C4 Module Introduction - WikispacesChemical+Patterns+SOW.pdf · C4 Module Introduction ... o Brainstorm everything that students remember about the different topics using the headings

OCR 21st Century Science: C4 Chemical patterns

COLLINS NEW GCSE SCIENCE © HarperCollinsPublishers Ltd 2011

C4 Module Introduction

Pages 98−99 in the Student Book provide an introduction to this module.

When and how to use these pages

These pages summarise what students should already know from KS3 or from previous GCSE units and provide

an overview of the content that they will learn in this module.

o Use these pages as a revision lesson before you start the first new topic.

o Brainstorm everything that students remember about the different topics using the headings as a starting point. Compare your list with the points on page 98.

o Use the questions on page 98 as a starting point for class discussions.

o Ask students if they can tell you anything about the topics on the right-hand page.

o Make a note of any unfamiliar / difficult terms and return to these in the relevant lessons.

Suitable answers to the questions on page 98 are:

o Carbon, oxygen, sodium, magnesium

o Magnesium oxide

o The forces between the particles are stronger, so more energy is needed to break them apart

o They will move apart because of electrostatic repulsion

You could revisit these pages at the following points:

o before lesson c4_06 on Group 1 elements, pages 112−113

o before lesson c4_11 on ionic compounds, pages 122−123

Overview of module

Students begin by learning about the discoveries that led to the Periodic Table, and the way the elements are

arranged. This leads to the structure of the atom and in particular the electron arrangement in atoms. Using the

modern form of the Periodic Table, students then study two families or groups of elements, the alkali metals and

the halogens. Finally students examine ionic compounds exemplified by the compounds of alkali metals and

halogens.

Obstacles to learning

Students may need extra guidance with the following terms and concepts:

Scientific progress

In this lesson students have to take themselves back to a time before the Periodic Table was commonplace and

ideas about the nature of elements were not widely accepted. This empathy for previous ways of thinking can be

difficult for some students.

Line spectra

An understanding of line spectra requires that students understand that white light is a mixture of all the colours

and that a prism (or diffraction grating) can disperse or spread out the colours of the spectrum.

Atomic structure

Students may forget that only protons and neutrons contribute to the mass of atoms and that it is the sum of the

number of these two particles that gives the atomic mass number (referred to as the relative atomic mass here).

Some students will confuse the atomic nucleus with the nucleus of cells. Explain that they are very different uses of

the same term.

Electron arrangements

Students may have difficulty in recalling and applying the rules by which electrons fill up shells in atoms.

Element properties

Each element has unique properties and it may be difficult to see trends in properties just by taking appearance

into consideration. Students need to consider other properties, such as electrical conductivity and melting point, to

see trends.

Some students may have difficulty in recognising that while each element is different, a group has similarities in

behaviour that show a trend within the group.

The difference in appearance in the halogens may mask the similarities in the behaviour of the elements.

C4 Module Introduction continued


Element names and symbols

Some students may have difficulty recalling the symbols Na for sodium and K for potassium.

It is easy for students to confuse halogens and halides. Emphasise what the endings of the names mean.

Understanding ions

It requires imagination to move from concrete experimental data (conduction by molten salts) to the visualisation of

charged particles moving in the liquid. Students will also have difficulty understanding that the loss of a negatively

charged electron leaves a positive charge on the ion (think: minus a minus makes a positive).

Some students may struggle to understand how to balance the charges between ions in the formulae of ionic salts

(required at Higher-tier) if they are not familiar with the concept of lowest common multiples.

Practicals in this module

In this module students will do the following practical work:

o flame tests

o modelling electron shells

o investigating the properties of elements (demonstration)

o reacting the alkali metals with oxygen (optional demonstration)

o reacting the alkali metals with water (demonstration)

o reacting sodium with chlorine (demonstration)

o investigating the properties of halogens (demonstration)

o displacement reactions of chlorine, bromine and iodine

o conductivity of molten salts (demonstration)

o conductivity of solutions of salts (demonstration)

Key vocabulary covered in this module

�� element �� relative atomic mass

�� Periodic Table �� line spectrum

�� electron �� proton �� neutron �� nucleus �� orbit/shell �� proton number

�� electron arrangement �� period �� energy level

�� melting point �� inert �� trend

�� group �� alkali metal �� tarnish

�� balanced symbol equation �� state symbols

�� halogen �� diatomic molecule

�� corrosive �� toxic �� halide �� displacement reaction

�� ion �� ionic compound

�� crystal lattice



C4 Planning and collecting

Pages104−105 in the Student Book prepare students for controlled assessment.


This activity provides an opportunity to build and assess the skills that students will use when planning an

investigation.

Ask students to:

o read through the context and tasks, listing any terms that they do not understand

o as a whole class or in small groups, discuss the tasks to ensure that all students understand the terminology used and to clarify what is required

o work individually to answer the questions for each task.

If time allows, ask the students to mark one another’s work using the mark scheme provided.

Notes

This activity relates to work that students will have already met – the use of flame colours to identify elements.

Here the task is to assess the design of an experiment to identify the metal in an unknown white powder. Students

need to use the information given and their knowledge of the technique to answer the questions relating to the

experiment plan. In particular they need to identify the hazards in the experiment and assess the risks.

Answers Task 1

� Flame wire, Bunsen burner, bench mat, dish

� Wore eye protection, held the handles on the flame wire

� Concentrated hydrochloric acid – corrosive and gives off toxic fumes; Bunsen burner flame – hot, powder might

spit when put in flame

Task 2

� They cleaned the wire before each test

� They photographed the flames

� They only used a very small amount of concentrated hydrochloric acid, and wore eye protection at all times

Task 3

� Copper wire would give the flame a colour

� Risk assessment:

Hazard Precaution to reduce risk Risk

Concentrated hydrochloric acid – corrosive and toxic

vapour

Small amount put in a beaker or evaporating

dish

Wear eye protection

low

Bunsen burner – hot Place on bench mat

Close hole when not in use

low

Flame wire – gets hot Hold by the handle

Dip in acid after use

low

Solids may spit when heated Wear eye protection

Only use small amount

low

C4 Planning and collecting continued


Mark scheme

For grade E, students should show that they can:

o comment on how to work safely

o choose the basic equipment needed to carry out the flame tests

o identify the hazards in the procedure

For grades D, C, in addition show that they can:

o select the equipment needed to collect all the data needed

o identify the risks and suggest safety precautions to take

For grades B, A, in addition show that they can:

o explain why certain pieces of equipment were needed

o complete a risk assessment and suggest ways of reducing the risk in the procedure



C4 Exam-style questions

Pages 128−129 in the Student Book are exam-style questions.


These questions are based on the whole of Module C4 and cover a range of different types of questions that

students will meet in their written exams.

o The questions could be used as a revision test once you’ve completed the module.

o Work through the questions as a class as part of a revision lesson.

o Ask students to mark each other’s work, using the mark scheme provided.

o As a class, make a list of the questions that most students did not get right. Work through these as a class.

Notes on the worked example

The question is about ionic compounds. Part (a) requires students to recall or work out the symbol for a sodium ion.

Students will have a copy of the Periodic Table in the examination. Part (b) asks students to select the correct

properties that are unique to ionic compounds. There are 2 marks for this which is a good indication that two

statements of the four are correct. Students need to apply logic, i.e. all ionic compounds are crystalline but not all

crystals are ionic compounds. The last part is a 6-mark extended-answer question on the formation of an ionic

bond. The answer given is complete and well written so is worth the full marks. Students should note all the points

made in the answer.

Assessment Objectives

These exam-style questions cover the Assessment Objectives as described below.

Assessment Objectives Questions

AO1 Recall, select and communicate their knowledge and understanding of science

1c, 2b(i), 3, 4a, 4b,

5a, 5b, 5c

worked example a, b(i), c

AO2 Apply skills, knowledge and understanding of science in practical and other contexts

1a, 1b, 2a, 2b(ii), 3, 4c, 5c

worked example b(i), c

AO3 Analyse and evaluate evidence, make reasoned judgements and draw conclusions based on evidence

2b(ii), 5d

worked example b(ii)

Answers

These answers are also supplied on the Teacher Pack CD so that students can mark their own or their peer’s work.

Question

number

Answer Additional notes Mark

1a Boiling point increases down the group 1

1b 59 °C 1

1c Bromine 1

1a A, C 1 mark for each 2

2b(i)

The bromine has displaced the iodide

from the solution forming iodine

because the bromine is more reactive

than iodine

Accept ‘iodine has been formed’

Accept ‘bromine forms compounds more

readily than iodine’

1

1

2b(ii) This test is not sufficient to state that

chlorine is more reactive than bromine, it

only shows that chlorine and bromine

1

C4 Exam-style questions continued


are both more reactive than iodine

3 Relevant points:

They are both metals (malleable,

conduct electricity)

They are both soft and tarnish rapidly

They both highly reactive metals

Rubidium is softer and reacts faster

Rubidium has a lower melting point

They both have one electron in their

outer shell

Rubidium has bigger atoms / more

electron shells than sodium / its outer

electron is further from the nucleus

For 5–6 marks:

Answer demonstrates knowledge of the

Group 1 elements and understanding of

how the electron arrangements determine

the properties. All the information given is

relevant, clear, organised and presented

in a structured and coherent format.

Specialist terms are used appropriately.

Few, if any, errors in grammar,

punctuation and spelling

For 3–4 marks:

Answer states and explains either the

similarities or the differences in the

elements. For the most part the

information given is relevant and

presented in a structured and coherent

format. Specialist terms are used for the

most part appropriately. There may be

occasional errors in grammar, punctuation

and spelling

For 1–2 marks:

Answer gives some properties but the

explanations are missing or incorrect.

Answer may be simplistic. There may be

limited use of specialist terms. Errors of

grammar, punctuation and spelling may be

intrusive

For 0 marks:

Insufficient or irrelevant science. Answer

not worthy of credit

6

4a 10 1

4b Sketch showing three rings with the

correct number of electrons marked in

each (2, 8, 8 going outwards)

1

4c 22 1

5a Lithium + oxygen → lithium oxide 1

5b(i) Hydrogen 1

5b(ii) LiOH 1

5c 2Li + Cl2 → 2LiCl 1 mark for symbols Li, Cl2, 1 mark for

balancing

2

5d Lithium is hazardous / flammable /

reacts with water to produce flammable

hydrogen

It is used in small amounts / in laptops

the risk is low but increasing the amount

used / using it to power cars could make

the risk high

Or similar comparison of risks

2



C4 Module Checklist

Pages 126−127 in the Student Book provide a student-friendly checklist for revision.


This checklist is presented in three columns showing progression, based on the grading criteria. Bold italic means

Higher tier only.

