MYP Chemistry · Topic 7
Types of chemical reaction
Most reactions are either a hand-over of protons or a hand-over of electrons. Learn to see which, and the products follow.
Topic: Types of chemical reaction (acids and bases, neutral solutions, acid/base reactions, pH and indicators, formation of salts, uses of salts; redox reactions, reactivity series; extraction of metals, and corrosion, electrochemical cells).
Prior knowledge: ions and ionic formulas, balancing equations and moles (Topic 6); metals and non-metals, and the trends in groups 1 and 17 (Topic 1).
Learning objectives
Objectives and contextOverview
By the end of this chapter you should be able to:
- describe acids, bases, alkalis and neutral solutions and state their typical pH values A
- use indicators and pH data to classify solutions and explain what a change in pH means A C
- predict the products of acids reacting with metals, bases and carbonates and determine the formula of the salt formed A
- explain oxidation and reduction as the loss and gain of electrons and use the reactivity series to predict displacement A
- outline how metals are extracted and how corrosion is prevented, and describe how an electrochemical cell produces a voltage A
- design and evaluate investigations involving pH, acids and metals, and discuss the impacts of these reactions B C D
Behind the huge variety of chemical reactions lie a few patterns. Two of them account for much of everyday chemistry. In acid–base reactions hydrogen ions are transferred — this is what happens when an antacid settles an upset stomach, when lime is spread on acidic soil, or when a bee sting is treated. In redox reactions electrons are transferred — this is what happens when iron rusts, when a battery powers a phone, and when metals are extracted from their ores. Recognising which pattern a reaction follows lets you predict its products without memorising each reaction separately.
02 / Acids and bases
Acids, bases and alkalisAcids and bases
An acid is a substance that produces hydrogen ions, H+(aq), when it dissolves in water. It is the hydrogen ion that gives acids their shared properties: a sour taste, the ability to turn blue litmus red, and reactions with metals, bases and carbonates. Common laboratory acids are hydrochloric acid (HCl), sulfuric acid (H2SO4) and nitric acid (HNO3); ethanoic acid in vinegar, citric acid in fruit and methanoic acid in ant stings are weaker, naturally occurring acids.
A base is a substance that neutralises an acid. Metal oxides, metal hydroxides and metal carbonates are bases. A base that dissolves in water is called an alkali; alkalis produce hydroxide ions, OH−(aq), in solution. Sodium hydroxide and calcium hydroxide are alkalis. Copper(II) oxide is a base but not an alkali, because it does not dissolve in water.
Acid: produces H+ ions in aqueous solution; a proton donor.
Base: a substance that reacts with an acid to neutralise it; a proton acceptor.
Alkali: a soluble base; produces OH− ions in aqueous solution.
Neutral solution: neither acidic nor alkaline; pH 7 at 25 °C — pure water and solutions of many salts.
Strong acids, such as hydrochloric acid, dissociate completely into ions in water. Weak acids, such as ethanoic and methanoic acid, dissociate only partly and set up an equilibrium (Topic 6). “Strong” describes how completely an acid dissociates; “concentrated” describes how much acid is dissolved in a given volume. A dilute solution of a strong acid and a concentrated solution of a weak acid can have similar pH values.
03 / pH
pH and indicatorspH and indicators
The pH scale measures how acidic or alkaline a solution is. It runs from about 0 to 14. Below 7 a solution is acidic, 7 is neutral and above 7 it is alkaline. The lower the pH, the higher the concentration of hydrogen ions and the more acidic the solution; each step of one pH unit is a tenfold change in hydrogen-ion concentration.
A decrease in pH means an increase in acidity. Examination feedback reports students confusing “decreased pH” with “decreased acidity”. When you describe a pH change, say which way the acidity changed as well: “the pH fell from 7 to 4, so the solution became more acidic.”
An indicator is a substance whose colour depends on pH. Litmus is red in acid and blue in alkali. Universal indicator is a mixture of indicators that changes through a range of colours, so it gives an approximate pH value — red for strongly acidic, green for neutral, purple for strongly alkaline. Many plant juices, such as red cabbage and cherry juice, are natural indicators. Because an indicator gives a colour that has to be matched by eye, it can only estimate pH; a pH probe (pH meter) gives a numerical reading, which is quantitative and more precise.
