Your revision route
What you’ll learn
- Use oxygen gain and loss to identify oxidation and reduction.
- Recall the reactivity series and infer metal order from water, acid and displacement evidence.
- Choose carbon reduction or electrolysis to extract a metal from its compound.
- Higher Tier: identify oxidation and reduction from electron transfer and write balanced half equations.
- Predict products of acids with metals, bases, alkalis and carbonates, and name the salt formed.
- Prepare a pure, dry soluble salt from an insoluble base and dilute acid.
- Use pH, hydrogen ions and hydroxide ions to explain acidity, alkalinity and neutralisation.
- Distinguish strong from weak acids and concentration from strength at Higher Tier.
- Predict products from molten and aqueous electrolysis and explain why ions must be mobile.
- Carry out the soluble-salt and aqueous-electrolysis practicals, plus titration for separate Chemistry.
Build the big picture
Key ideas
Oxidation and reduction pass oxygen between substances
When magnesium burns, brilliant white light announces a quieter accounting fact: magnesium has gained oxygen.
- Metals react with oxygen to form metal oxides; the vigour depends on the metal's reactivity.
- Oxidation is gain of oxygen, while reduction is loss of oxygen.
- In 2Mg + O₂ → 2MgO, magnesium gains oxygen and is oxidised.
- In carbon reduction, a metal oxide loses oxygen while carbon gains it, so oxidation and reduction occur together.
The punchline: Follow the oxygen: gained means oxidised; lost means reduced.
The reactivity series is a ladder of electron loss
A higher metal gives up electrons more readily, so it reacts more vigorously and can evict a metal below it from a compound.
- Recall: potassium, sodium, lithium, calcium, magnesium, carbon, zinc, iron, hydrogen, copper, from more to less reactive.
- At room temperature K, Na, Li and Ca react with water; Mg reacts only very slowly, while Zn, Fe and Cu show no reaction. Steam is outside this scope.
- Metals above hydrogen react with dilute acids to form a salt and hydrogen; copper is below hydrogen and does not.
- A more reactive metal displaces a less reactive metal from its compound; observations can therefore reveal an unknown order.
The punchline: For prediction, place both substances on the ladder before writing products.
Extraction reverses the chemistry that trapped a metal
Gold may occur native, but most metals arrive locked in compounds and need a reduction route chosen by reactivity.
- Very unreactive metals such as gold can occur as elements; most metals occur as compounds in ores.
- A metal below carbon can be extracted from its oxide by heating with carbon, which removes oxygen and reduces the oxide.
- A metal above carbon is too reactive for carbon reduction and is extracted by electrolysis of a molten compound.
- The extraction method follows position relative to carbon, not a claim that carbon works for every metal.
The punchline: Below carbon: carbon reduction can work. Above carbon: use electrolysis.
Higher Tier: redox is an electron handover
Electrons cannot simply vanish from one species; every loss is paired with a gain elsewhere in the reaction.
- Oxidation is loss of electrons and reduction is gain of electrons: OIL RIG.
- A metal atom is oxidised when it loses electrons to become a positive ion.
- The species accepting those electrons is reduced, so both processes always occur together.
- Use the chemical equation to name the species oxidised and reduced, rather than merely reciting definitions.
The punchline: Track electrons from donor to acceptor and identify both halves of the redox pair.
Reactivity chooses the extraction route
Start with the metal's position relative to carbon. That comparison decides whether carbon can remove oxygen from its oxide.
- Very unreactiveA metal such as gold may occur native because it does not combine readily with other elements.
- Below carbonHeat the metal oxide with carbon; the oxide loses oxygen and is reduced to the metal.
- Above carbonCarbon cannot reduce the compound, so electrolysis of a molten compound is required.
The acid supplies the salt's surname
Hydrochloric, sulfuric and nitric acids make chloride, sulfate and nitrate salts. For a hydrogen-producing metal reaction here, use dilute hydrochloric or sulfuric acid.
- Magnesium, zinc and iron react with dilute hydrochloric or sulfuric acid to make a salt and hydrogen; more reactive metals react faster under the same conditions.
