adding silver nitrate solution.
• Use clean samples and record the precipitate colour: chloride white, bromide cream and iodide yellow.
• For example, a cream silver-halide precipitate identifies bromide ions under the specified test conditions.
• A common error is to swap the cream and yellow results: silver bromide is cream, while silver iodide is yellow.
• Acidification removes interfering ions, while nitric acid is chosen because nitrate ions do not form a precipitate with silver ions.

Explanation

• Sulfate ions form a white precipitate when barium chloride solution is added under acidic conditions.
• Acidify the sample with dilute hydrochloric acid, then add barium chloride solution and observe any precipitate.
• For example, a white barium sulfate precipitate is the positive result for sulfate ions.
• A common error is to report only that the mixture turns white without naming the formation of a precipitate.
• The acidic conditions remove interfering carbonate ions before barium ions are added to test for sulfate.

Explanation

• Instrumental methods detect and identify elements or compounds using measured signals rather than only visible chemical-test observations.
• Choose an instrumental method when rapid analysis, high sensitivity or accurate measurement is important.
• For example, a sensitive instrument can detect a component at a concentration too low to give a clear precipitate or colour change.
• A common error is to list 'easy' as a specification advantage; the required comparison is that instrumental methods are accurate, sensitive and rapid.
• Instrumental methods generate signals that can be compared with reference or calibration data to identify or measure substances.

Explanation

• In flame emission spectroscopy, a solution sample enters a flame and the emitted light passes through a spectroscope.
• Identify metal ions by comparing the positions of lines in the sample spectrum with a reference set recorded in the same form.
• Line intensity can be compared with calibration data to measure concentration; for example, an intensity halfway between two standards gives an intermediate concentration when the calibration is linear.
• A common error is to use line brightness to identify the ion; line position identifies the ion, while intensity is used for concentration.
• A sample can contain several metal ions because each element contributes its characteristic set of emission lines.