Choose a pH probe when a question needs an exact value or small differences between samples; an indicator is enough to decide whether a solution is acidic, neutral or alkaline. A natural indicator that only changes colour at one point cannot give the exact pH of each fruit juice — it can only tell you which side of that point the pH lies.
04 / Neutralisation
Neutralisation and saltsAcid/base reactions
When an acid reacts with a base, the hydrogen ions of the acid combine with the base and the acidity is removed. This is neutralisation. With an alkali the essential reaction is between hydrogen ions and hydroxide ions, forming water:
H+(aq) + OH−(aq) → H2O(l)
The other ions — the metal ion from the base and the negative ion from the acid — form a salt. A salt is the compound formed when the hydrogen of an acid is replaced by a metal (or ammonium) ion. Neutralisation is exothermic: the temperature of the mixture rises. Four general reactions produce salts:
| Reaction | General equation | Example | Observation |
|---|---|---|---|
| Acid + alkali | acid + alkali → salt + water | HCl + NaOH → NaCl + H2O | Temperature rises; no visible change |
| Acid + metal oxide | acid + base → salt + water | H2SO4 + CuO → CuSO4 + H2O | Black solid dissolves; blue solution |
| Acid + carbonate | acid + carbonate → salt + water + carbon dioxide | CaCO3 + 2HCl → CaCl2 + H2O + CO2 | Fizzing; the gas turns limewater milky |
| Acid + reactive metal | acid + metal → salt + hydrogen | Mg + 2HCl → MgCl2 + H2 | Fizzing; the gas gives a squeaky pop with a lit splint |
The name of the salt comes from both parents. The first part is the metal (or ammonium); the second comes from the acid: hydrochloric acid gives chlorides, sulfuric acid gives sulfates and nitric acid gives nitrates. The formula is found by balancing the charges of the two ions, exactly as in Topic 6.
Question: determine the formula of the salt formed when calcium carbonate neutralises sulfuric acid.
Ions: Ca2+ from the carbonate, SO42− from the acid. The charges are equal and opposite, so one of each.
Answer: CaSO4. The full equation, CaCO3 + H2SO4 → CaSO4 + H2O + CO2, is useful working, but “determine” asks for one answer: examination feedback reports students who wrote the equation and never identified the salt.
A solution of a salt formed from a strong acid and a strong alkali, such as sodium chloride or calcium sulfate, is neutral — pH 7. When a question asks for the pH of such a salt solution, the answer is 7.
Salt maker
Choose an acid and something to react with it. The model names the products, works out the salt's formula from the charges of its ions, and writes the balanced equation.
Neutralisation and pH
Alkali is added in steps to 25.0 cm3 of a strong acid. The model calculates the concentration of hydrogen or hydroxide ions left over after each addition and converts it to pH (pH = −log[H+], with [H+][OH−] = 1.0 × 10−14 at 25 °C). The indicator colour is approximate.
05 / Salts in use
Making and using saltsFormation and uses of salts
A soluble salt can be made from an insoluble base or carbonate in three steps, each with a clear purpose. First, add the solid to warm dilute acid until no more dissolves, so that the base is in excess and all the acid has been used up. Second, filter to remove the unreacted solid, which leaves a solution of the salt only. Third, heat the filtrate gently to evaporate some water and leave it to cool, so that crystals form (crystallisation, Topic 5). Copper(II) sulfate crystals are made in this way from copper(II) oxide and sulfuric acid.
Salts have many uses, and each use follows from a property of the ions:
| Salt | Use | Property used |
|---|---|---|
| Sodium chloride | Food preservation, de-icing roads | Dissolved ions lower the freezing point of water |
| Calcium carbonate | Neutralising acidic soil and lakes; antacids | Reacts with acid (a base) |
| Ammonium nitrate, potassium phosphate | Fertilisers | Soluble sources of nitrogen, potassium and phosphorus |
| Strontium, barium and copper salts | Fireworks | Characteristic flame colours (Topic 1) |
| Calcium sulfate | Plaster, building materials | Sets hard when mixed with water |
Solutions of salts conduct electricity. When an ionic compound such as sodium chloride dissolves, its ions separate and are free to move; moving charged particles carry the current. That is why the electrical conductivity of water can be used to estimate how much salt it contains. Examination feedback reports that most students could not make this link — a salt is an ionic compound, so dissolving it releases ions, and the ions are what conduct.