- Hydrogen gives a squeaky pop with a lighted splint at the mouth of the tube.
- Hydrochloric acid makes chlorides, sulfuric acid makes sulfates and nitric acid makes nitrates in reactions with bases, alkalis or carbonates.
- Higher Tier: the metal loses electrons and is oxidised as it forms positive ions in an acid–metal reaction.
The punchline: For metal + acid → hydrogen, use magnesium, zinc or iron with dilute hydrochloric or sulfuric acid.
Neutralisation has two familiar product patterns
A base removes an acid's bite by consuming hydrogen ions; a carbonate adds a visible fizz of carbon dioxide.
- Acid + base → salt + water; bases include metal oxides and hydroxides, and a soluble hydroxide is an alkali.
- Acid + metal carbonate → salt + water + carbon dioxide, so the gas is carbon dioxide rather than hydrogen.
- The metal in the base or carbonate supplies the first part of the salt name.
- At ion level, neutralisation is H⁺ + OH⁻ → H₂O.
The punchline: Carbonate adds carbon dioxide; it does not switch the gas to hydrogen.
Excess solid makes the final solution cleaner
Add an insoluble base until some refuses to react: that stubborn remainder is evidence that all the acid has been used.
- Warm dilute acid, add an insoluble oxide or carbonate in portions, and continue until the solid is in excess.
- Filter away the unreacted excess, leaving salt solution with no acid remaining.
- Gently evaporate some water, allow crystals to form, then separate and dry them.
- This required practical produces a pure, dry soluble salt; do not evaporate the solution violently to dryness.
The punchline: Excess, filter, concentrate, crystallise, dry—the order protects purity.
pH reports the hydrogen-ion landscape
Universal indicator gives a colour map; a calibrated pH probe replaces the broad colour bands with a number.
- On the 0–14 scale, pH below 7 is acidic, 7 is neutral and above 7 is alkaline.
- Acids supply H⁺ ions in aqueous solution; alkalis supply OH⁻ ions.
- Universal indicator estimates pH from colour, while a pH probe or meter gives a more precise reading.
- During neutralisation, H⁺ and OH⁻ combine in a 1:1 particle ratio to form water.
The punchline: Use the measured pH for classification, then use ions to explain neutralisation.
Make the model move
Interactive checkpoint
Touch the science. Change a state, build a route or test a relationship.
Build the reactivity ladder
Order seven landmarks from most to least reactive
Arrange this representative subset, including the carbon and hydrogen reference points used for extraction and acid reactions.
Choose the first step below.
A higher metal loses electrons more readily. Carbon divides common extraction routes; hydrogen helps predict which metals react with dilute acids.
Free the ions, then list them
How do state and solvent change electrolysis?
Switch between a solid, a melt and two solutions. Watch ion mobility and competition change the electrode products.
Pb²⁺ and Br⁻ are charged, but they are fixed in a lattice and cannot carry current to the electrodes.
1 of 4 states explored
A solid fails because its ions are fixed. Melting frees only the compound's ions; dissolving also introduces H⁺ and OH⁻ from water.
Separate Chemistry: titration catches the end point drop by drop
Near the colour change, one careless squeeze can turn a precise reacting volume into an expensive guess.
- Use a pipette for a fixed acid or alkali volume, a burette for the other solution and one suitable indicator in a conical flask.
- Approach the end point dropwise, record initial and final burette readings, and subtract to obtain the titre.
- Repeat until concordant titres are obtained and calculate a mean without the rough result or clear anomalies.
- Higher Tier separate Chemistry: use reacting volumes and a known concentration to calculate an unknown in mol/dm³ or g/dm³.
The punchline: Read at eye level, add dropwise near the end point and average concordant titres.
Higher Tier: strength is not concentration
A crowded weak acid and a sparse strong acid describe different ideas: amount dissolved and fraction ionised.
- Strong acids such as hydrochloric, sulfuric and nitric acids ionise completely in water; weak acids ionise only partially.
- Ethanoic, citric and carbonic acids are weak acids; at equal concentration, a strong acid has lower pH because it releases more H⁺.
