Welcome to Chemical Analysis!
Ever wondered how forensic scientists solve crimes from tiny drops of ink, or how water treatment plants ensure our drinking water is safe? That is what chemical analysis is all about! In this chapter for CCEA GCSE Double Award Science (Unit C1), we will learn how to describe mixtures, separate different substances, identify common gases, and test for pure water. Don't worry if science has felt tough in the past—we will take this step-by-step with clear definitions and helpful memory tricks.
1. Pure Substances, Mixtures, and Formulations
What Does "Pure" Really Mean?
In everyday life, a carton of orange juice might say "100% pure orange juice", but in chemistry, that juice is a complex mixture of water, sugars, acids, and vitamins!
Here are the official definitions you need for your exam:
• Pure Substance: A single element or a single compound that is not mixed with any other substance. A pure substance melts and boils at specific, sharp, fixed temperatures (for example, pure water boils at exactly \(100\,^\circ\text{C}\) and melts at \(0\,^\circ\text{C}\)).
• Mixture: Two or more substances that are not chemically combined together. The substances in a mixture keep their own individual properties and can be separated easily by physical methods.
What is a Formulation?
A formulation is a very special kind of mixture. It has been designed as a useful product by mixing specific components in carefully measured quantities to ensure it has the exact properties needed.
Everyday examples of formulations: fuels, paints, medicines, and alloys.
Key Solution Vocabulary
• Solute: The solid or substance that dissolves in a liquid (e.g., salt).
• Solvent: The liquid in which the solute dissolves (e.g., water).
• Solution: The uniform mixture formed when a solute dissolves in a solvent (e.g., salt water).
• Soluble: Capable of dissolving in a particular solvent.
• Insoluble: Does not dissolve in a particular solvent.
• Miscible: Liquids that mix completely together to form a single layer (e.g., ethanol and water).
• Immiscible: Liquids that do not mix and separate into distinct layers (e.g., oil and water; separated using a separating funnel).
Key Takeaway: Pure substances have sharp melting and boiling points. Formulations are mixtures carefully measured for a specific purpose.
2. Separation Techniques
Because the parts of a mixture are not chemically joined, we can use physical separation methods to isolate them. Let's look at the main techniques you need to know:
A. Filtration
• Purpose: Separates an insoluble solid from a liquid or solution (e.g., sand from water).
• Apparatus: Filter funnel, filter paper, conical flask or beaker.
• Key terms to remember:
- Residue: The solid left behind on the filter paper.
- Filtrate: The liquid that passes through the filter paper into the flask.
B. Evaporation / Crystallisation
• Purpose: Separates a dissolved soluble solid (solute) from a solvent to obtain solid crystals (e.g., getting salt crystals from salt water).
• Apparatus: Evaporating basin, gauze, tripod, Bunsen burner, heatproof mat.
• Method: The solution is heated so the liquid solvent evaporates away, leaving the solid crystals behind.
C. Simple Distillation
• Purpose: Separates and collects a liquid solvent from a solution (e.g., collecting pure water from salt water).
• How it works:
1. The solution is heated in a flask until the solvent boils and turns into vapour.
2. Anti-bumping granules are added to the flask to promote smooth boiling.
3. The hot vapour rises and travels into a Liebig condenser.
4. The condenser is cooled by cold water flowing through an outer jacket. Crucial exam rule: Cold water must enter at the bottom and leave at the top so the jacket fills completely.
5. The vapour condenses back into a pure liquid, called the distillate, which is collected in a beaker or flask.
D. Fractional Distillation
• Purpose: Separates miscible liquids that have different boiling points (e.g., separating ethanol from water).
• Apparatus: Uses a fractionating column filled with glass beads, attached below the condenser, along with a thermometer.
• Why glass beads? They provide a large surface area for repeated condensation and evaporation. The liquid with the lower boiling point evaporates and reaches the top first, while the liquid with the higher boiling point condenses and drips back down.
Key Takeaway: Use filtration for insoluble solids, evaporation for dissolved solids, simple distillation to save the solvent, and fractional distillation for miscible liquids.
3. Paper Chromatography
Chromatography is used to separate mixtures of soluble coloured substances, such as food dyes or ink pigments.
The Two Phases
• Stationary Phase: The material that does not move — the chromatography paper.
• Mobile Phase: The liquid that moves up the paper — the solvent (such as water or ethanol).
Step-by-Step Practical Method
1. Draw a baseline in pencil about \(1\text{ to }2\text{ cm}\) from the bottom of the chromatography paper.
Why pencil? Pencil lead (graphite) is insoluble and will not dissolve in the solvent or run up the paper.
2. Place small spots of the ink or dye samples onto the pencil baseline.
3. Place the paper into a beaker containing a small depth of solvent. The solvent level must be below the pencil baseline so the spots do not wash directly into the solvent reservoir.