Remind students that they need to be able to use these ideas in various ways, such as:

o interpreting pictures, diagrams and graphs

o applying ideas to new situations

o explaining ethical implications

o suggesting some benefits and risks to society

o drawing conclusions from evidence they have been given.

These pages can be used for individual or class revision using any combination of the suggestions below.

o Ask students to construct a mind map linking the points on this checklist.

o Work through the checklist as a class and note the points that need further class discussion.

o Ask students to tick the boxes on the checklist worksheet (on the Teacher Pack CD) if they feel confident that they are well prepared for the topics. Students should refer back to the relevant Student Book pages to revise the points that they feel less confident about.

o Ask students to use the search terms at the foot of the relevant Student Book pages to do further research on the different points in the checklist.

o Students could work in pairs, and ask each other what points they think they can do, and why they think that they can do those, and not others.

C4 Module Checklist continued


Module summary

In the introduction to this module, students were presented with a number of new ideas. Work through the list

below as part of their revision. Ask students to write their own summaries and mind maps, using this list as a

starting point.

Elements and the Periodic Table

o Scientists (Dobereiner, Newlands) searched for patterns in elements

o Mendeleev produced his successful periodic table by putting elements in order of their relative atomic mass, and he made predictions of new elements

o New elements were discovered, in particular by their line spectra

o Atoms have a nucleus containing protons and neutrons, surrounded by electrons in shells

o Only certain electron arrangements in the shells are allowed

o In the modern Periodic Table elements are in order of their proton number; element properties and electron arrangements show patterns across periods (rows) and down groups (columns)

Group 1 and Group 7

o Group 1 consists of the alkali metals; they have similar appearance and properties, and similar reactions with oxygen, water and chlorine. They get more reactive down the group

o Group 7 consists of the halogens; these are non-metals that have similar chemical properties but different appearances. They have similar reactions with alkali metals. They get less reactive down the group – shown by displacement reactions

Ionic compounds

o Group 1 and Group 7 elements form ionic compounds

o Ionic compounds are made up of positive and negative ions formed by loss and gain of electrons respectively

o The ions pack together in a fixed regular pattern forming a crystal

o When the ionic substance melts or is dissolved in water, the ions become free to move and so the substance then conducts electricity



Checklist C4 Aiming for A

Use this checklist to see what you can do now. Refer back to pages 100–125 if you’re not sure. Look across the rows to see how you could progress – bold italic means Higher tier only.

Remember you’ll need to be able to use these ideas in many ways:

� interpreting pictures, diagrams and graphs � applying ideas to new situations � explaining ethical implications � suggesting some benefits and risks to society � drawing conclusions from evidence you’ve been given.

Look at pages 300–306 for information about how you’ll be assessed.

Working towards an A grade

Aiming for Grade C �� Aiming for Grade A �� understand that scientists, such as Döbereiner and Newlands, used the properties and relative atomic masses of elements to find patterns

understand what Newlands’ table of elements showed

recall that these ideas were initially dismissed by some other scientists

explain how data accounts for or conflicts with Newlands’ table of elements

understand that each element gives a particular line spectrum of light

understand the principles that Mendeleev used to arrange his Periodic Table

understand that the agreement between Mendeleev’s predictions and later observations increased confidence in his ideas

understand that the number of protons and electrons is the same in atoms of an element and that the modern Periodic Table lists elements in the order of their number of protons

use the information in the Periodic Table to work out the number of protons, neutrons and electrons in an atom

work out the electron arrangement of an atom given the number of electrons or protons, or the Periodic Table

understand that the shells represent energy levels filled from the bottom upwards

understand that in the rows of the Periodic Table, called periods, the number of outer shell electrons increases from 1 to 8

understand that the chemical properties of elements depend on the electron arrangement of their atoms and hence their position in the Periodic Table

OCR 21st Century Science: C4 Checklist


Aiming for Grade C �� Aiming for Grade A �� understand that the group number is the number of outer shell electrons

use data on the physical properties of Group 1 to identify patterns

understand that Group 1 metals are reactive because of their single outer electron

understand that the electron arrangements of the Group 1 elements explain both their similarities and trends in their properties down the group

describe the reaction of lithium, sodium and potassium with oxygen, with cold water, and with chlorine; understand that reactions of the alkali metals become more vigorous down the group

recall hazard symbols and explain how to handle alkali metals safely

interpret balanced symbol equations for the reactions, including state symbols

balance unbalanced equations and write balanced symbol equations for the reactions, including state symbols

recall that the halogens form diatomic molecules; interpret data on their physical properties to describe patterns in the group

understand that the electron arrangements of the Group 7 elements explain both their similarities and trends in their properties down the group

understand that the halogens become less reactive down the group, as shown by their reactions with alkali metals, with iron and by displacement reactions

explain the safety precautions necessary in handling the halogens


balance unbalanced equations and write balanced symbol equations for the reactions, including state symbols

understand that ions are charged atoms that have lost (+) or gained (−) electrons

understand that ionic compounds can conduct in the molten state because the ions can move

explain the formation of the charges on the ions in compounds formed between the alkali metals and the halogens

understand why ionic compounds are non-conductors when solid but become conductors when molten or when dissolved in water

work out the formulae of ionic compounds given the symbols and charges on the ions; work out the charge on one ion given the formula of the compound and the charge on the other ion



Checklist C4 Aiming for C

Use this checklist to see what you can do now. Refer back to pages 100–125 if you’re not sure. Look across the rows to see how you could progress – bold italic means Higher tier only.

Remember you’ll need to be able to use these ideas in many ways:

� interpreting pictures, diagrams and graphs � applying ideas to new situations � explaining ethical implications � suggesting some benefits and risks to society � drawing conclusions from evidence you’ve been given.

Look at pages 300–306 for information about how you’ll be assessed.

Working towards a C grade

Aiming for Grade E �� Aiming for Grade C �� recall what is meant by the relative atomic mass of an element; understand that elements put in order may show patterns

understand that scientists, such as Döbereiner and Newlands, used the properties and relative atomic masses of elements to find patterns

understand what Newlands’ table of elements showed

recall that these ideas were initially dismissed by some other scientists

recall that some elements have distinctive flame colours

understand that each element gives a particular line spectrum of light

recall that Mendeleev produced a successful ordering of elements in his Periodic Table, allowing elements to be predicted

understand the principles that Mendeleev used to arrange his Periodic Table

recall the relative mass and charge of the proton, neutron and electron and where they are found in atoms

understand that the number of protons and electrons is the same in atoms of an element and that the modern Periodic Table lists elements in the order of their number of protons

understand that electrons are arranged in shells, and interpret diagrams showing the arrangements

work out the electron arrangement of an atom given the number of electrons or protons, or the Periodic Table

understand that in the rows of the Periodic Table, called periods, the number of outer shell electrons increases from 1 to 8

OCR 21st Century Science: C4 Checklist


Aiming for Grade E �� Aiming for Grade C �� understand that groups of elements with similar chemical properties form a column in the Periodic Table

understand that the group number is the number of outer shell electrons

recall the names and symbols of the top three elements in Group 1 (the alkali metals)

use data on the physical properties of Group 1 to identify patterns

understand that Group 1 metals are reactive because of their single outer electron

describe the reaction of lithium, sodium and potassium with oxygen, with cold water, and with chlorine; understand that reactions of the alkali metals become more vigorous down the group

recall hazard symbols and explain how to handle alkali metals safely

recall the word equations for the reactions of alkali metals with water, and with chlorine


recall the names and appearance of the elements in Group 7 (the halogens) and recognise their symbols

recall that the halogens form diatomic molecules; interpret data on their physical properties to describe patterns in the group

understand that the halogens become less reactive down the group, as shown by their reactions with alkali metals, with iron and by displacement reactions

explain the safety precautions necessary in handling the halogens

recall the word equations for the reactions of halogens with alkali metals and with iron


recall the properties of compounds of the alkali metals with the halogens, and understand that they are called ionic compounds because they consist of ions

understand that ions are charged atoms that have lost (+) or gained (−) electrons

understand that ionic compounds can conduct in the molten state because the ions can move

understand that ionic compounds are crystalline because the ions form a regular crystal lattice

understand why ionic compounds are non-conductors when solid but become conductors when molten or when dissolved in water



c4_01 Looking for patterns

Resources

Student Book pages 100−101 � Homework pack c4_01

Files on Teacher Pack CD: c4_01_worksheet; c4_01_technician

Materials for demonstration

Learning outcomes C4.1.3 understand that early attempts to find connections between the chemical properties of the elements and

their relative atomic mass were dismissed by the scientific community

C4.1.4(part) recall the significant stages in the history of the development of the Periodic Table to include the

ideas of Döbereiner and Newlands

Ideas about Science IaS 3.1 scientific hypotheses, explanations and theories are not simply summaries of the available data – they are

based on data but are distinct from them

IaS 3.2 an explanation cannot simply be deduced from data, but has to be thought up creatively to account for the

data

IaS 3.3 a scientific explanation should account for most (ideally all) of the data already known – it may explain a

range of phenomena not previously thought to be linked, and it should also enable predictions to be made about

new situations or examples

IaS 4.1 scientists report their claims to other scientists through conferences and journals – scientific claims are

only accepted once they have been evaluated critically by other scientists

IaS 4.2 scientists are usually sceptical about claims that cannot be repeated by anyone else, and about

unexpected findings until they have been replicated (by themselves) or reproduced (by someone else)

Literacy focus: Reading about scientists’ ideas.

Numeracy focus: Understanding that relative atomic mass is a comparison between the masses of atoms.

ICT focus: Accessing information about scientists from the internet.

In this lesson students are learning to:

� understand that scientists used creative ideas in looking for patterns in elements

� explain how other scientists learned about and reacted to the new ideas

� understand that explanations are unlikely to be accepted if they cannot account for new evidence.

Key vocabulary

element �� relative atomic mass


We are accustomed to science having all the correct answers, but in this lesson students have to take themselves

back to a time before the Periodic Table was everywhere and ideas about the nature of elements were not widely

accepted. This ‘empathy’ for previous ways of thinking can be difficult for some students.

Stimuli and starter suggestions

� Review what students know about elements. Ask questions to elucidate the definition of an element, the number

of elements known, the symbols and properties of some. Note that, as of 2011, 112 elements have been named,

with a number still awaiting confirmation. Only 90 occur naturally.

Learning activities worksheet c4_01 Low demand � The worksheet provides information about scientists to build into a timeline. More information is

given in the Student Book. Note the period of time covered and discuss how scientists found out about the ideas of

other people (books, letters and journals). Explain that scientists investigated the properties of substances,

including elements, and that (from the time of Dalton) an important property was the relative atomic mass of the

element. Use models, e.g. balls of different size and mass, to explain the comparison of the mass of atoms.

Element cards showing name, symbol, relative atomic mass, and date of discovery would be useful.

