Edexcel International A Level Chemistry
Topic 8: Redox Chemistry and Groups 1, 2 and 7
Redox is the movement of electrons, tracked with oxidation numbers. This topic uses it to explain the metals of Groups 1 and 2 and the halogens of Group 7, and then adds the tests and titrations that identify their compounds.
The one idea
Chemistry in this topic is electrons moving, and one tool follows them: the oxidation number. A species whose oxidation number goes up has lost electrons (oxidised); one whose number goes down has gained them (reduced). The two always happen together.
Three families then show the idea at work:
- Groups 1 and 2 are strong reducing agents. They lose their outer electrons easily, and the ease rises down each group because the outer electrons are further from the nucleus. Their compounds follow trends you can explain with ionic radius and charge density.
- Group 7 halogens are oxidising agents. They gain an electron to form halide ions, and the oxidising power falls down the group. The halide ions are reducing agents, and the reducing power rises down the group.
- Chlorine with water or alkali is a disproportionation: the same element is oxidised and reduced in the same reaction.
Almost every question here is one of four moves:
- Assign numbers. Find each oxidation number with the rules, then say what rose and what fell.
- Build the equation. Write the two half-equations, balance the electrons, and add.
- Explain a trend. Name the size, the charge and the electrons, in that order, and finish with the energy needed.
- Do the practical. Give the reagent, the observation, the colour and the conclusion, or carry the titration through to a concentration.
Where to find things
| Question | Pages |
|---|---|
| What is an oxidation number, and what are the rules? | 1 and 2 |
| How do Roman numerals in names work? | 3 |
| What are oxidation, reduction and oxidising and reducing agents? | 4 and 5 |
| How do I write half-equations and combine them? | 6 and 7 |
| What is disproportionation, and how do I classify a reaction? | 8 |
| Why does reactivity rise down Groups 1 and 2? | 9 and 10 |
| What are the oxides, hydroxides and their solubility trends? | 11 and 12 |
| Why do carbonates and nitrates decompose more easily up the group? | 13 and 14 |
| Which colour is which in a flame test? | 15 |
| How do I test for carbonate, sulfate and ammonium, and identify an unknown? | 16 and 17 |
| How do titrations work, and how do I calculate from one? | 18 and 19 |
| The two core practicals | 20 and 21 |
| How certain is a titration result? | 22 |
| What do the halogens look like and how do they react? | 23, 24 and 25 |
| What does chlorine do with alkali? | 26 |
| What do halides do with sulfuric acid and silver nitrate? | 27 and 28 |
| How do I predict fluorine and astatine? | 29 |
How to use these notes
- Read the pages in order the first time. Pages 1 to 8 give you the language and equations of redox. Pages 9 to 17 use it for Groups 1 and 2 and the tests for their compounds. Pages 18 to 22 are titration and its practicals. Pages 23 to 29 are Group 7.
- Each page opens with its one idea in a sentence. If you can explain it to someone else, you know the page.
- Coloured boxes mean the same thing on every page: definitions you may be asked to quote, key ideas, common mistakes, exam technique, practical steps to describe, and beyond the spec (interesting, but you will not be examined on it).
- Each page ends with questions and fully worked answers. Cover the answers, try the questions, then mark yourself.
- Use the topic on a page the night before an exam.
This topic includes Core Practical 3 (the concentration of a hydrochloric acid solution by titration) on page 20 and Core Practical 4 (a standard solution from a solid acid) on page 21. All titres, masses and observations in the worked examples are made up to show the method and are marked illustrative. Reference values, such as flame colours and boiling temperatures, are standard rounded values.
Constants used in this topic: relative atomic masses are the Edexcel IAL values to one decimal place: H 1.0, C 12.0, N 14.0, O 16.0, Na 23.0, Mg 24.3, S 32.1, Cl 35.5, K 39.1, Ca 40.1, Ba 137.3. Titration volumes are in cm³ and concentrations in mol dm⁻³, so you convert volume to dm³ by dividing by 1000.