4. The solvent moves up the paper, carrying the substances with it. Different substances travel different distances based on how soluble they are in the mobile phase versus how strongly attracted they are to the stationary paper.
5. Remove the paper before the solvent reaches the top, and immediately mark the solvent front in pencil (the highest point reached by the solvent).
Interpreting the Chromatogram & Calculating \(R_f\) Values
• A pure substance produces only one single spot.
• A mixture produces two or more spots.
• We can calculate the retention factor (\(R_f\) value) for each spot using this formula:
\[R_f = \frac{\text{distance moved by spot}}{\text{distance moved by solvent front}}\]
Important \(R_f\) Value Rules:
• Always measure distances starting from the pencil baseline.
• Because the spot can never travel further than the solvent front, an \(R_f\) value is always less than or equal to 1.0 (\(R_f \le 1.0\)).
• \(R_f\) values have no units (they are dimensionless numbers).
• If two substances have the exact same \(R_f\) value in the same solvent, they are likely the same compound.
Key Takeaway: Always draw the baseline in pencil, keep the solvent below the baseline, and remember that \(R_f\) is spot distance divided by solvent front distance (\(\le 1.0\)).
4. Qualitative Chemical Tests
In the lab, you must be able to identify specific gases and verify whether a sample of liquid is water.
Gas Identification Tests
Make sure you memorise both the test method and the positive result for each gas:
• Hydrogen (\(\text{H}_2\)):
- Test: Hold a lit splint near the mouth of the test tube.
- Result: Burns with a 'squeaky pop' sound.
• Oxygen (\(\text{O}_2\)):
- Test: Insert a glowing splint into the test tube.
- Result: Relights the glowing splint.
• Carbon Dioxide (\(\text{CO}_2\)):
- Test: Bubble the gas through limewater (calcium hydroxide solution).
- Result: Limewater turns milky / cloudy (due to the formation of a white precipitate, \(\text{CaCO}_3\)).
• Chlorine (\(\text{Cl}_2\)):
- Test: Hold damp blue litmus paper in the gas.
- Result: Turns red, then bleaches white.
Tests for Water
Examiners love testing the difference between a chemical test for the presence of water and a physical test for pure water!
1. Chemical Test for Water (Presence Test):
Add the liquid to anhydrous copper(II) sulfate.
- Result: Colour changes from white to blue.
- Equation: \[\text{CuSO}_4\text{ (s)} + 5\text{H}_2\text{O (l)} \rightarrow \text{CuSO}_4\cdot5\text{H}_2\text{O (s)}\]
Note: This test proves water is present, but it does not prove the water is pure!
2. Physical Test for Pure Water (Purity Test):
- Pure water has a sharp boiling point of \(100\,^\circ\text{C}\) and a sharp melting point of \(0\,^\circ\text{C}\) (at \(1\text{ atm}\) pressure). If water contains impurities like salt, it will boil above \(100\,^\circ\text{C}\) and freeze below \(0\,^\circ\text{C}\).
Key Takeaway: Hydrogen pops, oxygen relights, carbon dioxide turns limewater milky, and chlorine bleaches damp litmus. Anhydrous copper(II) sulfate turns white to blue with water.
Common Exam Pitfalls & How to Avoid Them
• Using pen on chromatography paper: Ink contains soluble dyes that will run and ruin your results. Always use a pencil.
• Too much solvent in chromatography: If the solvent level is above the baseline spots, the spots will wash off into the beaker.
• Inverting \(R_f\) calculations: If your calculated \(R_f\) value is greater than 1.0, you divided upside down! It is always \(\frac{\text{spot distance}}{\text{solvent distance}}\).
• Condenser water direction: Water goes in at the bottom (coldest point) and out at the top. Drawing arrows the wrong way loses easy marks.
• Chemical purity vs. everyday purity: Never say tap water or bottled mineral water is "chemically pure"—they contain dissolved minerals and are mixtures.
Quick Review Summary
• Pure substances: Single element/compound with sharp melting/boiling points.
• Formulation: A mixture designed for a specific use with exact quantities.
• Filtration: Insoluble solid + liquid \(\rightarrow\) leaves residue on paper, collects filtrate.
• Distillation: Simple distillation collects the solvent; fractional distillation separates miscible liquids.
• Chromatography: \(R_f = \frac{\text{distance moved by spot}}{\text{distance moved by solvent front}}\) (always \(\le 1.0\)).
• Gas tests: Lit splint for \(\text{H}_2\) (pop), glowing splint for \(\text{O}_2\) (relights), limewater for \(\text{CO}_2\) (milky), damp blue litmus for \(\text{Cl}_2\) (red then bleaches white).
• Water tests: Anhydrous copper(II) sulfate turns white \(\rightarrow\) blue; pure water boils at exactly \(100\,^\circ\text{C}\) and melts at \(0\,^\circ\text{C}\).