Chemical analysis

Welcome to Chemical Analysis

Think of chemical analysis as chemistry detective work. When scientists find an unknown substance — whether at a crime scene, in a polluted river, or in a batch of medicine — they need reliable tests to identify exactly what is in it. In this chapter of Unit 1, you will learn how to determine whether a substance is pure, how to separate mixtures, and how to carry out the standard flame and test-tube tests used to identify unknown positive and negative ions.

Don't worry if all the chemical tests and colours seem overwhelming at first! We will break them down into easy-to-remember patterns and mnemonics.

1. Purity, Formulations, and Chromatography

Pure Substances vs. Impure Mixtures

In everyday life, a label on a bottle might say "pure orange juice," but in chemistry, pure has a very strict meaning:

• A pure substance consists of only a single element or a single compound that is not mixed with any other substance.
• Melting and boiling points: A pure substance melts and boils at specific, sharp, fixed temperatures (for example, pure water melts at exactly \(0^\circ\text{C}\) and boils at exactly \(100^\circ\text{C}\)).
• The effect of impurities: Adding an impurity lowers the melting point and causes the substance to melt over a broad range of temperatures. Impurities also raise the boiling point.

Formulations

Not all mixtures are accidental. Many products around us are carefully designed:

• A formulation is a mixture designed as a useful product, manufactured with carefully defined proportions of specific chemicals to ensure it has the exact properties needed.
• Everyday examples: Paints, medicines, foods, cleaning agents, and fertilisers.
• Analogy: Baking a cake requires an exact recipe so that the cake rises and tastes right. A formulation is a chemist's precise recipe!

Potable Water

Water that is safe to drink is called potable water. Sea water contains high amounts of dissolved salts and is not potable. Pure water can be separated from dissolved salts in seawater by distillation.

Paper Chromatography

Paper chromatography is an analytical technique used to separate and identify dissolved substances (such as dyes in inks or food colourings) based on their different solubilities.

• Stationary phase: The chromatography paper (it stays still).
• Mobile phase: The solvent (such as water) that moves up the paper.
• How it works: Substances that are more soluble in the mobile phase spend more time in the solvent and travel further up the paper. Substances with a stronger attraction to the stationary phase do not travel as far.

Important Practical Rules for Chromatography:
1. Always draw the baseline/origin line in pencil. Pencil lead (graphite) is insoluble in the solvent and will not run or contaminate the results. Never use pen/ink!
2. The solvent level must start below the pencil baseline. If the solvent is above the line, the samples will dissolve directly into the solvent beaker rather than traveling up the paper.

Calculating the Retention Factor (\(R_f\)):
The \(R_f\) value allows chemists to identify substances by comparing their values to known standards.

\(R_f = \frac{\text{Distance moved by substance}}{\text{Distance moved by solvent front}}\)

• The distance moved by the substance is measured from the pencil baseline to the centre of the spot.
• The distance moved by the solvent front is measured from the pencil baseline to where the solvent stopped.
• Remember: An \(R_f\) value is always a decimal less than or equal to 1 (\(R_f \le 1\)) and has no units.

Key Takeaway: Pure substances have sharp melting/boiling points. Formulations are useful mixtures with set recipes. In chromatography, \(R_f\) is always \(\le 1\) and baseline lines are always drawn in pencil.

2. Identifying Metal Cations (Positive Ions)

Metal ions carry a positive charge and are called cations (remember: ca+ions are paws-itive). We use two main methods to identify them: Flame Tests and Precipitation Tests.

Method 1: Flame Tests (CCEA 1.9.12–1.9.13)

Certain metal ions give off distinctive colours when heated in a flame.

Step-by-Step Procedure:
1. Dip a clean nichrome wire loop into concentrated hydrochloric acid (\(\text{HCl}\)) and hold it in a flame to clean off any contaminants.
2. Dip the clean wire into the solid sample or test solution.
3. Place the wire loop into the edge of a blue / roaring Bunsen burner flame.
4. Observe and record the colour of the flame.

Required CCEA Flame Colours:
• Lithium (\(\text{Li}^+\)): Crimson
• Sodium (\(\text{Na}^+\)): Yellow / Orange
• Potassium (\(\text{K}^+\)): Lilac
• Calcium (\(\text{Ca}^{2+}\)): Brick red
• Copper(II) (\(\text{Cu}^{2+}\)): Blue-green / Green-blue

Memory Trick: Potassium = Purple/Lilac; Sodium = Sunny Yellow/Orange; Calcium = Clay Brick red.