06 / Redox
Oxidation and reductionRedox reactions
Originally, oxidation meant gaining oxygen and reduction meant losing it: magnesium is oxidised when it burns to magnesium oxide, and iron oxide is reduced when carbon removes its oxygen. The more general and more useful definition is in terms of electrons.
Oxidation is the loss of electrons. Reduction is the gain of electrons. (One way to remember it: OIL RIG — oxidation is loss, reduction is gain.)
The two always happen together, because the electrons lost by one substance are gained by another. A reaction in which both occur is a redox reaction.
When magnesium burns, each magnesium atom loses two electrons and becomes Mg2+; each oxygen atom gains two and becomes O2−. Magnesium is oxidised and oxygen is reduced. The changes can be written as half-equations:
Mg → Mg2+ + 2e− (oxidation) O2 + 4e− → 2O2− (reduction)
A useful bookkeeping tool is the oxidation number: 0 for an uncombined element, and equal to the charge for a simple ion. An increase in oxidation number is oxidation; a decrease is reduction. When iron reacts to form iron(II) carbonate, FeCO3, iron goes from 0 to +2: it has lost two electrons, so it has been oxidised.
A full description names the substance, the direction of electron transfer and the term: “iron loses electrons (oxidation number increases from 0 to +2), so iron is oxidised.” At the electrodes of a cell, the same reasoning gives: at one electrode electrons are lost (oxidation), at the other electrons are gained (reduction).
07 / Reactivity
The reactivity seriesReactivity series
Metals react by losing electrons to form positive ions, and they differ in how readily they do so. Arranging metals in order of how vigorously they react with water, oxygen and acids gives the reactivity series. Hydrogen and carbon are included as reference points: metals above hydrogen react with dilute acids to release hydrogen; metals below carbon can be extracted from their oxides by heating with carbon.
| Metal | With water | With dilute acid |
|---|---|---|
| Potassium, sodium | Violent with cold water | Dangerously violent |
| Calcium | Steady with cold water | Vigorous |
| Magnesium | Very slow with cold water; reacts with steam | Rapid fizzing |
| Aluminium, zinc, iron | React with steam | Fizzing, slower down the list |
| (carbon) | reference point for extraction | |
| Tin, lead | No reaction | Very slow |
| (hydrogen) | reference point for reaction with acids | |
| Copper, silver, gold | No reaction | No reaction |
The rate at which hydrogen is produced when metals react with the same acid reflects their reactivity: magnesium fizzes rapidly, zinc steadily, iron slowly and copper not at all.
Displacement
A more reactive metal displaces a less reactive metal from a solution of its compound. The more reactive metal gives up its electrons to the ions of the less reactive one:
Zn(s) + CuSO4(aq) → ZnSO4(aq) + Cu(s) ionic: Zn + Cu2+ → Zn2+ + Cu
Observations: a pink-brown coating of copper forms on the zinc, the blue colour of the copper(II) ions fades, and the mixture warms up. Zinc is oxidised (loses electrons) and copper ions are reduced (gain electrons). If the metal is less reactive than the metal in the solution — copper placed in zinc sulfate — there is no reaction. Displacement reactions can be used to place metals in order when their reactions with water and acid are too slow to compare.
Displacement predictor
Choose a metal and a salt solution. The model uses the reactivity series to decide whether displacement occurs and, if it does, writes the ionic equation and identifies what is oxidised and what is reduced.
08 / Extraction
Extracting metalsExtraction of metals
Only the least reactive metals, such as gold, are found uncombined in the Earth's crust. Most metals occur in ores — rocks containing a metal compound, often an oxide — and extracting the metal means reducing the metal ions back to atoms. The method depends on the metal's position in the reactivity series.