Teaching and learning notes: Students need to understand that relative atomic mass is a property unique to each

element.

c4_01 Looking for patterns continued


Standard demand � Discuss what was known about elements in the early 19th century (about 30–40 elements

were known then, and scientists were only just beginning to understand what an ‘element’ was). Note that some

scientists looked for patterns while others dismissed ideas that elements were related to each other (trying to avoid

the mistakes of the old ‘four-elements’ theory). Discuss the evidence for each side, e.g. Döbereiner’s triads versus

discoveries of new and different elements. Show samples of lithium, sodium and potassium; and carry out the

demonstration. (Health and Safety: see the Technician sheet c4_01.) Point out the similarities in the properties of

the metals, and show that their relative atomic masses obey Döbereiner’s rule, i.e. the relative atomic mass of Na

is the average of the relative atomic masses of Li and K.

Discuss the place of imagination and creativity in devising new hypotheses (e.g. de Chancourtois) and the role of

peer review in acceptance of new theories. Do not mention Mendeleev at this stage. The worksheet invites

students to carry out further research on aspects of this story and provides additional tasks. Students could work in

groups and report their findings to the rest of the class in some form of presentation.

High demand � Students carry out a study of John Newlands’ work on the Law of Octaves and the response of the

Chemical Society of London (see Student Book p. 101). They should explore the reasoning behind Newlands’ idea

and the reasons for its dismissal by other chemists. Students could report the controversy as an article in a

contemporary newspaper or magazine. The worksheet provides guidance for this activity and additional tasks.

Plenary suggestions Tell the students that the date is 1869. About 60 elements are now known. Propose a vote on whether there is or is

not a pattern relating the elements to each other. Ask students to make statements in support of both sides. There

is no correct answer at this date.

Student Book answers Q1 From his book; Q2 7

Q3 Cl (35.5), Br (80), I (127); average (35.5 + 127)/2 = 81, so prediction almost correct

Q4 Each element had a unique set of properties. The differences between the elements obscured the similarities.

The pattern was too complex to see.

Q5 For: Li, Na, K are similar in a row (also G=Be, Mg, Ca). Against: O, S, Fe are not similar (nor are N, P, Mn).

Q6 Newlands did not leave any spaces for the newly discovered elements.

Worksheet answers Activity 1 (Low demand)

Q1 … mass … atom … lightest … carbon … elements … 12 … twice … relative

Q2 Aristotle 4th century BCE, Boyle 1661, Lavoisier 1789 , Dalton 1803, Döbereiner 1829, Newlands 1864,

Sigurd Hofmann present day

Activity 2 (Standard demand)

Q2 a) They each used data on relative atomic masses and observations of the reactions and properties of

elements, but the use of imagination to see the patterns was also important, particularly perhaps for Newlands,

who related the pattern in the elements with the musical scale.

b) Printing was the only way that scientists across Europe (and elsewhere) could share knowledge and ideas.

Scientists either wrote books detailing their ideas or contributed papers to journals.

c) The response of other scientists was important to share and disseminate ideas: Döbereiner was ignored,

de Chancourtois was not understood (because of the missing diagram) and Newlands was ridiculed.

Activity 3 (High demand)

Q2 a) Relative atomic mass

b) He thought similar elements recurred every eighth element – an octave.

c) No. 3, G = beryllium Be; No. 4, Bo = boron B

d) Li, Na, K are similar and obey Döbereiner’s rule (23 = (7 + 39)/2)

H, F, Cl are almost correct, although H is not very similar to F and Cl; but generally no.

e) It is almost correct for H, F, Cl, for Li, Na, K, and for Be, Mg, Ca; but then breaks down.

f) H, F, Cl (H not as reactive); C, Si, Ti (Ti a typical metal); N, P, Mn (Mn metallic); O, S, Fe; generally it breaks

down in the later third (and fourth) columns.

g) Cr and Ti

h) To make room for extra elements, Newlands would have to shift all the later elements, destroying whatever

pattern there was.




Technician sheet

Equipment and materials

Demonstration of Döbereiner’s triads

� samples of lithium, sodium, potassium

� white tile

� sharp knife

� tongs/tweezers

Method

Use the tongs/tweezers to take a lump of the metal out of the storage bottle. Cut a sliver of the metal from the lump on the tile.

Show students that the metal is shiny silver when freshly cut. Use a paper towel to wipe off the oil used to protect the metal in storage. Show students how the surface of the metal tarnishes rapidly by reaction with oxygen in the air.

Repeat with the other metals.

Notes

� If you wish you could demonstrate the reaction of the metals with cold water. However, this is covered in more detail in a later lesson, and instructions are given on Technician sheet c4_06.

Health and Safety

� If a member of CLEAPSS, refer to Hazcards 58, 88 and 76.

� Lithium, sodium and potassium are FLAMMABLE and REACT VIOLENTLY with water.

� The metals should not be allowed to come into contact with skin.

� The oil used to protect the metals in storage is FLAMMABLE.

� Use prepared pieces of the metals with sides no larger than 4 mm.

� Never attempt to constrain the metals; all them to roam freely on a large surface of water.

� Place the safety screen as close to the trough as possible to avoid pieces flying out over the class.

� The behaviour of the metals is very unpredictable, and they may explode and eject particles. In place of safety screens, use an acrylic sheet can be placed on top of the trough.




1 Ideas about elements

1 Fill in the missing words in the following sentences. The words are given below. Each word should be used only once.

carbon mass twice lightest atom relative elements 12

The relative atomic mass of an element is the ……………… of its atoms compared to

the mass of another ……………… Originally atoms were compared to the mass of the

……………… element, hydrogen. A ……………… atom has the mass of twelve

hydrogen atoms. Today we compare the mass of atoms of other ……………… with

the mass of carbon atoms, which are given the value ……………… A magnesium atom

is ……………… the mass of a carbon atom, so has a ……………… atomic mass of 24.

2 Find out when the people mentioned in the boxes below were alive. Cut out the boxes

and arrange them as a timeline. You could look for extra information to add to the timeline on the internet or in books.

Antoine Lavoisier writes Elements of Chemistry. He says oxygen is an element. He lists 33 elements and proves that the mass of products of a reaction is the same as the mass of the reactants.

Robert Boyle writes The Sceptical Chemist. In the book he says elements are simple substances that cannot be broken down.

Greek philosophers such as Aristotle think that all substances are made up of Earth, Air, Fire and Water.

Johann Döbereiner reports his theory of ‘triads’ of elements that have similar properties and a pattern in their relative atomic masses.

John Dalton writes a book containing his theory of atoms. He lists the relative atomic masses of 20 elements.

John Newlands describes his Law of Octaves showing a pattern in the elements when arranged in order of their relative atomic mass. The members of the Chemical Society of London think it is nonsense.

The 112th element to be discovered, by Sigurd Hofmann and his team, is given the name copernicium.



2 Evidence and patterns in elements

1 Investigate the contribution made by certain scientists to knowledge and theories about elements. For each scientist, note the date when the work was done or published, what the scientist suggested or discovered, how other scientists found out about the ideas and what the response was. Some of the scientists involved in the development of ideas about elements were Robert Boyle, Antoine Lavoisier, John Dalton, Johann Döbereiner, Alexandre-Emile de Chancourtois and John Newlands.

2 a) Which was more important in the work of Döbereiner, Chancourtois and Newlands: evidence from observations and measurements, or the use of imagination to see patterns in behaviour? Explain your answer.

b) How important was the printing of books, magazines and journals in spreading ideas about elements?

c) What part did other scientists play in approving or disapproving of new ideas?

3 John Newlands and the Law of Octaves

1 Work in pairs to investigate the ideas of John Newlands and the response to them by the members of the Chemical Society of London in 1866.

One person should look at Newlands’ proposals and the evidence he presented to support them. The other should examine why the Chemical Society generally rejected Newlands’ ideas. Then combine your work to produce a report for The Times newspaper of 1866.

Things to think about:

• What did Newlands’ Law of Octaves suggest?

• Why did Newlands arrange the elements in the pattern that he presented to the meeting?

• What evidence was there for the pattern?

• Was Newlands known to the members of the Chemical Society of London before the meeting?

• Was the use of a musical term (octaves) a good or bad way of Newlands getting his ideas across to other chemists?

• Could Newlands’ ideas be confirmed by performing experiments to test predictions?

• Were there any other faults with Newlands’ idea?

2 Look at Newlands’ table of the elements and answer the questions that follow.



In John Newlands’ table of elements the numbers next to each symbol are simply the number order of the elements.

a) What data did Newlands use to put the elements in order?

b) Why did Newlands put seven elements in each column?

c) Some elements in the table do not have symbols that we recognise today. Which elements have changed their symbols?

d) Do Döbereiner’s triads feature in Newlands’ table? Give examples.

e) Chancourtois’ spiral table suggested that similar elements were separated by 16 relative atomic mass units. Is this pattern obeyed in Newlands’ table?

f) Give examples of where the similarity between elements in horizontal rows breaks down.

g) Give an example of elements in the wrong order in Newlands’ table due to inaccuracies in measuring relative atomic masses.

h) Vanadium (relative atomic mass 51) was discovered in the early 19th century but Newlands left it out of his table. What would happen to the table when vanadium and other newly discovered elements were added to it?



c4_02 Finding new elements

Resources

Student Book pages 102−103 � Interactive Book: Matching pairs ‘Spectroscopy’ � Homework pack c4_02

Files on Teacher Pack CD: c4_02_worksheet; c4_02_practical; c4_02_technician

Equipment for class practical

Learning outcomes C4.2.6 recall that some elements emit distinctive flame colours when heated (for example lithium, sodium and

potassium) (recall of specific flame colours emitted by these elements is not required)

C4.2.7 understand that the light emitted from a particular element gives a characteristic line spectrum

C4.2.8 understand that the study of spectra has helped chemists to discover new elements

C4.2.9 understand that the discovery of some elements depended on the development of new practical

techniques (for example spectroscopy)

C4.1.4(part) recall the significant stages in the history of the development of the Periodic Table to include the

ideas of Mendeleev

C4.1.5 understand how Mendeleev used his Periodic Table to predict the existence of unknown elements

Ideas about Science As for lesson c4_01.

Literacy focus: Writing an appreciation of Mendeleev’s work.

ICT focus: Accessing images of flame colours and line spectra from the internet.


� remember that some elements produce distinctive flame colours

� explain how flame colours led to the discovery of new elements

� understand that Mendeleev’s Periodic Table was accepted because it made successful predictions.

Key vocabulary

Periodic Table �� line spectrum


An understanding of line spectra requires that students understand that white light is a mixture of all the colours

and that a prism (or diffraction grating) can disperse or spread out the colours of the spectrum.