Method 2: Sodium Hydroxide (\(\text{NaOH}\)) and Ammonia (\(\text{NH}_3\)) Tests

When dilute sodium hydroxide solution or dilute aqueous ammonia solution is added to solutions containing metal cations, insoluble metal hydroxides form as solid precipitates. We observe the colour of the precipitate when adding a few drops, and then check what happens when we add the reagent in excess.

Group A: The Coloured Precipitates

• Copper(II) (\(\text{Cu}^{2+}\)):
- With \(\text{NaOH}\): Blue precipitate of \(\text{Cu(OH)}_2\), which is insoluble in excess.
- With \(\text{NH}_3\): Blue precipitate, which dissolves in excess to form a deep blue solution.

• Iron(II) (\(\text{Fe}^{2+}\)):
- With \(\text{NaOH}\): Green precipitate of \(\text{Fe(OH)}_2\), insoluble in excess.
- With \(\text{NH}_3\): Green precipitate, insoluble in excess.
- Note: The green precipitate turns brown at the surface over time because iron(II) oxidises to iron(III) when exposed to air.

• Iron(III) (\(\text{Fe}^{3+}\)):
- With \(\text{NaOH}\): Brown / Red-brown precipitate of \(\text{Fe(OH)}_3\), insoluble in excess.
- With \(\text{NH}_3\): Brown / Red-brown precipitate, insoluble in excess.

Group B: The White Precipitates (\(\text{Mg}^{2+}\), \(\text{Al}^{3+}\), \(\text{Zn}^{2+}\))

All three of these ions form a white precipitate when a few drops of \(\text{NaOH}\) or \(\text{NH}_3\) are added. How do we tell them apart? Follow this systematic order:

1. Add excess \(\text{NaOH}\):
• Magnesium (\(\text{Mg}^{2+}\)): White precipitate \([\text{Mg(OH)}_2]\) remains insoluble.
• Aluminium (\(\text{Al}^{3+}\)): White precipitate \([\text{Al(OH)}_3]\) dissolves to form a colourless solution.
• Zinc (\(\text{Zn}^{2+}\)): White precipitate \([\text{Zn(OH)}_2]\) dissolves to form a colourless solution.

2. How to distinguish \(\text{Al}^{3+}\) from \(\text{Zn}^{2+}\):
Test a fresh sample with aqueous ammonia (\(\text{NH}_3\)) in excess:
• Aluminium (\(\text{Al}^{3+}\)): White precipitate is insoluble in excess \(\text{NH}_3\).
• Zinc (\(\text{Zn}^{2+}\)): White precipitate dissolves in excess \(\text{NH}_3\).

Ionic Equations for Cation Precipitation

In the CCEA exam, you must write balanced ionic equations with state symbols:

• \(\text{Cu}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Cu(OH)}_2\text{(s)}\)
• \(\text{Fe}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_2\text{(s)}\)
• \(\text{Fe}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Fe(OH)}_3\text{(s)}\)
• \(\text{Mg}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Mg(OH)}_2\text{(s)}\)
• \(\text{Al}^{3+}\text{(aq)} + 3\text{OH}^-\text{(aq)} \rightarrow \text{Al(OH)}_3\text{(s)}\)
• \(\text{Zn}^{2+}\text{(aq)} + 2\text{OH}^-\text{(aq)} \rightarrow \text{Zn(OH)}_2\text{(s)}\)

Key Takeaway: \(\text{Cu}^{2+}\) is blue, \(\text{Fe}^{2+}\) is green, \(\text{Fe}^{3+}\) is brown. For white precipitates, only \(\text{Al}^{3+}\) and \(\text{Zn}^{2+}\) dissolve in excess \(\text{NaOH}\); only \(\text{Zn}^{2+}\) dissolves in excess \(\text{NH}_3\).

3. Identifying Anions (Negative Ions)

Negative ions are called anions (remember: a negative ion). Here are the chemical tests for carbonates, sulfates, and halides.

1. Carbonate Ion (\(\text{CO}_3^{2-}\))

• Reagent: Add dilute hydrochloric acid (\(\text{HCl}\)).
• Observation: Effervescence (fizzing/bubbling).
• Confirmatory test: Bubble the gas produced through limewater. The limewater turns milky / cloudy, confirming the presence of carbon dioxide gas (\(\text{CO}_2\)).