- Each fall of one pH unit means a tenfold increase in hydrogen-ion concentration.
- Concentrated and dilute describe amount per volume; strong and weak describe the extent of ionisation.
The punchline: Ask two separate questions: how much acid is present, and what fraction ionises?
Electrolysis needs ions with room to move
Current can split an ionic compound only when its charged particles are free to travel through the electrolyte.
- Melting an ionic compound or dissolving it in water frees its ions; a solid cannot be electrolysed because its ions are fixed.
- Positive ions move to the negative cathode, while negative ions move to the positive anode.
- Ions are discharged at the electrodes to form elements, decomposing the electrolyte.
- For a simple molten compound, the metal forms at the cathode and the non-metal at the anode.
The punchline: Write the electrode charges first, then send each ion to the opposite charge.
Molten compounds offer one ion choice at each electrode
Remove water from the puzzle and the product rule becomes spare: each ion turns into its element.
- Molten lead bromide contains mobile Pb²⁺ and Br⁻ ions but no competing ions from water.
- Lead forms at the negative cathode; bromine forms at the positive anode.
- Higher Tier explanations use electron gain at the cathode and electron loss at the anode; half equations must balance electrons.
The punchline: For a molten binary compound, map cation to metal and anion to non-metal.
Two electrodes, opposite electron events
Electrode names, charges and electron changes travel as one package. Keep the cathode and anode stories side by side.
- Cathode (−)Positive ions arrive, gain electrons and are reduced. A metal or hydrogen forms.
- Anode (+)Negative ions arrive, lose electrons and are oxidised. A halogen or oxygen forms in aqueous electrolysis.
- Whole circuitThe number of electrons released at the anode equals the number accepted at the cathode.
Aluminium extraction lowers the melting-point bill
Aluminium is too reactive for carbon reduction, so industry melts its oxide—and uses cryolite to avoid an even hotter process.
- Aluminium oxide is mixed with molten cryolite, lowering the operating temperature and therefore the energy cost.
- Aluminium forms at the cathode, while oxygen forms at the carbon anodes.
- Hot oxygen reacts with the carbon anodes, so they burn away and need regular replacement.
- Electrolysis remains expensive because energy is needed both to keep the mixture molten and to drive the current.
The punchline: Mention cryolite, energy use and sacrificial carbon anodes for a complete extraction explanation.
Water adds rival ions to every aqueous electrolysis
An aqueous electrolyte is a crowded contest: ions from the solute compete with H⁺ and OH⁻ supplied by water.
- At the cathode, hydrogen forms unless the metal ion is from a metal less reactive than hydrogen; then the metal is deposited.
- At the anode, a halide ion forms its halogen; without a halide, oxygen forms from hydroxide ions.
- Copper chloride solution gives copper at the cathode and chlorine at the anode.
- Identify hydrogen by a squeaky pop, oxygen by relighting a glowing splint and chlorine by bleaching damp litmus paper.
The punchline: List all ions first, then apply the separate cathode and anode rules.
Higher Tier: half equations expose electron traffic
A half equation is a close-up at one electrode, with electrons written on the side that makes charge balance.
- At the cathode, positive ions gain electrons and are reduced: 2H⁺ + 2e⁻ → H₂.
- At the anode, negative ions lose electrons and are oxidised: 2Cl⁻ → Cl₂ + 2e⁻.
- Balance atoms first, then add electrons so total charge is equal on both sides.
- Across the complete electrolysis, electrons lost at the anode equal electrons gained at the cathode.
The punchline: Cathode reduction consumes electrons; anode oxidation produces them.
Words worth knowing
Key definitions
- oxidation
- Gain of oxygen, or at Higher Tier loss of electrons.
- reduction
- Loss of oxygen, or at Higher Tier gain of electrons.
- redox reaction
- A reaction in which oxidation and reduction occur together.
- reactivity series
- An ordering of elements by how readily they undergo chemical reactions.
- displacement reaction
- A reaction in which a more reactive element replaces a less reactive element in a compound.