� Issue cards with information about elements (appearance, relative atomic mass, simple properties) but no names

(or made-up names) and ask students to put them in order in rows and columns as Mendeleev did. (Miniature

examples are given on the worksheet.) Give students a few minutes, and then discuss whether the task is easy

or difficult and whether they have enough information.

Learning activities worksheet c4_02 + practical c4_02 Low demand � Explain how Mendeleev published his new Periodic Table in 1871 and in particular that he

predicted that some elements would exist that had not been discovered. Discuss how new elements could be

discovered. Explain that it was not necessary to isolate a sample of the new element but just to show that it was

present in compounds and mixtures. Note that particular elements have distinctive flame colours (see Student

Book p. 102). Explain that chemists saw new colours in new minerals that revealed previously unknown elements.

Some of these proved to be the elements that Mendeleev predicted.

Teaching and learning notes: Students need to understand the reason why Mendeleev’s Table was successful

and the importance of flame tests in identifying elements.

Standard demand � Discuss how new techniques resulted in discoveries of new elements (some are mentioned in

the Student Book p. 102). (Methods for liquefying and distilling gases and radioactivity measurements were

responsible for later bursts of discovery.) Explain that looking at flame colours and the technique of spectroscopy

provided a new method in the 1860s and 1870s. Students should carry out the flame tests described on the

practical sheet and if possible use handheld spectrometers to observe the line spectra. If these are not available,

use pictures to reveal the different spectra of elements. Explain how the prism/diffraction grating splits the light and

that every element produces a unique pattern of lines. Hence discuss how the discovery of new elements aided

c4_02 Finding new elements continued


Mendeleev in his bid to get his Periodic Table accepted. It was only after the first of his predictions was proved

correct that chemists began to accept that he was on to something. The worksheet gives further tasks.

Teaching and learning notes: Students will need access to line spectra of the various elements and copies of

Mendeleev’s Periodic Table showing where the elements he predicted would fit (see Student Book p. 103).

High demand � Discuss the rules that Mendeleev adopted to produce his table, the evidence he used (particularly

the pattern in the formula of oxides, e.g. Na2O, MgO, Al2O3, SiO2, P2O5, SO3, Cl2O7), and the predictions he made

(to include various new elements and that some measurements of relative atomic masses would prove to be

inaccurate and hence conform to the order of the elements that he suggested). Note at that time ‘atomic number’

was simply the place of the element in order in the Periodic Table. Discuss how discoveries that Mendeleev had

not predicted also bolstered his Periodic Law, i.e. the discovery of the ‘noble gases’. Point out that the modern

Periodic Table is not Mendeleev’s – there were problems with his table, e.g. the division of columns lower down

into subgroups of very different elements, the dumping of a number of elements into one position, and the

persistence of the ‘wrong’ relative atomic masses of tellurium (Te) and iodine (I).

Plenary suggestions Ask students to discuss the impact of Mendeleev’s Table on chemistry: it urged chemists to search for the

elements he had predicted; it showed that there was a connection between the elements and encouraged chemists

to hunt for it; and it made chemistry neat and orderly and not a jumble of unconnected facts. Students could write a

brief article ‘Why Mendeleev should have been awarded a Nobel Prize’.

Student Book answers Q1 So that only the colour of the flame of the element being studied would show up.

Q2 Mendeleev left gaps in his table for elements yet to be discovered.

Q3 New apparatus provided new ways of examining elements, and produced evidence that could not be obtained

by other methods.

Q4 Spectroscopy was used to discover rare elements, which may not have been found using the older methods.

Q5 They all form oxides with a similar formula, MO.

Q6 Tellurium/iodine Te/I (this was Mendeleev’s main problem); argon/krypton Ar/K (Ar was discovered later).


Answer should include: news of the discovery; it was found because of its unique violet flame; he called it after his

home; it is one of the elements that Mendeleev predicted and hence helps to prove the Periodic Table is true.


Q1 radium − D; oxygen − A; sodium − B; caesium − E; neon − C

Q2 Answers should include: only a small amount needed, every element has a unique line spectrum, new lines

can be seen even in mixtures of elements, and little preparation is required (no long extractions are needed).

Q3 Points may include: encouraged search for new elements, brought order to chemistry, made it easier to learn

about elements.


Q1 a) Similarities: elements listed in order of relative atomic mass; seven rows/columns of elements.

Differences: Newlands did not leave gaps; and had elements with different properties in same group.

b) Predictions could be tested and found to be correct, which built confidence in Mendeleev’s table. It coped

with new data.

Q2 a) Places where Sc, Ga, Ge were found. b) Te/I, Ar/K

c) They just added an extra column, with no alterations needed to others.

d) He split columns into two groups (A and B subgroups).

Practical sheet answers 1 Answer depends on students’ observations of the salts they were given.

2 red – lithium, yellow – sodium, lilac/mauve – potassium, calcium – orange, copper – green

3 The salt contains an element that has not be detected before.

4 Many elements have similar coloured flames but the line spectra are unique for each element so allow the

element to be identified.



c4_02 Identifying elements

P Flame tests

Objectives

In this activity you will:

� observe the colours of flames of various elements

� identify an unknown element from its flame colour

� observe the line spectrum of some elements.

You must wear eye protection when carrying out this activity.

The acid used to clean the flame wire is CORROSIVE.

Some of the compounds used in the tests are IRRITANT. Avoid touching them or inhaling them.


flame wires • watch glass or small dish • Bunsen burner • bench mat • handheld spectrometer (optional) •

hydrochloric acid • samples of salts to test

Method

1 Pour a little hydrochloric acid into a small dish. Light the Bunsen burner and turn it to a blue flame.

2 Heat the flame wire in the Bunsen flame. Dip the hot wire into the acid and then hold it in the flame again. Repeat this until the wire gives no colour to the flame.

3 Dip the wire into the acid and then into a sample of one of the salts. Hold the wire and salt in the Bunsen flame and record the colour that you see.

4 (optional) Look through the spectrometer at the coloured flame and describe what you see.

5 Clean the wire as described in part 2, then repeat part 3 with a different salt.

6 Finally, use a clean wire to test the unknown salt. Record the colour that you see.

Results

Record the names of the salts tested and the observations that you made in a suitable table.

Questions

1 What are the flame colours for the metals that you tested?

2 What metal was in the unknown salt?

3 What would it mean if, for the unknown salt, you saw a flame colour that did not match any of the colours that had been seen before?

4 (optional) Why are line spectra more useful than just looking at the flame colours for identifying elements?



c4_02 Identifying elements

Technician sheet


Flame tests

Each group of students will need:

� flame wires – pieces of nichrome wire about 20 cm long with one end bent into a large loop as a handle and the other end in a small loop will do instead of proprietary flame wires in glass handles (or wires can be fixed in corks to be used as handles)

� Bunsen burner

� bench mat

� watch glass or small dish/beaker

� (optional) handheld spectrometer (not for basic level students)

� hydrochloric acid (2 mol/dm3)

� samples of test solids – lithium chloride, sodium chloride, potassium chloride, calcium chloride, copper chloride or copper sulfate

� an ‘unknown’ salt (suggest sodium chloride or potassium chloride)

Method

Full instructions are given on practical sheet c4_02.

Notes

� Clean Bunsen burners are a necessity.

� Students only need a tiny amount of each solid to test.

� The hydrochloric acid will need replacing after each test.

� Students should be told to leave testing sodium salts to last, as sodium produces a very intense flame colour that can colour Bunsen flames for some time and prevent other colours being seen.

� If students are working in pairs, then one can hold the sample in the flame while the other uses the spectrometer to observe the line spectrum.

� Observing the line spectrum can be difficult, and if there are not enough spectrometers to go round, this can be done as a demonstration at the end of the practical session.

� There are alternative ways of demonstrating flame colours using alcoholic solutions that are more dramatic. For example, see the ‘Practical Chemistry’ website. It is important to practise this before the lesson to find a safe distance for the students.

Health and Safety

� Hydrochloric acid is CORROSIVE. Only a few drops are needed at a time.

� Lithium and copper salts are HARMFUL.

� Students should wear eye protection and avoid contact with the acid and the salts.

� If a member of CLEAPSS, refer to Hazcards 47A, 58 and 27A.

OCR 21st Century Science: C

4 Chemical patterns

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c4_02 Findin

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S Starte

r activity: sortin

g elements

Cut o

ut th

e following element cards. Shuffle each set. Sort th

e cards into order of relative

atomic mass and place them into a table with seven columns.

Aquagen

1

Flammable

gas

Stonium

7

Reactive

metal

Mabelium

9

Fairly reactive

metal

Yawnium

11

Semi-metal

Burnon

12

Black solid

that burns

Azote

14

Unreactive

gas

Sourgen

16

Reactive

gas

Lumenine

19

Highly reactive

gas

Saltium

23

Reactive

metal

Epsomium

24

Fairly reactive

metal

Hallium

27

Shiny

metal

Quartzon

28

Semiconductor

looks metallic

Glowon

31

Flammable solid

non-metal

Brimstone

32

Yellow solid

that burns

Greenine

35.5

Highly reactive

gas

Kalium

39

Reactive

metal

Limeium

40

Fairly reactive

metal



1 Discovering gallium

Imagine you are Pierre Lecoq de Boisbaudran, a French chemist. In 1875 he discovered one of the new elements that Mendeleev predicted. It colours flames a deep violet. He called the new element ‘gallium’ after his home (‘Gaul’ is the old name for France).

Write a letter to Mendeleev describing your discovery, how you did it and why you think it is important.

2 Finding new elements

1 Match up the name of the element with the technique used to discover it.

radium A Collected using glass jars and tubes in the 1770s

oxygen B Obtained by electrolysing salt by Davy in 1807

sodium C Extracted by liquefying air by Ramsay in 1898

caesium D Radioactive element detected by Curie in 1898

neon E Discovered by its line spectrum by Bunsen in 1861

2 Imagine you are a chemist in the 1870s keen to discover new elements. Write a letter to a friend saying why looking at the line spectrum of compounds will be a help in your quest.

3 Design a presentation or poster to convince the Nobel committee that Mendeleev should be awarded the prize.



3 Probing Mendeleev

1 Compare Mendeleev’s Periodic Table with Newlands’ table from his Law of Octaves.

a) What are the similarities and differences?

b) Why was Mendeleev’s table eventually accepted while Newlands’ was rejected?

2 Mendeleev’s main rules for designing his Periodic Table were:

• the elements must be in order of their relative atomic mass

• elements with similar properties must be in the same column.

Compare Mendeleev’s Periodic Table (below) with the modern table. Answer the questions that follow.