The skills that earn most marks here are writing the sign first on an oxidation number, balancing electrons before you add half-equations, explaining trends with size and charge rather than "the atoms get bigger" alone, naming reagents and colours exactly, and never writing a titre average that includes the rough one.
Pages
- Oxidation number and the rules
Oxidation number is a bookkeeping device that tracks electrons in ionic and covalent substances alike. Eight short rules, applied in order, give every atom its number.
- Calculating oxidation numbers
Write the known numbers under the atoms, multiply by the number of atoms, set the total equal to the overall charge, and solve for the unknown. Peroxides and metal hydrides are the two exceptions the rules must be bent for.
- Roman numerals and writing formulae
A Roman numeral in a name is the oxidation number of that element. Read it to find the charge on an ion, and use it to write a formula whose numbers add up to zero.
- Oxidation and reduction
Oxidation is loss of electrons and a rise in oxidation number; reduction is gain of electrons and a fall in oxidation number. Metals form positive ions this way, non-metals negative ions.
- Oxidising and reducing agents
An oxidising agent takes electrons and is reduced; a reducing agent gives electrons and is oxidised. Group 1 and 2 metals are reducing agents and oxygen and the halogens are oxidising agents, in every s-block and p-block reaction you need.
- Building half-equations
A half-equation shows one species gaining or losing electrons. Balance the atoms (O with water, H with H⁺), then add electrons to balance the charge, and check against the change in oxidation number.
- Full ionic equations
Multiply the two half-equations until the electrons match, add them, and cancel anything that appears on both sides. Then check atoms and charge, and make sure no electrons are left.
- Disproportionation and classifying reactions
In disproportionation one element in one species is oxidised and reduced at the same time. Write the oxidation numbers on both sides of any equation to decide whether it is not redox, redox, or disproportionation.
- Ionisation energy and reactivity down Groups 1 and 2
Down both groups the outer electron is further from the nucleus and more shielded, so ionisation energy falls. Group 1 and 2 reactions are oxidations of the metal, so the metals become more reactive down the group.
- Reactions of Group 1 and 2 metals
The metals burn in oxygen and chlorine to give oxides and chlorides, and react with water to give a hydroxide and hydrogen. In every case the metal is oxidised, and the vigour increases down the group.
- Oxides and hydroxides of Groups 1 and 2
Metal oxides and hydroxides are bases. Oxides react with water to form hydroxides, and both oxides and hydroxides neutralise dilute acids to give a salt and water. No oxidation number changes, so none of these is redox.
- Solubility of Group 2 hydroxides and sulfates
Down Group 2 the hydroxides become more soluble and the sulfates become less soluble. The two trends run in opposite directions, and the insolubility of barium sulfate is the basis of the sulfate test.
- Thermal stability: size and charge of the cation
A small, highly charged cation has a high charge density and polarises the carbonate or nitrate ion, so the compound decomposes more easily. Stability to heat rises down each group, and Group 1 is more stable than Group 2.
- Decomposing carbonates and nitrates
Group 2 carbonates give the oxide and carbon dioxide. Group 1 nitrates give a nitrite and oxygen, but lithium and Group 2 nitrates give the oxide, brown nitrogen dioxide and oxygen. Heat equal amounts and time the tests to compare.
- Flame colours and electron transitions
Heat lifts an electron in a metal ion to a higher energy level. When it falls back it gives out light with a frequency set by the energy gap, so each metal gives its own flame colour. Learn the colours and the flame-test method.
- Tests for carbonate, hydrogencarbonate and sulfate ions
Dilute acid releases carbon dioxide from carbonates and hydrogencarbonates, which turns limewater milky. Acidified barium chloride gives a white precipitate of barium sulfate with sulfate ions, and the acid must go in first.
- Ammonium ions and identifying an unknown salt
Sodium hydroxide and gentle warming release ammonia from ammonium ions, which turns damp red litmus blue and makes white smoke with hydrogen chloride. Combine the ion tests, flame tests and decomposition tests to identify an unknown.
- Concentration and titration
Concentration is the amount of solute in 1 dm³ of solution, in mol dm⁻³ or g dm⁻³. A titration measures exactly how much of one solution reacts with a known volume of another, and an indicator shows the end point.