- native metal
- An uncombined metal found naturally as the element rather than in a compound.
- acid
- A substance that produces hydrogen ions, H⁺, in aqueous solution.
- base
- A substance that neutralises an acid; metal oxides and hydroxides are common bases.
- alkali
- A soluble base that produces hydroxide ions, OH⁻, in aqueous solution.
- neutralisation
- A reaction between an acid and a base in which hydrogen ions are consumed and a salt is formed.
- salt
- An ionic compound formed when the hydrogen ions of an acid are replaced by positive ions.
- strong acid
- An acid that ionises completely in aqueous solution.
- weak acid
- An acid that ionises only partially in aqueous solution.
- electrolyte
- A molten or aqueous ionic substance whose mobile ions carry electric current.
- electrolysis
- The decomposition of an electrolyte by passing an electric current through it.
- cathode
- The negative electrode in electrolysis, where positive ions are reduced.
- anode
- The positive electrode in electrolysis, where negative ions are oxidised.
- end point
- The stage in a titration when the indicator just changes colour.
- titre
- The volume delivered from a burette between its initial reading and the end point.
Calculate with confidence
Equations
Balanced magnesium oxidation
2Mg + O₂ → 2MgO
Magnesium combines with oxygen and is oxidised to magnesium oxide.
Exam tip: Use the coefficient 2 before Mg and MgO; do not change MgO's formula.
Dilute hydrochloric or sulfuric acid and a metal
dilute hydrochloric or sulfuric acid + magnesium, zinc or iron → salt + hydrogen
Magnesium, zinc or iron reacts with dilute hydrochloric or sulfuric acid to release hydrogen.
Exam tip: Hydrochloric acid selects chloride and sulfuric acid selects sulfate; do not use nitric acid in this hydrogen pattern.
Neutralisation product patterns
acid + base → salt + water; acid + metal carbonate → salt + water + carbon dioxide
Bases neutralise acids; carbonates also release carbon dioxide gas.
Exam tip: Do not predict hydrogen for an acid–carbonate reaction.
Neutralisation ionic equation
H⁺(aq) + OH⁻(aq) → H₂O(l)
Hydrogen ions and hydroxide ions combine in a 1:1 ratio to make water.
Exam tip: Check both atom count and total charge: each side is neutral overall.
Higher Tier electrode examples
cathode: 2H⁺ + 2e⁻ → H₂; anode: 2Cl⁻ → Cl₂ + 2e⁻
Reduction consumes electrons at the cathode; oxidation releases electrons at the anode.
Exam tip: Place electrons so atoms and total charge balance on both sides of each half equation.
Follow it step by step
Processes to remember
How to deduce a reactivity order
- Compare reactions under the same conditions, including metal size, acid concentration and temperature.
- Rank faster or more vigorous water and acid reactions above slower reactions; place no-reaction results lower.
- For displacement, place the free metal above any metal it displaces from a compound.
- Combine all comparisons into one order and check that every observation fits.
Exam tip: An observation earns the comparison; do not rank metals from memory when the question supplies evidence.
How to choose a metal extraction route
- Decide whether the metal is so unreactive that it may occur native.
- Locate the metal relative to carbon in the reactivity series.
- Choose carbon reduction for an oxide below carbon.
- Choose electrolysis of a molten compound for a metal above carbon and mention the high energy demand.
Exam tip: Name the position relative to carbon before naming the method.
How to predict and name an acid reaction
- Identify the acid: hydrochloric makes chlorides and sulfuric makes sulfates; nitric acid makes nitrates with bases, alkalis or carbonates.
- Take the first part of the salt name from the metal, base or carbonate.
- Choose the product pattern: magnesium, zinc or iron with dilute hydrochloric or sulfuric acid gives hydrogen; a base gives water; a carbonate gives water and carbon dioxide.
- Write correct formulae, then balance the symbol equation with coefficients if asked.
Exam tip: Copper oxide plus sulfuric acid makes copper sulfate and water.
How to predict aqueous electrode products
- List ions from the dissolved compound and include H⁺ and OH⁻ from water.