© Science Photo Library

a) Give two examples of where Mendeleev left gaps for elements that had not yet been discovered.

b) Give an example of where Mendeleev’s first rule is broken.

c) What effect did the discovery of the gases helium, neon, argon, krypton, xenon and radon in the 1890s have on Mendeleev’s Periodic Table?

d) How did Mendeleev squeeze the elements into seven main columns (plus an extra dumping ground that he called Group VIII)?



c4_03 Building atoms

Resources

Student Book pages 106−107 � Interactive Book: Quick starter ‘Atomic odd one out’; Quick starter ‘Can you unjumble the sentences to find out about the particles in an atom?’ � Homework pack c4_03

Files on Teacher Pack CD: c4_03_worksheet

Learning outcomes C4.1.1 understand that atoms of each element have different proton numbers

C4.1.2 understand that arranging the elements in order of their proton numbers gives repeating patterns in the

properties of elements

C4.1.6 use the Periodic Table to obtain the names, symbols, relative atomic masses and proton numbers of

elements

C4.2.1 describe the structure of an atom in terms of protons and neutrons in a very small central nucleus with

electrons arranged in shells around the nucleus

C4.2.2 recall the relative masses and charges of protons, neutrons and electrons

C4.2.3 understand that in any atom the number of electrons equals the number of protons

C4.2.4 understand that all the atoms of the same element have the same number of protons

C4.2.5 understand that the elements in the Periodic Table are arranged in order of proton number

C4.2.10 use the Periodic Table to work out the number of protons, electrons and neutrons in an atom

Literacy focus: Reading about the discoverers of atomic structure.

Numeracy focus: Developing a sense of scale in the context of atomic structure.

ICT focus: Using spreadsheets to store data on the structure of atoms and calculate numbers of protons, neutrons

and electrons.


� remember the properties of the particles that make up atoms

� explain how atoms are arranged in the Periodic Table.

Key vocabulary

electron �� proton �� neutron �� nucleus �� orbit/shell �� proton number �� relative atomic mass


Students may forget that only protons and neutrons contribute to the mass of atoms and that it is the sum of the

number of these two particles that gives the atomic mass number (referred to as the relative atomic mass here).

Note that there is no reference to isotopes in this specification.

Some students will confuse the atomic nucleus with the nucleus of cells. Explain that they are very different uses of

the same term.


� Display the Periodic Table, or students can use p. 311 of the Student Book. Ask students to find named

elements.

Learning activities worksheet c4_03 Low demand � Assess students’ sense of scale. Ask them to name things that are 1–2 m in size (humans,

bicycles), 1–2 mm (grains of sand), 1–2 µm (cells), 1–2 nm (large molecules). Show pictures at each level of scale.

Note each of these are a thousand times smaller. Explain that atoms are 10 times smaller again (0.1 nm) and that

they are made up of particles 100 000 times smaller.

Describe the three particles that make up atoms and explain how they are arranged. Note that the diameter of the

largest nucleus is about 1/10 000 of the atom. Most diagrams exaggerate the size of the nucleus. It is relatively

about the size of a pea at the centre of Wembley football stadium (including all the stands).

The worksheet provides tasks to help students remember the particles and their properties.

Teaching and learning notes: Students need to recall the properties of the particles in atoms.

c4_03 Building atoms continued


Standard demand � Discuss the rules of atoms and the Periodic Table: all the atoms of an element have the same

proton number; in atoms the number of electrons is equal to the number of protons; elements are arranged in

proton number order in the Periodic Table; the Periodic Table gives the name, symbol, proton number (lower

number) and relative atomic mass (upper number) of the element.

To develop an understanding of these ideas, students may investigate the work of the scientists involved in

discoveries related to atomic structure. The worksheet gives some background to this activity. Students could

design a spreadsheet to display data obtained from the Periodic Table.

Teaching and learning notes: Students will need access to the Periodic Table (p. 311 of the Student Book).

High demand � Higher-tier students are expected to be able to calculate the numbers of particles in atoms given

the data from the Periodic Table. The number of neutrons is found by subtracting the number of protons from the

atomic mass number. The Periodic Table provided in the specification and in the Student Book (p. 311) gives the

relative atomic mass. Students may use this instead of the atomic mass number. Students may design a

spreadsheet to calculate the data for them.

Students who question why chlorine has a relative atomic mass of 35.5 should be told about the existence of

isotopes. Similarly, there are no bromine atoms with 45 neutrons (the relative atomic mass of 80 is the average for

the isotopes of bromine). If students see other versions of the Periodic Table, they may notice that most relative

atomic masses are not whole numbers. This is also largely explained by the existence of isotopes of most elements

(nuclear binding energy also results in relative atomic mass not being a whole number).

Plenary suggestions Ask students to pick out elements from the Periodic Table by name, symbol or proton number, and state the

number of protons and electrons. Higher-tier students should also be able to use the relative atomic mass to give

the number of neutrons.

Student Book answers Q1 a) proton b) electron c) proton

Q2 a) 6 b) 7

Q3 sodium, Na, 23

Q4 a) p 2, n 2, e 2 b) p 16, n 16, e 16 c) p 19, n 20, e 19 d) p 26, n 30, e 26


Q1 a) photo 28.5 mm; diagram 57 mm b) 2 c) about 100

d) The nucleus is much too small to be seen. e) 10 000 cm = 100 m

Q2

Particle Charge Mass Position in atom

proton +1 same as hydrogen atom in nucleus at centre of atom

electron –1 almost zero in shells around nucleus

neutron 0 same as hydrogen atom in nucleus at centre of atom


Q2 John Dalton – Idea that atoms are hard particles that cannot be split, and that atoms of each element are

different.

Dmitri Mendeleev – The Periodic Table, and predictions of the existence of further elements

J. J. Thomson – Discovered the electron, devised the plum pudding model of atom

Ernest Rutherford – Evidence for atoms having a tiny positively charged nucleus surrounded by electrons,

discovered the proton

Henry Moseley – Found that an element’s properties depended on the atomic or proton number

Niels Bohr – Suggested that the electrons were in orbits or shells around the nucleus

James Chadwick – Discovered the neutron



c4_03 Building atoms

1 The scale of atoms and the atomic nucleus

1

Photo of a £2 coin (actual size).

Diagram of a £2 coin to scale.

a) Use a ruler to measure the diameter of the coin in the photo and in the diagram.

………………… …………………

b) How many times bigger is the diagram of the coin compared to its actual size?

………………… This is called the magnification scale of the diagram.

c) A human hair is about 0.1 mm thick. What is the magnification scale of the photo?

…………………

Photo of a human hair magnified.

d) Why can’t you see the nucleus in the diagram of the atom?

………………………………………………………………………………………………

Diagram of an atom with a magnification scale of about 100 000 000 (one hundred million).

e) The diameter of an atom is about 10 000 (ten thousand) times the diameter of the atomic nucleus. How big would the atom be if the nucleus was 1 cm in diameter?

…………………

c4_03 Building atoms continued


2 Fill in the gaps in the table below for the particles found in atoms.

Particle Charge Mass Position in atom

proton same as hydrogen atom

electron –1

0 in nucleus at centre of atom

2 Discovering atomic structure

1 Design a spreadsheet to show:

name symbol proton number relative atomic mass number of electrons

for atoms of elements in the Periodic Table. Use a Periodic Table (page 309 of your Student Book) to fill in the spreadsheet for the first 20 elements.

2 Design a poster to show how ideas about the structure of atoms developed. The scientists named below each contributed an important idea to the development of the Periodic Table and theories of atomic structure. Carry out research to find out about their important contributions.

John Dalton – 1766–1844, British, worked in Manchester

Dmitri Mendeleev – 1834–1907, Russian, did his important in St. Petersburg

J. J. Thomson – 1856–1940, British, ran the Cavendish Laboratory at the University of Cambridge, received a Nobel Prize 1906

Ernest Rutherford – 1871–1937, New Zealander, did his important work at the Universities of Manchester and Cambridge in the UK, received a Nobel Prize 1908

Henry Moseley –1887–1915, British, worked with Rutherford at the University of Manchester, killed in the First World War

Niels Bohr – 1885–1962, Danish, first theories while working with Rutherford in Manchester but returned to Copenhagen in 1918, received Nobel Prize 1922

James Chadwick – 1891–1974, worked with Rutherford at Cambridge, received Nobel Prize 1935

3 Counting particles

1 Design a spreadsheet to show

name symbol proton number relative atomic mass

and to calculate

number of protons number of neutrons number of electrons

for atoms of elements in the Periodic Table. Use a Periodic Table (page 309 of your Student Book) to fill in the spreadsheet for the first 20 elements.



c4_04 Arranging electrons

Resources


Files on Teacher Pack CD: c4_04_worksheet

Small balls of two colours

Learning outcomes C4.2.11 use simple conventions, such as 2.8.1 and dots in circles, to represent the electron arrangements in the

atoms of the first 20 elements in the Periodic Table, when the number of electrons or protons in the atom is given

(or can be derived from the Periodic Table)

C4.2.12 understand that a shell (or energy level) fills with electrons across a period

ICT focus: Using models of electron arrangements in atoms.


� describe how electrons are arranged in shells around the nucleus of an atom

� explain how the electron arrangement of elements is related to the Periodic Table.

Key vocabulary

electron arrangement �� period �� energy level


Students may have difficulty in recalling and applying the rules by which electrons fill up shells in atoms.


� Test the recall of the properties of the particles in the atom from the previous lesson.

Learning activities worksheet c4_04 Low demand � Explain how electrons fill the shells in atoms. Tell students that we can use models to show this. A

model can be a simple numerical pattern, e.g. 2.8.1, or a diagram (show examples of shells as rings with electrons

as dots or crosses), or other models (e.g. 3D). Explain that the model has some of the qualities of the real thing, i.e.

in this case the arrangement of electrons, but is not the real thing. A useful model is to use students to represent

electrons. Designate one student as the nucleus and give them coloured ball(s) to represent proton(s). Start by

modelling a hydrogen atom. Give another student a different coloured ball to represent the negative charge of the

electron. Add more proton balls and more electron students to build up larger atoms. The students can move in

circles around the nucleus.

The worksheet gives further tasks.

Teaching and learning notes: Students need to know the rules. Rule 1: start with first shell closest to the nucleus;

when it is full, go to second shell and so on. Rule 2: first shell can take up to two electrons, second shell eight

electrons, and third shell eight electrons (up to element number 20). Students may say that the third shell is full

when it has eight electrons in it. This is incorrect, but is accepted at GCSE.

Standard demand � Having described the arrangement of electrons in atoms (see above), explain how the

Periodic Table is arranged according to the electron arrangement of atoms. The first row/period (H, He) has just

one shell occupied; the second row/period has the second shell occupied and so on. The shell fills up from left to

right across the Periodic Table. The number of the column (group) is the number of electrons in the outer occupied

shell. Hence, pick examples from the first 20 elements in the Periodic Table, give students the proton number and

ask them to work out the electron arrangement.