- Metals below carbon (zinc, iron, tin, lead, copper) are extracted by heating the oxide with carbon. Carbon is more reactive, so it removes the oxygen. In the blast furnace, iron(III) oxide is reduced by carbon monoxide made from coke:
Fe2O3 + 3CO → 2Fe + 3CO2
Iron is reduced (it loses oxygen; its oxidation number falls from +3 to 0) and carbon monoxide is oxidised. - Metals above carbon (aluminium, magnesium, sodium, potassium) hold on to oxygen too strongly for carbon to remove it. They are extracted by electrolysis of the molten compound, in which an electric current supplies the electrons that reduce the metal ions. Electrolysis uses large amounts of electrical energy, which is why aluminium is more expensive to produce than iron and why recycling it saves so much energy.
Every extraction has impacts: mining and quarrying scar landscapes and produce waste rock; the carbon used in reduction is released as carbon dioxide; electrolysis needs a large, constant electricity supply. Recycling metals conserves ores and energy and reduces landfill.
09 / Corrosion
CorrosionCorrosion
Corrosion is the gradual oxidation of a metal by substances in its environment — extraction running in reverse. The corrosion of iron and steel is called rusting. It needs both oxygen and water; iron kept in dry air, or in boiled water with a layer of oil to keep air out, does not rust. Dissolved salt and acids speed rusting up, which is why cars rust faster where salt is spread on winter roads, and why chloride-containing de-icers can damage aircraft.
4Fe(s) + 3O2(g) + 2xH2O(l) → 2Fe2O3·xH2O(s) (hydrated iron(III) oxide — rust)
Rust is a problem because it is weak and crumbly and flakes away, exposing fresh iron underneath to further attack. Other metals, such as aluminium, form a thin, tough oxide layer that sticks to the surface and protects the metal beneath — which is why aluminium, although reactive, does not corrode away.
| Method | How it works |
|---|---|
| Painting, oiling, greasing, plastic coating | A barrier keeps oxygen and water away from the iron |
| Galvanising (coating with zinc) | A barrier; and zinc is more reactive, so it corrodes in preference to iron even if the coating is scratched |
| Sacrificial protection (blocks of magnesium or zinc on ships' hulls and pipelines) | The more reactive metal is oxidised instead of the iron, and is replaced periodically |
| Alloying (stainless steel with chromium) | Chromium forms a protective oxide layer on the surface |
Zinc and magnesium are above iron in the reactivity series, so they lose electrons more readily. When they are in contact with iron, they are oxidised first and supply electrons to the iron, which therefore is not oxidised. This is the same electron-transfer reasoning as displacement.
10 / Cells
Electrochemical cellsElectrochemical cells
In a displacement reaction, electrons pass directly from the more reactive metal to the ions of the less reactive one, and the energy is released as heat. If the two halves of the reaction are separated and connected by a wire, the electrons must travel through the wire — and a flow of electrons is an electric current. This arrangement is an electrochemical cell: it converts chemical energy into electrical energy.
Each metal dips into a solution of its own ions. At the more reactive metal, atoms lose electrons and go into solution as ions — oxidation; this electrode is the negative terminal. At the less reactive metal, ions from the solution gain electrons and are deposited as metal — reduction; this electrode is the positive terminal. The circuit is completed by a salt bridge, or a strip of filter paper soaked in an electrolyte, through which ions move between the two solutions. Without it the charges would build up and the current would stop immediately; a dry filter paper therefore gives no voltage.
Zn(s) → Zn2+(aq) + 2e− Cu2+(aq) + 2e− → Cu(s)
The further apart the two metals are in the reactivity series, the larger the voltage of the cell. As the cell runs, the zinc electrode loses mass and the copper electrode gains mass, and the blue colour of the copper(II) solution fades as its ions are used up — so the voltage slowly falls. Batteries are one or more cells in series; rechargeable batteries, such as lithium-ion batteries, use reactions that can be driven backwards by an external supply.
Voltages measured in a school laboratory rarely match published values exactly. The published values are measured under standard conditions (fixed concentration of 1 mol dm−3, 25 °C, pure metals with clean surfaces). Differences in concentration, temperature, surface oxide on the electrodes and the resistance of the salt bridge all change the measured voltage — useful points when a question asks you to evaluate a cell investigation.