- At the cathode, choose hydrogen unless the competing metal is less reactive than hydrogen.
- At the anode, choose the halogen if a halide is present; otherwise choose oxygen.
- Name suitable observations or tests, then write Higher Tier half equations if required.
Exam tip: Do not apply the molten rule to a solution: water changes the list of possible products.
Higher Tier: how to balance a half equation
- Write the ion and the element produced at that one electrode.
- Balance atoms, remembering that hydrogen, chlorine and other elemental gases may be diatomic.
- Add electrons to the more positive side until total charge is equal.
- Check the direction: cathode reduction gains electrons; anode oxidation loses them.
Exam tip: Electrons belong on the left for reduction and on the right for oxidation.
See the thinking
Worked example
Worked example: predict electrolysis products
Concentrated aqueous sodium chloride is electrolysed using inert electrodes. Predict the products, give a test for each gas and, at Higher Tier, write suitable half equations.
- List Na⁺ and Cl⁻ from the solute, plus H⁺ and OH⁻ from water.
- At the cathode, sodium is more reactive than hydrogen, so hydrogen is produced.
- At the anode, chloride is a halide, so chlorine is produced.
- Test hydrogen with a lighted splint for a squeaky pop; chlorine bleaches damp litmus paper.
- Higher Tier: 2H⁺ + 2e⁻ → H₂ and 2Cl⁻ → Cl₂ + 2e⁻.
Answer: Hydrogen forms at the negative cathode and chlorine at the positive anode. Hydrogen gives a squeaky pop; chlorine bleaches damp litmus. Higher Tier half equations are 2H⁺ + 2e⁻ → H₂ and 2Cl⁻ → Cl₂ + 2e⁻.
Sodium does not form because it is more reactive than hydrogen, while chloride wins the anode rule because it is a halide. The half equations show reduction at the cathode and oxidation at the anode.
Protect the marks
Common mistakes
Watch out: Reversing oxidation and reduction in electron terms.
Do this instead: Higher Tier: oxidation is electron loss; reduction is electron gain.
Watch out: Saying every metal reacts with dilute acid.
Do this instead: Only suitable metals above hydrogen do; copper is below hydrogen and does not.
Watch out: Teaching reactions with steam in this reactivity comparison.
Do this instead: The specified comparison here is at room temperature; steam reactions are outside this scope.
Watch out: Claiming carbon extracts every metal from its oxide.
Do this instead: Carbon reduction works for metals below carbon; more reactive metals need electrolysis.
Watch out: Predicting hydrogen from an acid–carbonate reaction.
Do this instead: A carbonate gives salt, water and carbon dioxide.
Watch out: Treating a strong acid as the same thing as a concentrated acid.
Do this instead: Strength is extent of ionisation; concentration is amount dissolved per volume.
Watch out: Evaporating a salt solution rapidly until completely dry.
Do this instead: Concentrate gently, stop heating and let crystals form while cooling.
Watch out: Calling the cathode positive during electrolysis.
Do this instead: The cathode is negative and attracts positive ions; the anode is positive.
Watch out: Trying to electrolyse a solid ionic compound.
Do this instead: Its ions are fixed; melt it or dissolve it so ions can move.
Watch out: Using molten-electrolysis rules for an aqueous solution.
Do this instead: Water adds H⁺ and OH⁻, so competing ions alter the possible products.
Watch out: Putting electrons on the wrong side of a half equation.
Do this instead: Reduction gains electrons on the left; oxidation releases electrons on the right.
Watch out: Averaging a rough titration with precise titres.
Do this instead: Use only concordant precise titres and exclude justified anomalies.
Plan it like the exam
Required practicals
Prepare a pure, dry soluble salt
Combined Science and separate Chemistry
Aim: Prepare pure, dry copper sulfate crystals from dilute sulfuric acid and insoluble copper oxide.
Method
- Wear eye protection and warm a measured volume of dilute sulfuric acid gently in a water bath.
- Add small portions of copper oxide while stirring until fresh solid no longer reacts and some remains in excess.