The worksheet gives further tasks to reinforce this.

Teaching and learning notes: It is time-consuming but nevertheless useful for students to draw diagrams of the

electron arrangements of the first 20 elements.

High demand � Explain that electrons in different shells have different energies, and that an energy level diagram

is an alternative way of describing the electron arrangement. Note that the first shell has the lowest energy level.

Students should be able to calculate electron arrangements and draw energy level diagrams for any of the first 20

elements when given the Periodic Table and name of an element. Students may like to know that Niels Bohr was

one of the founders of the science of quantum mechanics, and his idea that electrons could only have the energies

c4_04 Arranging electrons continued


of the shells in the atoms (and no energy in between) gave rise to the whole weird and wonderful world of the

quantum.

Plenary suggestions Write out the names and proton numbers of the first 20 elements in the Periodic Table on slips of paper. Give each

pupil a slip and a sheet of paper. Ask them to sketch the electron arrangement of the element they have been

given (no names or symbols). Split the class into teams. In turn, members of the team hold up their diagrams and

the other team(s) have 10 seconds to identify the element using a Periodic Table.

Student Book answers Q1 a) 1st, 2; 2nd, 8; 3rd, 1

b) 11

Q2 Diagram showing two electrons in the inner circle and four in the outer circle.

Q3 a) 2.1 b) 2.8 c) 2.8.2 and diagrams to match

Q4 2.8.5 and diagram to match

Q5 Third period, last element on right. Appropriate energy level diagram. Proton number = 18.

Q6 Proton number = 20. Four shells, so in the fourth row/period. Two electrons in outer shell, so is the second

element in the period (in Group 2). It is calcium, Ca.


Q1 He 2, 2; Li 2.1, 3; N 2.5, 7; Si 2.8.4, 14; Ca 2.8.8.2, 20

Q2 Appropriate diagrams for C (2.4), Ne (2.8) and P (2.8.5)

Q3 a) 2.3 b) 2.8.1 c) 2.8.8.1


Q1 Diagrams completed correctly.

Q2 a) beryllium b) fluorine c) aluminium d) argon

Q3 a) True

b) False. The first shell takes up to two electrons. The second and third shells take up to eight electrons (in the

first 20 elements).

c) False. The rows in the Periodic Table are called periods. OR The columns in the Periodic Table are called

groups.

d) False. From left to right in a period, the number of electrons in the outermost shell increases. OR From right

to left in a period, the number of electrons in the outermost shell decreases.

e) True.


Q1 Correct diagrams for the electron arrangements in Question 2 in Activity 2.

Q2 a) second row, sixth column, 2.6

b) third row, seventh column, 2.8.7

c) third row, second column, 2.8.2

d) third row, eighth column, 2.8.8



c4_04 Arranging electrons

1 Filling shells

1 a) Write down the electron arrangement for each atom underneath the diagram.

b) Write down the proton number for each element underneath the diagram.

2 Draw electron arrangement diagrams for the following atoms:

a) carbon b) neon c) phosphorus

electron arrangement electron arrangement electron arrangement

2.4 2.8 2.8.5

3 Write down the electron arrangement for the following atoms:

a) boron, proton number 5 …………………

b) sodium, proton number 11 …………………

c) potassium, proton number 19 …………………



2 Arranging the periods

1 Complete the electron arrangements for the first 18 elements in the Periodic Table.



2 Use a Periodic Table to name the elements with the following electron arrangements:

a) 2.2

b) 2.7

c) 2.8.3

d) 2.8.8

3 Decide which of the following statements are true.

Write corrections for the false statements.

a) Electron shells in atoms are filled from the inner shells to the outer shells.

b) Each shell can take up to eight electrons.

c) The columns in the Periodic Table are called periods.

d) From left to right in a period, the number of electrons in the outermost shell decreases.

e) Except for helium, the Group 0 elements have eight electrons in their outermost shell.

3 On a level

1 Draw energy level diagrams of the electron arrangements of the elements in Question 2 in Activity 2 above.

2 Find the following elements in the Periodic Table. Describe their position in the Periodic Table and write out their electron arrangements.

a) oxygen

b) chlorine

c) magnesium

d) argon



c4_05 The Periodic Table

Resources



Materials for demonstration of elements

Learning outcomes C4.1.8 recall that a period is a row of elements in the Periodic Table

C4.1.9 use the Periodic Table to classify an element as a metal or non-metal

C4.1.10 use patterns in the Periodic Table to interpret data and predict properties of elements (candidates will be

given a copy of the Periodic Table with the examination paper)

C4.2.14 (part) understand that the chemical properties of an element are determined by its electron

arrangement

Ideas about Science IaS 3.3 a scientific explanation should account for most (ideally all) of the data already known – it may explain a



Literacy focus: Describing elements.

Numeracy focus: Extracting information from the Periodic Table. Plot, draw and interpret graphs and charts from

secondary data.

ICT focus: Using databases to access data or store data on the properties of elements.


� use the Periodic Table to describe and predict the properties of elements.

Key vocabulary

proton number �� period �� melting point �� inert �� trend


Each element has unique properties and it may be difficult to see trends in properties just by taking appearance

into consideration. Students need to consider other properties, such as electrical conductivity and melting point, to

see trends.


� Show students Mendeleev’s Periodic Table again (see Student Book p. 103). Tell students that Mendeleev

predicted that an element would be discovered to fit in the space below aluminium. He predicted that it would be

a silvery metal with quite high melting point that conducted electricity, and that it would react with oxygen to form

a solid oxide. Ask students to discuss how Mendeleev was able to make these predictions.

Learning activities worksheet c4_05 Low demand � Issue copies of the Periodic Table (or display a table that all can see). Display as many elements

as possible and demonstrate their conductivity and malleability (see the Technician sheet). Ask students to

describe the elements and their demonstrated properties, find their position in the Periodic Table, and note their

symbols and proton numbers. Students could start to develop a database of information about elements. The

worksheet gives a template for an element card that could alternatively be used for this.

Ask students to discuss what they have learned and state any patterns that they have found. In particular, make

sure that they note that the non-metals are on the right of the Periodic Table. Note the position of the division line

(see Student Book p. 110). Note that, across a period, elements change from metals to non-metals.

Teaching and learning notes: Students need to recall the typical properties of metals and non-metals and the

changes from metallic to non-metallic behaviour across the Periodic Table.

Standard demand � Ensure that students collect data on the properties of the elements of period 3 (Na to Ar) (see

the Technician sheet). The worksheet gives data on the elements of period 3, which students can use to draw

charts and comment on the trends across the period. Note that the trends observed are repeated in other periods,

and so predictions can be made about the properties of elements. Emphasise that it was the ability to make

c4_05 The Periodic Table continued


predictions about undiscovered elements that was the strength of Mendeleev’s Periodic Table and the reason why

it was accepted. Note that, although we now know all the elements up to about 112, we can still use the Periodic

Table to help us recall the properties of elements and to predict the properties of new artificial elements.

Teaching and learning notes: Students will need to be able to use and interpret data from bar charts and line

graphs of the physical properties of elements across a period. There are various interactive Periodic Tables that

provide data on elements.

High demand � Complete the data collecting and interpretation covered above. Ask students to record the electron

arrangement of each element and discuss the patterns (see Student Book p. 111). Make sure students understand

that it is the electron arrangement of atoms that determines the element’s physical and chemical properties.

Plenary suggestions In groups, students should prepare a summary of what they have discovered about the relationship between the

Periodic Table and the properties of elements, for example:

� metals on the left, non-metals on the right

� conductors on the left, insulators on the right

� solids on the left, gases on the top right

� melting point increases then falls left to right

� density increases and then falls left to right.

Student Book answers Q1 a) Any appropriate answers (include semi-metals as metals)

b) Any appropriate answers (include semi-metals as metals)

Q2 The predicted properties of elements have been shown to be correct.

Q3 They conduct electricity and heat, are malleable and shiny.

Q4 Melting points increase to a peak at silicon and then fall rapidly. (The last two should be below the metals.

Ignore the fact that phosphorus has a slightly lower melting point than sulfur.)

Q5 Metal. It only has one electron in its outer shell.

Q6 20, 2.8.8.2, metal


Q2 a) F, non-metal

b) Rb, metal

c) Kr, non-metal

d) Co, metal

e) Sr, metal

f) Te, non-metal


Q2 a) Bar chart – check scales and accuracy (the melting point has a very wide range).

b) Melting points rise to a maximum at carbon, then fall very abruptly and continue to fall.

c) Yes, but there is not such a sharp fall after the peak.

Q3 a) Density rises from Na to Al.

b) Density falls from P to Ar.

c) Si is less dense than Al and P. It is a semi-metal.


Q1 If number of electrons in the outer shell < 3, then it is a metal. If number > 4, it is a non-metal.

Q2 Yes. Increasing number of electrons in outer shell causes a rise in m.p. up to 3 or 4 electrons, then it falls.

Q3 a) Bottom, far right, because seven shells = seventh period, and eight outer electrons = eighth column (group)

b) Non-metal, because more than four electrons in outer shell.

c) Gas. Trend in period shows that elements on far right are gases. All the elements with eight electrons in

their outer shell are gases. Noble gases have low melting and boiling points.




Technician sheet


Demonstration: elements in the Periodic Table

Samples of as many elements as possible are required, but in particular:

� sodium (kept under oil – cut a piece to show its silver colour)

� magnesium (ribbon)

� aluminium (foil or sheet)

� silicon (lump or powder)

� phosphorus (red)

� sulfur (powder or lump)

� chlorine (in fume cupboard – can be prepared by reaction of potassium manganate(VII) with hydrochloric acid – see CLEAPSS recipe and hazard notes)

� argon (an old clear incandescent light bulb can be used to show it is a colourless gas)

[Note: other hazardous elements such as lithium, potassium, barium and bromine should be kept in a fume cupboard.]

� white tile

� sharp knife

� tongs or tweezers

� battery or power pack

� 3 leads, light bulb in holder, crocodile clips

Method

Set up an electrical circuit with the ends of two of the leads touching a sample of one of the elements. Set the power pack to 6 V and turn the power on. Demonstrate that the metals conduct electricity by the lighting of the bulb.

This can be done with sodium, as well as magnesium, aluminium and silicon, if a new surface is cut and the oil wiped off the sodium sample.

Phosphorus is a soft solid. It can be tested as described above, but make sure that no phosphorus is left on the end of the leads.

Demonstrate the malleability of sodium, magnesium and aluminium. The brittleness of sulfur can be demonstrated with a thin sliver shipped off a lump.

Notes

� The aim is for students to note the appearance of each element, the state of the element and, if solid, whether it conducts electricity.

� A lump of silicon looks metallic, but it will not conduct electricity sufficiently to light the bulb.