Build a cell
Choose two metals. The model identifies the negative electrode (more reactive, oxidised) and the positive electrode (less reactive, reduced), writes both half-equations and estimates the voltage from standard electrode potentials (reference values, not the examination data).
11 / Practice
Examination practicePast examination tasks
The tasks below come from past on-screen examinations and are grouped by the criterion they assess. Criterion A items test acids, salts and redox directly; criteria B and C use investigations of pH, acid–carbonate reactions and burning candles; criterion D applies redox to capturing carbon dioxide. Attempt each one before opening the marking guidance.
A Criterion A · Knowing and understanding
In the 1950-1960s, the radioactive element radium was used to coat the numbers on watches. The numbers glowed in the dark so it was possible to see the time at night. The workers that painted the dials often became ill after long exposure to the paint.
As society became fascinated with nuclear power, children’s toys were produced to encourage them to learn all about radioactive elements. The Atomic Energy Lab in the image contained radioactive lead (Pb-210), ruthenium (Ru-106), zinc (Zn-65) and polonium (Po-210).
Today, toys that glow in the dark do not contain radioactive materials but rely instead on the properties of strontium nitrate and zinc sulphide to emit light.
Strontium nitrate (Sr(NO3)2) can be made by reacting strontium carbonate (SrCO3) with nitric acid (HNO3). State the name of the products to complete the word equation for this reaction.
Strontium carbonate + nitric acid → Strontium nitrate + ………………………….+………………………
Marking guidance
- Strontium carbonate + nitric acid → strontium nitrate + Accept correct formulae
- Carbon dioxide
- Water
Hydrangea flowers are known for changing colour in response to changes in soil pH.
Calcium carbonate can be used to neutralize sulfuric acid, which produces a salt. Select the other two products that are formed when acid is neutralized by a metal carbonate:
sulfuric acid + calcium carbonate → salt + product A + product B
Product A: CO / CO2 / H2O Product B: CO / CO2 / H2O
Suggested answer
Not from an official marking scheme — a worked answer written for these notes.
Product A: CO2 (carbon dioxide). Product B: H2O (water). Acid + metal carbonate → salt + water + carbon dioxide.
Determine the formula of the salt that is produced when calcium carbonate neutralizes sulfuric acid.
Suggested answer
Not from an official marking scheme — a worked answer written for these notes.
CaSO4 (calcium sulfate): Ca2+ from the carbonate and SO42− from sulfuric acid, in a 1 : 1 ratio.
Candles can be used for a variety of reasons, such as relaxing or in ceremonies. The length of time that a candle burns will determine its use.
Select numbers to balance the equation showing complete combustion of C5H12. ….. C5H12(l) + …..O2(g) → ….CO2(g) + ……H2O(g)
Marking guidance
- Do not accept ? for pentane
- Reactants correctly balanced
- Products correctly balanced
Worked answer: C5H12(l) + 8O2(g) → 5CO2(g) + 6H2O(g). Check: 5 C, 12 H and 16 O on each side.
A group of students used a simulation to model the behaviour of electrochemical cells using different metals. They decided to use the simulation with different combinations of metals and 1 mol dm–3 solutions of their ions. The students recorded the voltage. Below is the table for the data they wanted to record:
Using ideas about redox reactions and transfer of electrons, outline the processes taking place at each electrode.
Anode
Cathode
Marking guidance
- Anode
- oxidation or electrons are lost 1 mark for correct identification of
- electrons are lost and oxidation reduction and oxidation in terms of electron gain or loss wherever seen
- Cathode
- electrons are gained or reduction
- electrons are gained and reduction
B Criterion B · Inquiring and designing
When families are at home, their typical diets will depend on their location and lifestyle but will not change much. In this globalized world, many people travel to countries with traditions and cultures that may be new to them. Besides seeing the sights, travellers will also try local foods and drinks. In some cases, there will be a significant change to their diet. In some countries, spicy food and ice-cold fruit smoothies are a regular part of the local diet. Some travellers suffer from heartburn when they are abroad because of this change to their diet. Some researchers believe that acidic foods may lead to heartburn. Heartburn happens when acid leaves the stomach and moves up into the food pipe just above the stomach. This will irritate the food pipe and will create a burning discomfort.