- Titration calculations
Write the equation, find the amount of the solution you know, use the ratio to get the amount of the other, then divide by its volume. Give the answer in mol dm⁻³ and g dm⁻³ to the right number of significant figures.
- Core Practical 3: concentration of hydrochloric acid
Titrate a measured volume of hydrochloric acid with a standard solution of sodium hydroxide, using methyl orange, until you have concordant titres, and calculate the concentration of the acid.
- Core Practical 4: a standard solution from a solid acid
Weigh a solid acid by difference, dissolve it and make it up to 250 cm³ in a volumetric flask to give a standard solution, then use it to find the concentration of a solution of sodium hydroxide.
- Measurement uncertainty in volumetric analysis
Every reading has an uncertainty set by the apparatus. Express each as a percentage of the value measured, add them to estimate the overall uncertainty in the result, and reduce it by making measurements large and using precise apparatus.
- Trends in Group 7
Down Group 7 the molecules have more electrons, so London forces and boiling temperatures rise (gas, liquid, solid). The atoms get bigger, so electronegativity and the ability to gain an electron fall: halogens become less reactive down the group.
- Colours of the halogens and displacement reactions
A more reactive halogen oxidises the halide ion of a less reactive one and displaces the halogen. Learn the colour of each halogen in water and in a non-polar solvent, so you can read the result of each test.
- Halogens with metals, and chlorine with water
Halogens oxidise Group 1 and 2 metals to ionic halides, and the reactions are most vigorous for chlorine. Chlorine reacts with water by disproportionation to give HCl and chloric(I) acid, which is why it is used to treat water.
- Chlorine with alkali: bleach and chlorate(V)
Chlorine disproportionates in sodium hydroxide. Cold, dilute alkali gives chloride and chlorate(I), which is bleach; hot alkali gives chloride and chlorate(V). Bromine and iodine behave in the same way.
- Halides with concentrated sulfuric acid
Every solid sodium halide gives a steamy hydrogen halide with concentrated sulfuric acid, but only HBr and HI go on to reduce the acid. Iodide reduces sulfur furthest, from +6 to −2, so the reducing ability is I⁻ > Br⁻ > Cl⁻.
- Silver nitrate test and hydrogen halides
Acidified silver nitrate gives a white, cream or yellow precipitate with chloride, bromide or iodide ions, which dissolve in ammonia in a set order. Hydrogen halides make white smoke with ammonia and strong acids with water.
- Predicting fluorine and astatine
Fluorine and astatine sit at the two ends of Group 7. Apply the trends you know, in the same direction, to predict their state, colour, oxidising power and reactions, and the properties of their compounds.
Also in this topic: Topic 8 on a page
Spec checklist
Every point you can be examined on in this topic, and where it is explained. Tick them off as you go.
| Spec | You should be able to | Where |
|---|---|---|
| 8.1 | know what is meant by the term 'oxidation number' and understand the rules for assigning oxidation numbers | Why we need oxidation numbers The rules Writing oxidation numbers correctly |
| 8.2 | be able to calculate the oxidation number of elements in compounds and ions, including in peroxides and metal hydrides | A method that always works Two exceptions: peroxides and metal hydrides |
| 8.3 | be able to indicate the oxidation number of an element in a compound or an ion, using a Roman numeral | Roman numerals in names |
| 8.4 | be able to write formulae given oxidation numbers | Writing formulae from oxidation numbers |
| 8.5 | understand oxidation and reduction in terms of electron transfer and changes in oxidation number, and the application of these ideas to reactions of s-block and p-block elements | Electron transfer Oxidation numbers follow the electrons Reactions of s-block and p-block elements |
| 8.6 | know that oxidising agents gain electrons and reducing agents lose electrons | The two agents |
| 8.7 | understand that a disproportionation reaction involves an element in a single species being simultaneously oxidised and reduced | Disproportionation |