Health and Safety

� Sodium is FLAMMABLE and REACTS VIOLENTLY with water. Students should not touch it.

� Phosphorus is FLAMMABLE and should not be touched.

� Chlorine is TOXIC and CORROSIVE. If a member of CLEAPSS, follow their guidelines for chlorine production in Recipe 24. Chlorine should be kept in a fume cupboard.

� There is no need for students to touch any elements.

� If a member of CLEAPSS, refer to Hazcards 88, 73 and 22A.




1 Locating elements in the Periodic Table

1 Investigate the properties of elements in the Periodic Table. Fill out a table or cards with details about each element or design a spreadsheet or database. An example of an element card is shown below.

2 Find the elements listed below on the Periodic Table. The proton number is given after the name. From their position, state whether you think each is a metal or a non-metal.

a) fluorine (9) …………………… b) rubidium (37) ……………………

c) krypton (36) …………………… d) cobalt (27) ……………………

e) strontium (38) …………………… f) tellurium (52) ……………………

Name of element

Symbol Proton number

Electron arrangement

State at room temperature

Appearance

Conductor / insulator

Metal / non-metal



2 Patterns in the Periodic Table

1 Investigate the elements in the third period of the Periodic Table. Complete a set of data cards for these elements, using information from the table below. You could use the format in the sample data card for silicon, shown below the table.

Name sodium magnesium aluminium silicon phosphorus sulfur chlorine argon

Symbol Na Mg Al Si P S Cl Ar

Proton

number

11 12 13 14 15 16 17 18

Relative

Atomic

Mass

23 24 27 28 31 32 35.5 40

Electron

arrangement

2.8.1 2.8.2 2.8.3 2.8.4 2.8.5 2.8.6 2.8.7 2.8.8

Melting

point (°°°°C)

98 649 660 1410 44 113 –101 –89

Boiling

point (°°°°C)

883 1107 2467 2355 280 445 –35 –86

2 The table below shows the melting points of the elements in the second period of the Periodic Table.

Element lithium beryllium boron carbon nitrogen oxygen fluorine neon

Proton number

3 4 5 6 7 8 9 10

Melting point (°C)

181 1278 2300 3652 –210 –218 –220 –248

a) Draw a suitable chart to display this information.

b) Describe the patterns that you can see in your chart.

c) Do the melting points of elements in the third period follow a similar pattern?

Silicon, Si Proton number 14 RAM 28 Electron arr. 2.8.4

Melting point 1410 °C Boiling point 2355 Conductivity semi-conductor Appearance Grey crystalline

solid. Brittle. Unreactive.



3 The chart below shows the pattern of the density of the elements in the third period. Density is the mass of 1 cm3 of the solid element. For elements that are gases at room temperature, the density is measured at a temperature below their melting point.

a) What is the pattern for the density for the metals in the third period?

b) What is the pattern for the density of the non-metals in the third period?

c) Are there any elements that do not fit the patterns? Suggest a reason for this.

3 Predicting properties

Look at your data on elements, in particular those in the third period of the Periodic Table.

1 How is the electron arrangement related to whether the element is a metal or a non-metal?

2 Is the electron arrangement of an element related to its melting point? Explain your answer.

3 Scientists are hoping to produce enough atoms of the element with the proton number 118 to be able to test its properties. It will have seven shells of electrons, with eight electrons in its outer shell.

a) Where will it be placed in the Periodic Table? Explain your answer.

b) Is it expected to be a metal or a non-metal? Explain your answer.

c) Is it likely to be a solid or a gas at room temperature? Explain your answer.



c4_06 Group 1 – the alkali metals

Resources

Student Book pages 112−113 � Interactive Book: Matching pairs ‘Alkali metals’ � Homework pack c4_06


Equipment for demonstration; large Periodic Table for display

Learning outcomes C4.1.7 understand that a group of elements is a vertical column in the Periodic Table and that the elements in a

group have similar properties

C4.1.11 recall and recognise the chemical symbols for the Group 1 metals (also known as the alkali metals)

lithium, sodium and potassium

C4.1.12 recall that the alkali metals are shiny when freshly cut but tarnish rapidly in moist air due to reaction with

oxygen

C4.1.13 use qualitative and quantitative data to identify patterns and make predictions about the properties of

Group 1 metals (for example, melting point, boiling point, density, formulae of compounds and relative reactivity)

C4.2.13 understand that elements in the same group have the same number of electrons in their outer shell and

how this relates to group number

C4.2.14 understand that the chemical properties of an element are determined by its electron arrangement,

illustrated by the electron configurations of the atoms of elements in Groups 1 and 7

Numeracy focus: Drawing and interpreting graphs and charts from secondary data; extracting information from

the Periodic Table; extracting information from charts and graphs, including patterns in the properties of elements.

ICT focus: Accessing data from the internet; preparing data on elements in spreadsheets or databases.


� explain that groups of elements in the Periodic Table have similar properties

� remember the symbols and properties of the Group 1 metals

� use data and patterns in groups to make predictions about the properties of elements.

Key vocabulary

group �� alkali metal �� tarnish


Some students may have difficulty in recognising that while each element is different, a group has similarities in

behaviour that show a trend within the group. Some may have difficulty recalling the symbols Na for sodium and K

for potassium.


� Display a Periodic Table. Ask students to name the elements in groups 0,1, 2 and 7. Ask if anyone can say what

is the same about the electron shells of the elements that make up a group.

Learning activities worksheet c4_06 Low demand � Tell students that this lesson is about the first group, Group 1, on the left of the Periodic Table. Go

through the symbols and names of each element. Ask students to state whether they are metals or non-metals

(see previous lesson). Emphasise that Mendeleev put these elements in the same group because they have similar

properties. Demonstrate the appearance and properties of the alkali metals lithium, sodium and potassium: colour,

shininess when freshly cut, hardness, speed of tarnishing, and (optionally) the reaction of sodium with oxygen. See

the Technician sheet. Draw student’s attention to the variation in hardness and rate of tarnishing down the group.

(Note that the reaction with air is more complex than discussed here with peroxides and superoxides formed by

some elements. If the air is moist, then hydroxides and carbonates will be formed, and tarnishing will occur faster.)

Students should record their observations of the demonstrations for inclusion in their database of elements.

Teaching and learning notes: Students need to recall the symbols of the first three members of Group 1.

Standard demand � Explain in more detail why the alkali metals are put in the same group (similar properties, one

electron in the outer shell), and write word equations for the reactions with oxygen and the formulae of the oxides.

All the metals form an oxide with the formula M2O.

c4_06 Group 1 – the alkali metals continued


Discuss patterns (trends) in the physical properties of the members of the group (see Student Book p. 113 for a

graph of melting points). The worksheet has further tasks on other properties.

Students may add further information to their database on elements. They should note similarities and differences

within the group and produce a summary of the trends of properties in Group 1. Individually or in small groups they

could be asked to discuss the expected appearance and properties of rubidium, caesium and/or francium.

Teaching and learning notes: Students will need access to data on the Group 1 elements.

High demand � Ask Higher-tier students to write or sketch the electron arrangements of Li, Na and K. They should

compare the data to the structures and suggest correlations: one outer shell electron explains the similarities in

behaviour, while the increasing size/number of shells explains the trends down the group (do not go into the

formation of ions, or metallic bonding at this point). The worksheet gives further tasks on this.

Plenary suggestions Ask students to give quick definitions or descriptions of the following terms:

� group – column in the Periodic Table / elements with same number of outer shell electrons

� shell – area around the nucleus of an atom occupied by electrons

� metal – element that is shiny, malleable and conducts electricity (and heat)

� alkali metal – Group1 / metals that are soft, reactive, tarnish rapidly / burn in oxygen / oxides formula M2O.

Student Book answers Q1 a) lithium, Li; sodium, Na; potassium, K

b) They have similar properties. They are all reactive metals. They have one electron in their outer shells.

Q2 Lithium + oxygen → lithium oxide

Q3 More quickly. Rb is below K in the Periodic Table and the trend is that they are more reactive down the group.

Q4 The melting point decreases down the group. It is a curve with a decreasing negative gradient.

Q5 The sixth shell

Q6 All the alkali metals react rapidly with oxygen to form the oxide because they have one electron in their outer

shell. (The reaction becomes more violent down the group because: the atoms get bigger / the outer electron

is further away / the atoms have more shells further down the group.)


Q1 lithium Li, sodium Na, potassium K

Q2 … groups … elements … similar … left … potassium … metals … alkali …


Q2 a) lithium, sodium, potassium (density less than 1 g/cm3)

b) The conclusion is generally true. There is a positive correlation between proton number and density.

Potassium is an anomaly. The increase in density decreases as proton number increases: the graph levels off.

c) An answer between 2.0 and 2.5 g/cm3 is acceptable.

Q3 b) There is a negative correlation between proton number and boiling point. Boiling point decreases as proton

number increases, but the change gets smaller at higher proton numbers: the graph levels off.

c) An answer between 650 and 669 °C is acceptable.

Q4 Na2O, K2O


Q1 a) True

b) False. There is a negative correlation between the boiling point of the alkali metals and the number of shells

occupied by electrons in the atoms.

c) False. The formula of the Group 1 metal oxides has the pattern M2O because all the metal atoms have one

electron in their outer shell.

d) True.

Q2 Rubidium atoms have an extra shell of electrons.

Q3 Element E. It has the same outer electron arrangement as Group 1 / alkali metals.




Technician sheet


Demonstration – properties of lithium, sodium and potassium

� samples of lithium, sodium and potassium (stored in oil)

� white tile

� sharp knife or scalpel


� paper towel

Reaction with oxygen (optional)

� supply of oxygen gas (if a member of CLEAPSS, see Recipe 64 if a cylinder of oxygen is not available – oxygen may be collected over water)

� gas jar and lid

� combustion spoon

� Bunsen burner

� red litmus paper or pH paper

� goggles and protective screen

Method

Demonstration – properties of lithium, sodium and potassium

Cut a piece of each metal and wipe off the oil. Show students the shiny silver surface and wait while the surface tarnishes and turns dull.

Reaction with oxygen (optional)

It is suggested that this is just done with sodium.

Collect a gas jar of oxygen. Cut a pea-sized piece (less than 1 g) of sodium and wipe off the oil. Place the sodium in a combustion spoon. Hold the spoon in a blue Bunsen flame until it has melted and is glowing yellow. Quickly transfer the spoon to the gas jar of oxygen. Make sure the spoon is held in the oxygen and does not fall to the bottom of the gas jar, where there may be some water. The sodium should burn with a bright yellow flame, producing a white smoke of sodium oxide. When combustion is complete, the smoke may be tested with an indicator to show it is alkaline.

Health and Safety notes

� The reaction of sodium with oxygen can be violent. Make sure protective screens are in place. As lithium has a higher melting point, it can burn explosively; and potassium burns rather violently. It is therefore recommended that these metals are NOT used.