Michelle, Yuri and Pedro have recently moved to a tropical country and love the fruit juices. Fruit juices are acidic outside the body and stay acidic when digested. Fruit juices can therefore increase the acidity in the stomach and cause heartburn. The friends are interested in finding out which fruit juice is most acidic and would lower the pH in the stomach the most.
State the variables in this experiment. Independent variable: Dependent variable:
Control variable 1: Control variable 2:
Marking guidance
- Independent variable: type of fruit
- Dependent variable: pH accept “acidity”
- Control variables: mass of fruit, time of mixing fruit Do not accept “amount”
- same volume of water added, same volume of mixture tested, same equipment used
State the research question that this experiment would investigate.
Marking guidance
- a research question linking pH with type of fruit Accept “acidic”
A student did a comparison between two candles: one fragranced with vanilla and one fragranced with strawberry. His results are shown below.
| Vanilla candle in a glass jar | Mass / g | Burn time / hours |
|---|---|---|
| 1 | 185 | 29 |
| 2 | 185 | 24.3 |
| 3 | 185 | 28 |
| Average | 27.1 |
| Strawberry candle in a glass | Mass / g | Burn time / hours |
|---|---|---|
| 1 | 114 | 32 |
| 2 | 114 | 28.5 |
| 3 | 114 | 29.5 |
| Average |
State the independent variable and the dependent variable in this investigation. Independent variable: Dependent variable:
Marking guidance
- Independent variable: fragrance Accept “flavour”, “ingredient”
- Dependent variable: burn time
C Criterion C · Processing and evaluating
The results from the experiment are shown in the table below.
| Type of fruit juice | pH |
|---|---|
| apple | 3.5 |
| pear | 3.8 |
| grape | 2.9 |
| cranberry | 2.4 |
| tomato | 4.1 |
Using data in the table, identify the fruit juice and pH that could cause the most severe heartburn. Justify your answer using scientific reasoning.
Marking guidance
- cranberry and (pH =) 2.4
- this has the lowest pH or this has the highest acidity
- the higher the acidity in the stomach, the more severe the heartburn This mark is awarding for linking stomach acidity with heartburn so it can be awarded for an incorrect juice
The data in the table in part (c) was from one trial for each fruit. Outline the benefits of carrying out more than one trial for each experiment.
Marking guidance
- (with more than one trial you are able to) calculate the mean
- reduces experimental errors or increases accuracy
Calculate the average burn time for the strawberry candle.
Marking guidance
- 30(.0 hours) Ignore any units
- 30.0 correctly expressed to 3 sig figs ECF from first marking point for transcription error
The student hypothesized that the candle with vanilla fragrance would take a longer time to burn because the candle had a smaller surface area. Use the data in the tables above part (a) to evaluate the validity of the student’s hypothesis.
Marking guidance
- Not valid because the strawberry has a longer time to burn
- Accept two any additional points from the list below [max 2] Ignore references to surface area
- different containers
- different masses so no direct comparison
- insufficient data to test the hypothesis
- in first data set would need to repeat investigation due to the 24.3 hours
- different wick sizes
D Criterion D · Reflecting on the impacts of science
Climate change is a global threat recognized by the Intergovernmental Panel on Climate Change (IPCC). The cause of climate change is thought to be due to increased levels of greenhouse gases, mainly carbon dioxide. Increased levels of carbon dioxide have resulted in an increase in the average temperature of the Earth. In the future, severe droughts are predicted in some regions, while floods are predicted in other areas. Both situations will lead to famine and increased poverty.
In the natural carbon cycle, carbon dioxide comes from sources such as respiration, volcanic eruptions, fossil fuels and industrial processes such as cement making. One way that carbon dioxide can be removed from the atmosphere is by photosynthesis in plants. The plants act as a natural carbon dioxide sink. Extensive deforestation means that less carbon dioxide is trapped in trees so levels are increasing in the atmosphere.
One way to reduce carbon dioxide in the atmosphere is to react CO2 with other chemicals to form a solid material, for example as shown in the reaction below:
Fe(..) + CO2(..) + H2O(..) → FeCO3(s) + H2(g)
Select the state symbols for the reactants in the equation if the reaction was at 25 °C.