- Filter the mixture to remove the excess copper oxide, collecting the blue copper sulfate solution as filtrate.
- Heat the filtrate gently in an evaporating basin until some water has evaporated and the solution is concentrated.
- Stop heating before the solution dries out and leave it to cool so crystals form.
- Separate the crystals from the remaining solution and dry them between clean filter papers.
Variables
- Independent
- No investigative variable; the prescribed acid and insoluble base determine the salt made.
- Dependent
- Formation and quality of the pure, dry copper sulfate crystal sample.
- Controls
- volume and concentration of sulfuric acid
- warming and evaporation conditions
- addition of copper oxide until a visible excess remains
- clean apparatus and drying method
Analysis: Excess copper oxide shows that all acid has reacted; filtration removes that solid excess. Controlled evaporation concentrates the filtrate, and slow cooling forms crystals. Record crystal appearance and dry mass only after removing surface solution.
Safety
- Wear eye protection; dilute sulfuric acid and copper sulfate solution can irritate or harm eyes and skin.
- Avoid inhaling copper oxide powder and clean spills using the school's chemical-safety procedure.
- Use a water bath or gentle controlled heating, handle hot glassware carefully and do not heat the solution to dryness.
Improvements
- Add copper oxide in small portions so a clear excess is reached without wasting large amounts.
- Rinse the reaction vessel into the filter with a small amount of distilled water to reduce transfer loss.
- Allow slow cooling and dry crystals consistently before comparing mass.
Titrate a strong acid and strong alkali
Separate Chemistry only
Aim: Determine accurately the volumes of a strong acid and strong alkali that react completely.
Method
- Rinse the pipette with the solution it will measure, then transfer a fixed volume into a conical flask.
- Add a few drops of one suitable indicator, such as phenolphthalein or methyl orange, and place the flask on a white tile.
- Rinse and fill the burette with the other solution, remove the funnel and record the initial reading at eye level.
- Run solution into the flask while swirling; use a rough titration to locate the end point.
- Repeat, adding solution dropwise near the sharp colour change, then record the final burette reading.
- Calculate titre = final reading − initial reading, recording burette readings to the precision of the scale.
- Repeat until concordant titres are obtained and calculate their mean, excluding the rough result and justified anomalies.
Variables
- Independent
- Volume of solution delivered from the burette.
- Dependent
- Indicator end point and the corresponding titre.
- Controls
- fixed pipetted volume
- concentrations and identities of the acid and alkali
- indicator identity and number of drops
- apparatus rinsing and eye-level reading technique
Analysis: Find a mean from concordant titres. Higher Tier: convert volume from cm³ to dm³ by dividing by 1000, use reacting ratios to find an unknown concentration in mol/dm³, and multiply by Mr when g/dm³ is required.
Safety
- Wear eye protection and rinse acid or alkali splashes from skin immediately with plenty of water.
- Clamp the burette securely and fill it below eye level, using a funnel that is removed before readings.
- Handle glass pipettes and burettes carefully; use a pipette filler, never mouth pipetting.
Improvements
- Use a white tile and add titrant one drop at a time near the end point.
- Read the bottom of the meniscus at eye level to avoid parallax error.
- Obtain several concordant titres before calculating the mean.
Identify products of aqueous electrolysis
Combined Science and separate Chemistry
Aim: Investigate the products formed at inert electrodes when different aqueous ionic solutions are electrolysed.
Method
- Add a measured volume and concentration of one aqueous electrolyte to a labelled container.
- Insert clean inert electrodes at a fixed depth and separation, then connect them to a low-voltage direct-current supply.
- Switch on for a fixed time and record bubbles, colour changes or solid deposits separately at each electrode.
- Collect any gas in small inverted test tubes without swapping the cathode and anode samples.
- Test hydrogen with a lighted splint, oxygen with a glowing splint and chlorine with damp litmus paper, using only small samples.
- Switch off, clean the electrodes and repeat with the other solutions under the same conditions.
Variables
- Independent
- Identity of the dissolved ionic compound.
- Dependent
- Substance formed and identified at each electrode.