| 8.8 | know that oxidation number is a useful concept in terms of the classification of reactions as redox and as disproportionation | Classifying a reaction |
| 8.9 | understand that metals, in general, form positive ions by loss of electrons with an increase in oxidation number whereas non-metals, in general, form negative ions by gain of electrons with a decrease in oxidation number | Metals and non-metals |
| 8.10 | be able to write ionic half-equations and use them to construct full ionic equations | What a half-equation is The four-step method Half-equations you will meet Combining two half-equations Simple redox equations |
| 8.11 | understand reasons for the trend in ionisation energy down Groups 1 and 2 | First ionisation energy Group 2 loses two electrons |
| 8.12 | understand reasons for the trend in reactivity of the elements down Group 1 (Li to K) and Group 2 (Mg to Ba) | Why reactivity increases down the group |
| 8.13 | know the reactions of the elements of Group 1 (Li to K) and Group 2 (Mg to Ba) with oxygen, chlorine and water | With oxygen and chlorine With water |
| 8.14 | know the reactions of: i oxides of Group 1 and 2 elements with water and dilute acid; ii hydroxides of Group 1 and 2 elements with dilute acid | Oxides with water Oxides and hydroxides with dilute acid |
| 8.14(i) | i oxides of Group 1 and 2 elements with water and dilute acid | Oxides with water Oxides and hydroxides with dilute acid |
| 8.14(ii) | ii hydroxides of Group 1 and 2 elements with dilute acid | Oxides and hydroxides with dilute acid |
| 8.15 | know the trends in solubility of the hydroxides and sulfates of Group 2 elements | Two trends in opposite directions Seeing the trends in the laboratory Why this matters |
| 8.16 | understand the reasons for the trends in thermal stability of the nitrates and the carbonates of the elements in Groups 1 and 2 in terms of the size and charge of the cations involved | Thermal decomposition Charge density and polarisation Reading the trends What each compound gives |
| 8.17 | understand the formation of characteristic flame colours by Group 1 and 2 compounds in terms of electron transitions. Students will be expected to know the flame colours for Group 1 and 2 compounds. | Where the colour comes from The colours |
| 8.18 | know experimental procedures to show: i patterns in the thermal decomposition of Group 1 and 2 nitrates and carbonates (Students will be expected to know tests for carbon dioxide and oxygen; and to recognise nitrogen dioxide by its colour and acidic pH.); ii flame colours in compounds of Group 1 and 2 elements | What each compound gives Recognising the gases Practical: comparing the carbonates Practical: comparing the nitrates Practical: flame tests |
| 8.18(i) | i patterns in the thermal decomposition of Group 1 and 2 nitrates and carbonates (Students will be expected to know tests for carbon dioxide and oxygen; and to recognise nitrogen dioxide by its colour and acidic pH.) | What each compound gives Recognising the gases Practical: comparing the carbonates Practical: comparing the nitrates |
| 8.18(ii) | ii flame colours in compounds of Group 1 and 2 elements | Practical: flame tests |
| 8.19 | know reactions, including ionic equations where appropriate, for identifying: i carbonate ions, CO₃²⁻, and hydrogencarbonate ions, HCO₃⁻, using an aqueous acid to form carbon dioxide (and testing the gas with limewater); ii sulfate ions, SO₄²⁻, using acidified barium chloride solution; iii ammonium ions, NH₄⁺, using sodium hydroxide solution and warming to form ammonia (and testing with litmus and HCl fumes) | Carbonate and hydrogencarbonate ions Sulfate ions Ammonium ions Identifying an unknown salt |
| 8.19(i) | i carbonate ions, CO₃²⁻, and hydrogencarbonate ions, HCO₃⁻, using an aqueous acid to form carbon dioxide (and testing the gas with limewater) | Carbonate and hydrogencarbonate ions |
| 8.19(ii) | ii sulfate ions, SO₄²⁻, using acidified barium chloride solution | Sulfate ions |
| 8.19(iii) | iii ammonium ions, NH₄⁺, using sodium hydroxide solution and warming to form ammonia (and testing with litmus and HCl fumes) | Ammonium ions Identifying an unknown salt |