Marking guidance
- one correct
- all correct Do not accept H2O(aq)
Select the name of FeCO3.
Marking guidance
- Iron (II) carbonate
Describe what happens to iron in the reaction in part (b) in terms of redox chemistry.
Marking guidance
- oxidation number of iron has increased or has lost electrons
- (from 0 to) +2
- (so) iron has been oxidised
Quick check: acids, salts and redox
Original practice questions for retrieval — not past examination items.
12 / Examination feedback
Examiner's overall observationEvidence from examination feedback
Examiner's overall observation
Most students knew that pure water has a pH of 7, and many could link the rate at which hydrogen is produced to the reactivity of the metal reacting with an acid. Determining the formula of a salt from the overall reaction was generally well prepared, yet on one task — the salt formed from calcium carbonate and sulfuric acid — many students wrote the whole equation without identifying the salt, and some wrote the formula without correct subscripts. Linking a pH value to the cause of a pH change was reported as difficult. A number of students did not know that the solution of this salt would be neutral, pH 7. Students confused a decrease in pH with a decrease in acidity. Most could not explain why electrical conductivity reveals the salt content of water: the link that a salt is an ionic compound which releases free ions when it dissolves, and that these ions carry the current, was rarely made. In investigations involving acids, some students named a calculated quantity such as rate as the dependent variable instead of something measured — the volume of gas, the number of bubbles or the change in mass — and without it could not choose suitable equipment; others did not name all the materials or metals to be tested, so they could not collect sufficient data. Asked how to improve the validity of a pH investigation, most suggested more trials rather than recognising what needed to be measured.
13 / Summary
Summary and knowledge organiserRevision
Essential knowledge
- Acids produce H+(aq); alkalis are soluble bases and produce OH−(aq). pH < 7 acidic, 7 neutral, > 7 alkaline; lower pH means more acidic.
- Indicators estimate pH from a colour; a pH probe gives a numerical value.
- acid + alkali/base → salt + water; acid + carbonate → salt + water + carbon dioxide; acid + metal → salt + hydrogen. Neutralisation is exothermic: H+ + OH− → H2O.
- Chlorides from hydrochloric acid, sulfates from sulfuric acid, nitrates from nitric acid; formula from the charges of the ions.
- Oxidation is loss of electrons (oxidation number increases); reduction is gain (oxidation number decreases).
- A more reactive metal displaces a less reactive one from its compounds. Metals below carbon are extracted by reduction with carbon; metals above carbon by electrolysis.
- Rusting needs oxygen and water; prevented by barriers, galvanising and sacrificial metals.
- In a cell the more reactive metal is oxidised (negative electrode) and ions of the less reactive metal are reduced (positive electrode); ions move through the salt bridge.
Definitions
- Acid — H+ donor in water
- Alkali — soluble base, gives OH−
- Salt — acid's H replaced by a metal or ammonium ion
- Oxidation — loss of electrons
- Reduction — gain of electrons
- Corrosion — oxidation of a metal by its environment
Equations
- CaCO3 + 2HCl → CaCl2 + H2O + CO2
- Mg + 2HCl → MgCl2 + H2
- Zn + Cu2+ → Zn2+ + Cu
- Fe2O3 + 3CO → 2Fe + 3CO2
Observations and tests
- CO2 turns limewater milky
- H2 gives a squeaky pop with a lit splint
- O2 relights a glowing splint
- Universal indicator: red → green → purple
- Displacement of copper: pink-brown solid, blue fades
Must-remember distinctions
- Strong/weak ≠ concentrated/dilute
- Base vs alkali (alkali dissolves)
- Lower pH = higher acidity
- “Determine the salt” = its formula, not the equation
- Indicator (qualitative) vs probe (quantitative)
Examination checklist
- Name the salt from both parent compounds
- Balance charges, then use subscripts
- Say who loses and who gains electrons
- Measured DV: gas volume, mass change, time
- Link conductivity to free ions
Other chapters: Criteria A–D · 1 · Periodic table · 2 · IUPAC naming · 3 · Atmosphere · 4 · Matter · 5 · Pure and impure · 6 · Bonding · 7 · Types of reaction
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