2. Sulfate Ion (\(\text{SO}_4^{2-}\))

• Reagents: Add dilute hydrochloric acid (\(\text{HCl}\)), followed by barium chloride solution (\(\text{BaCl}_2\)).
• Observation: A white precipitate of barium sulfate (\(\text{BaSO}_4\)) forms.
• Ionic Equation:
\(\text{Ba}^{2+}\text{(aq)} + \text{SO}_4^{2-}\text{(aq)} \rightarrow \text{BaSO}_4\text{(s)}\)

3. Halide Ions (\(\text{Cl}^-\), \(\text{Br}^-\), \(\text{I}^-\))

• Reagents: Add dilute nitric acid (\(\text{HNO}_3\)), followed by silver nitrate solution (\(\text{AgNO}_3\)).
• Observations:
- Chloride (\(\text{Cl}^-\)): White precipitate of silver chloride (\(\text{AgCl}\))
- Bromide (\(\text{Br}^-\)): Cream precipitate of silver bromide (\(\text{AgBr}\))
- Iodide (\(\text{I}^-\)): Yellow precipitate of silver iodide (\(\text{AgI}\))

Memory Trick: Milk, Cream, Butter \(\rightarrow\) White (\(\text{Cl}^-\)), Cream (\(\text{Br}^-\)), Yellow (\(\text{I}^-\)).

Ionic Equations for Halides:
• \(\text{Ag}^+\text{(aq)} + \text{Cl}^-\text{(aq)} \rightarrow \text{AgCl(s)}\)
• \(\text{Ag}^+\text{(aq)} + \text{Br}^-\text{(aq)} \rightarrow \text{AgBr(s)}\)
• \(\text{Ag}^+\text{(aq)} + \text{I}^-\text{(aq)} \rightarrow \text{AgI(s)}\)

Key Takeaway: Carbonates fizz with acid; sulfates give a white ppt with \(\text{BaCl}_2\); halides give white, cream, or yellow precipitates with \(\text{AgNO}_3\).

4. Prescribed Practical C4 & Examiner Tips

Why Do We Acidify Before Adding \(\text{AgNO}_3\) or \(\text{BaCl}_2\)?

• The Reason: Adding dilute acid destroys and removes any dissolved carbonate ions (\(\text{CO}_3^{2-}\)) that might be present as impurities.
• If carbonate ions were present, they would react to form silver carbonate (\(\text{Ag}_2\text{CO}_3\)) or barium carbonate (\(\text{BaCO}_3\)), both of which are white precipitates. This would give a false-positive result!
• Choosing the correct acid:
- For the halide test: You must use dilute nitric acid (\(\text{HNO}_3\)). Never use hydrochloric acid (\(\text{HCl}\)), because \(\text{HCl}\) contains chloride ions (\(\text{Cl}^-\)), which would react with \(\text{AgNO}_3\) and give a false white precipitate of \(\text{AgCl}\)!
- For the sulfate test: You use dilute hydrochloric acid (\(\text{HCl}\)).

Common Exam Pitfalls to Avoid

• Measuring \(R_f\) distances: Always measure from the pencil baseline, never from the very bottom edge of the paper.
• Flame test wire: If the nichrome wire is not cleaned in concentrated acid between tests, residual sodium will create an intense yellow/orange flame that masks every other colour.
• Describing precipitate changes: Be precise with your words. Say "the white precipitate dissolves in excess sodium hydroxide to form a colourless solution" rather than just "it disappears" or "it turns clear".

Quick Summary Checklist

Before sitting your exam, make sure you can:
1. State the definitions of a pure substance and a formulation.
2. Calculate \(R_f\) values and explain why the baseline must be in pencil.
3. Recall all five flame test colours (\(\text{Li}^+\), \(\text{Na}^+\), \(\text{K}^+\), \(\text{Ca}^{2+}\), \(\text{Cu}^{2+}\)).
4. Distinguish between \(\text{Mg}^{2+}\), \(\text{Al}^{3+}\), and \(\text{Zn}^{2+}\) using excess \(\text{NaOH}\) and excess \(\text{NH}_3\).
5. State the reagents and results for \(\text{CO}_3^{2-}\), \(\text{SO}_4^{2-}\), \(\text{Cl}^-\), \(\text{Br}^-\), and \(\text{I}^-\).
6. Write balanced ionic equations including state symbols for all precipitation reactions.