| 8.20 | be able to calculate solution concentrations, in mol dm⁻³ and g dm⁻³, including simple acid-base titrations using the indicators methyl orange and phenolphthalein | Concentration in mol dm⁻³ and g dm⁻³ What a titration does Indicators Good technique The five-step method When the sample has been diluted Finding a molar mass |
| 8.21 | CORE PRACTICAL 3 Finding the concentration of a solution of hydrochloric acid. | Aim and idea Method Results and calculation Evaluating the practical |
| 8.22 | understand how to minimise the sources of measurement uncertainty in volumetric analysis and estimate the overall uncertainty in the calculated result | Uncertainty in one measurement Reducing the uncertainty Comparing with the expected value |
| 8.23 | CORE PRACTICAL 4 Preparation of a standard solution from a solid acid and use it to find the concentration of a solution of sodium hydroxide. | Aim and idea Part 1: making the standard solution Part 2: titrating the sodium hydroxide |
| 8.24 | understand reasons for the trends for Group 7 elements in: i melting and boiling temperatures and physical state at room temperature; ii electronegativity; iii reactivity down the group | Melting and boiling temperatures and state Electronegativity Reactivity down the group |
| 8.24(i) | i melting and boiling temperatures and physical state at room temperature | Melting and boiling temperatures and state |
| 8.24(ii) | ii electronegativity | Electronegativity |
| 8.24(iii) | iii reactivity down the group | Reactivity down the group |
| 8.25 | understand the trend in reactivity of Group 7 elements in terms of the redox reactions of Cl₂, Br₂ and I₂ with halide ions in aqueous solution. Students are expected to know the colours of the elements in standard conditions, in aqueous solution and in a non-polar organic solvent. | Colours of the halogens Displacement reactions |
| 8.26 | understand, in terms of changes in oxidation number, the following reactions of the halogens: i oxidation reactions with Group 1 and 2 metals; ii the disproportionation reaction of chlorine with water and the use of chlorine in water treatment; iii the disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach; iv the disproportionation reaction of chlorine with hot alkali; v reactions analogous to those specified above | Oxidation of Group 1 and 2 metals Chlorine with water Chlorine in water treatment Cold, dilute alkali: bleach Hot alkali: chlorate(V) Analogous reactions |
| 8.26(i) | i oxidation reactions with Group 1 and 2 metals | Oxidation of Group 1 and 2 metals |
| 8.26(ii) | ii the disproportionation reaction of chlorine with water and the use of chlorine in water treatment | Chlorine with water Chlorine in water treatment |
| 8.26(iii) | iii the disproportionation reaction of chlorine with cold, dilute aqueous sodium hydroxide to form bleach | Cold, dilute alkali: bleach |
| 8.26(iv) | iv the disproportionation reaction of chlorine with hot alkali | Hot alkali: chlorate(V) |
| 8.26(v) | v reactions analogous to those specified above | Analogous reactions |
| 8.27 | understand the following reactions: i solid Group 1 halides with concentrated sulfuric acid, to illustrate the trend in reducing ability of the hydrogen halides; ii precipitation reactions of the aqueous anions Cl⁻, Br⁻ and I⁻ with aqueous silver nitrate solution and nitric acid, and the solubility of the precipitates in aqueous ammonia solution; iii hydrogen halides with ammonia gas (to produce ammonium halides) and with water (to produce acids) | Two reactions in one What you see Explaining the trend Silver nitrate and halide ions Hydrogen halides with ammonia and water |
| 8.27(i) | i solid Group 1 halides with concentrated sulfuric acid, to illustrate the trend in reducing ability of the hydrogen halides | Two reactions in one What you see Explaining the trend |
| 8.27(ii) | ii precipitation reactions of the aqueous anions Cl⁻, Br⁻ and I⁻ with aqueous silver nitrate solution and nitric acid, and the solubility of the precipitates in aqueous ammonia solution | Silver nitrate and halide ions |
| 8.27(iii) | iii hydrogen halides with ammonia gas (to produce ammonium halides) and with water (to produce acids) | Hydrogen halides with ammonia and water |
| 8.28 | be able to make predictions about fluorine and astatine and their compounds, in terms of knowledge of trends in halogen chemistry | Using the trends Fluorine Astatine How to answer a prediction question |