- Controls
- solution concentration and volume
- electrode material, area, depth and separation
- supply voltage and electrolysis time
- gas-collection and identification method
Analysis: Match each observation and gas test to a product, then compare it with the aqueous rules. A deposit may be weighed only after the electrode is rinsed and dried; identify cathode and anode evidence separately.
Safety
- Wear eye protection and use the low-voltage direct-current supply with dry hands.
- Use small quantities in good ventilation; chlorine is toxic, so do not inhale gases and stop after enough is collected for the test.
- Follow school hazard guidance for each electrolyte, including harmful copper compounds, and wash hands after the practical.
Improvements
- Clean electrodes between solutions to prevent contamination.
- Keep electrode geometry, concentration, voltage and time constant.
- Repeat observations and use fresh test reagents to confirm gas identities.
Try it before you move on
Quick check
Say your answer first, then open the card to check it.
Copper oxide loses oxygen and becomes copper. Has the copper oxide been oxidised or reduced?
Answer: Reduced.
Reduction is loss of oxygen; the copper oxide has lost oxygen.
Will zinc displace copper from copper sulfate solution?
Answer: Yes. Zinc is more reactive than copper.
A metal higher in the reactivity series displaces a lower metal from its compound.
Why is aluminium extracted by electrolysis rather than reduction with carbon?
Answer: Aluminium is more reactive than carbon.
Carbon cannot remove oxygen from aluminium oxide, so the molten compound is electrolysed.
What products form when calcium carbonate reacts with nitric acid?
Answer: Calcium nitrate, water and carbon dioxide.
The carbonate gives the calcium name; nitric acid makes a nitrate, and carbonates also produce water and carbon dioxide.
Higher Tier: hydrogen-ion concentration rises by a factor of 100. How many pH units does pH fall?
Answer: Two pH units.
Higher Tier: each one-unit fall represents a tenfold increase, and 10 × 10 = 100.
What forms at the cathode during electrolysis of molten lead bromide?
Answer: Lead metal.
Positive lead ions move to the negative cathode and are discharged.
What usually forms at the anode in an aqueous solution containing no halide ions?
Answer: Oxygen.
Without chloride, bromide or iodide, hydroxide ions from water are discharged to form oxygen.
Higher Tier: is Cl₂ + 2e⁻ → 2Cl⁻ oxidation or reduction?
Answer: Reduction.
Chlorine gains electrons; electron gain is reduction.
Good questions, clear answers
Frequently asked questions
What is the quickest way to remember oxidation and reduction?
For oxygen, oxidation is gain and reduction is loss. At Higher Tier, use OIL RIG: Oxidation Is Loss of electrons; Reduction Is Gain.
Why are carbon and hydrogen in a metal reactivity series?
They are non-metal reference points. Carbon helps choose an extraction method, while hydrogen predicts whether a metal reacts with dilute acid.
How do I know the name of a salt?
The metal, base or carbonate gives the first name. Hydrochloric acid makes chloride, sulfuric acid sulfate and nitric acid nitrate.
Why add an insoluble base in excess when making a salt?
A visible excess shows that all acid has reacted. The remaining solid can be filtered off, leaving salt solution without unreacted acid.
Is a concentrated acid always strong?
No. Concentration is amount per volume; strength is the fraction ionised. A concentrated weak acid is possible, as is a dilute strong acid.
Why can molten ionic compounds conduct but solids cannot?
Both contain ions, but only the molten compound has ions free to move through the liquid and carry charge.
How do aqueous electrolysis products differ from molten products?
Water adds H⁺ and OH⁻ ions. These compete with solute ions, so hydrogen or oxygen may form instead of the metal or non-metal expected from a molten compound.
Which practical is separate Chemistry only?
The strong acid–strong alkali titration is separate-only. Preparing a soluble salt and electrolysing aqueous solutions are shared with Combined Science: Trilogy.
15 illustrated pages in this topic
Pro unlocks every illustrated page, syllabus-matched Test Me recall and a printable PDF for this topic, plus the same revision loop throughout your saved course and tier. Test Me is available in all 25 topics across the full catalogue.