� Wear goggles.

� Lithium, sodium and potassium are FLAMMABLE metals with a VIOLENT REACTION with water. Keep the metals away from water or moisture. Avoid contact with skin.

� Students should not handle the metals.

� If a member of CLEAPSS, read the Model Risk Assessment on Hazcard 88.




1 Symbols and groups

1 Match up the names and symbols of these elements in Group 1:

Lithium Na

Sodium K

Potassium Li

2 Fill in the gaps in the following sentences with words from the list below.

metals elements left groups alkali potassium similar

The vertical columns in the Periodic Table are called ………………… The

………………… in a group have ………………… chemical properties. Group 1 is the

group on the ………………… of the Periodic Table, made up of the elements lithium,

sodium, …………………, rubidium, caesium and francium. All the elements in Group 1

are reactive ………………… The Group 1 elements are also known as the

………………… metals.

2 Properties and patterns

1 Record the appearance and properties of the alkali metals in a table, on data cards or on a spreadsheet or database. Include the melting point, boiling point, density and the formula of the oxide for each metal.

2 The graph shows how the density of the alkali metals changes with proton number. Look at the graph and answer the questions on the next sheet.

c4_06 Group 1 – the alkali metals continued


a) The density of water is 1 g/cm3. Which alkali metals will float on water?

b) A conclusion from these data could be that density increases with proton number. Evaluate how well the data support this conclusion.

c) Suggest a density for francium, proton number 87.

3 The table below shows data on the boiling points of the alkali metals.

Element lithium sodium potassium rubidium caesium

Proton number

3 11 19 37 55

Boiling point (°C)

1342 883 760 686 669

a) Plot a suitable chart to display the data.

b) Describe the pattern in the boiling points of Group 1.

c) Suggest a boiling point for francium, proton number 87.

4 The formula of lithium oxide is Li2O. What are the formulae of sodium oxide and potassium oxide?

3 Comparing electron arrangements

To answer these questions, you may need to use data given in the questions in Activity 2. In question 1c) below, M stands for the symbol of any Group 1 element.

1 Which of the following statements are true? Write corrections for the statements that are false.

a) There is a positive correlation between the density of the alkali metals and the number of shells occupied by electrons in the atoms.

b) There is a positive correlation between the boiling point of the alkali metals and the number of shells occupied by electrons in the atoms.

c) The formula of the Group 1 metal oxides has the pattern M2O because all the metal atoms have one electron in the first shell.

d) Group 1 alkali metals tarnish quickly in air because they have the same electron arrangement in their outer shells.

2 Why are rubidium atoms bigger than potassium atoms?

3 Which of the elements in the following table is likely to be a soft metal that tarnishes easily? Explain your answer.

Element Proton number Number of electrons in outer shell

A 83 5

B 84 6

C 85 7

D 86 8

E 87 1



c4_07 Group 1 – reactions with water

Resources

Student Book pages 114−115 � Interactive Book: Practical investigations ‘Acids, alkalis and bases’ � Homework pack c4_07


Materials for demonstration; copies of hazard symbols

Learning outcomes C4.1.14 describe the reactions of lithium, sodium and potassium with cold water

C4.1.15 recall that alkali metals react with water to form hydrogen and an alkaline solution of a hydroxide with the

formula MOH

C4.1.17 understand and give examples to show that the alkali metals become more reactive as the group is

descended

C4.1.18 recall the main hazard symbols and be able to give the safety precautions for handling hazardous

chemicals (limited to explosive, harmful, toxic, corrosive, oxidising, and highly flammable)

C4.1.19 state and explain the precautions necessary when working with Group 1 metals and alkalis

C4.1.28 (part) recall the formulae of hydrogen and water molecules

C4.1.29 write word equations for reactions of alkali metals and halogens in this module and for other reactions

when given appropriate information

C4.1.30 interpret symbol equations, including the number of atoms of each element, the number of molecules of

each element or covalent compound and the number of ‘formulae’ of ionic compounds, in reactants and products

(in this context, ‘formula’ is used in the case of ionic compounds as an equivalent to molecules in covalent

compounds; the concept of the mole is not covered in the specification)

C4.1.31 balance unbalanced symbol equations

C4.1.32 (part) write balanced equations, including the state symbols (s), (g), (l) and (aq), for reactions of

alkali metals in this module and for other reactions when given appropriate information

C4.1.33 recall the state symbols (s), (l), (g) and (aq) and understand their use in equations

Literacy focus: Describing observations.

Numeracy focus: Balancing chemical equations.


� describe what happens when Group 1 metals react with water

� understand the hazards posed by these reactions

� interpret symbol equations for the reactions of alkali metals.

Key vocabulary

balanced symbol equation �� state symbols


Students may not be able to maintain concentration if they are standing some distance from the demonstration.


� Ask students to make a list of the things they should be looking for in reaction if water with alkali metals. e.g.

evolution of gas, tests to identify gases, colour changes, heat given out or taken in, identity of the products.

Learning activities worksheet c4_07 Low demand � Demonstrate the reaction of sodium with water (see the Technician sheet). Make sure students

observe and record all the points. Ask them to identify the products of the reaction: flammable gas – hydrogen

(note that careful observation should show that the gas burns but not the metal); alkaline solution – sodium

hydroxide. Write the word equation for the reaction.

Ask students whether they think the reaction is hazardous (yes!) and how the risk can be reduced: keep the

Group 1 metals away from water, only keep small pieces, ensure proper labelling and warning signs. Discuss

which warning signs are needed: ‘FLAMMABLE, EXPLOSIVE’ on the sodium bottle, ‘CORROSIVE’ on the

container holding the solution that is formed.

Ask students what they might expect to happen with lithium and potassium, and write word equations for the

reactions.

c4_07 Group 1 – reactions with water continued


Teaching and learning notes: Students may need help in seeing and recording all the relevant points. They may

need copies of hazard symbols for reference.

Standard demand � Students should also observe and record the reactions of lithium and potassium with cold

water, (see the Technician sheet). Discuss precautions taken to keep everyone safe. Note the similarities in the

reactions and the pattern – reactivity increases down the group. Explain that the similarities are due to the single

electron in the outer shell, and that the differences are due to the number of shells in the atoms.

Ask students to predict observations for the reaction of rubidium and/or caesium. (The metals are denser than

water and sink. The hydrogen is formed very quickly and explodes.) Video clips of these reactions can be shown.

Show students the balanced chemical equation with state symbols for the reaction of sodium with water. Ask them

to explain the meaning of the balancing numbers. The worksheet has further tasks to reinforce this work.

Teaching and learning notes: The demonstrations in this lesson help to develop observing skills. Recording

observations develops literacy skills.

High demand � Explain trends in terms of the electron arrangement – the increasing size of the atoms makes it

easier for the outer shell electrons to be involved in the reaction (do not go into the formation of ions at this point).

Show how balanced chemical equations are written for these reactions and for the reactions of the metals with

oxygen. The worksheet gives more tasks.

Plenary suggestions Ask students to list the main points of the reaction of the alkali metals with water: vigorous, hydrogen produced,

alkali (soluble metal hydroxide) formed, reactions more violent down the group.

Student Book answers Q1 Sodium hydroxide solution turns pH indicator blue. Q2 Lithium hydroxide solution and hydrogen gas

Q3 potassium + water → hydrogen + potassium hydroxide Q4 LiOH, KOH

Q5 Violent fizzing with immediate flames and explosions. Solution remaining would turn pH indicator blue/violet.

Products are hydrogen gas and rubidium hydroxide.

Q6 a) 2Li(s) + 2H2O(l) → H2(g) + 2LiOH(aq) b) 4Na(s) + O2(g) → 2Na2O(s)

c) 2Rb(s) + 2H2O(l) → H2(g) + 2RbOH(aq)


Q2 When sodium reacts with water, hydrogen is given off because water contains hydrogen atoms.

Lithium hydroxide would be formed if lithium is put in water.

Sodium is called a reactive metal because it reacts with cold water.

Lithium would be less reactive than sodium but potassium would be more reactive than sodium.

Q3 Highly flammable (because hydrogen is given off and sodium metal burns); explosive (because the reaction

can be explosive); corrosive (because a strong alkali is formed).


Q1 a) A flammable gas is formed. b) The indicator colour shows that an alkaline solution formed.

c) Each metal has one electron in the outer shell of its atoms.

d) Answers will vary, but should be based on: speed of production of gas; and the ease with which the gas

catches light or an explosion occurs.

Q2 a) Two molecules of water are involved in the reaction. b) Two

c) The numbers of atoms of lithium (2), hydrogen (4) and oxygen (2) are the same on the two sides of the

equation.

d) aq means aqueous solution / dissolved in water; g means gas. e) Yes – how fast or violent the reaction is.


Q1 All the atoms have the same outer shell electron arrangement, which determines how they react.

Q2 a) 4K(s) + O2(g) → 2K2O(s) b) 2K(s) + 2H2O(l) → 2KOH(aq) + H2(g)

Q3 a) 4Rb(s) + O2(g) → 2Rb2O(s) b) 2Cs(s) + 2H2O(l) → 2CsOH(aq) + H2(g)

Q4 2Na(s) + O2(g) → Na2O2(s)




Technician sheet


Demonstration – reacting lithium, sodium and potassium with water

� samples of lithium, sodium, potassium (stored in oil)

� large bowl or glass trough

� white tile

� knife or scalpel


� paper towel

� Bunsen burner and splints or taper

� indicator paper or solution (red litmus, phenolphthalein, methyl orange or full range)

� goggles/eye protection and safety screens

Method

Half fill the trough, bowl or beaker with cold water. Cut a pea-sized piece of sodium and wipe off the oil. Using the tongs or tweezers drop the sodium into the water. Light a splint or taper and hold the flame near the piece of sodium (but do not constrain the sodium). When all the sodium has reacted, test the water with the indicator paper or add indicator solution to the water.

Repeat the experiment with lithium and potassium.

Notes

Expected results

� Sodium: Moves around on the surface of the water, fizzing. It turns into a ball (the heat of reaction melts the sodium). Sodium burns with a yellow flame in a ring around the metal. Sometimes the sodium will explode, producing splashes of solid sodium oxide/hydroxide, particularly if it sticks to the side of the trough. The solution formed is alkaline.

� Lithium: Moves around the surface, fizzing more leisurely than sodium. Does not melt and retains its irregular shape. The gas can be ignited and burns with a red flame, but usually does not continue to burn. The solution is alkaline.

� Potassium: Always explodes on contact with water, scattering mauve sparks (the heat of reaction ignites the hydrogen). The solution is alkaline.

� The similarities and the relative reactivity of the metals should be readily apparent.

� Video clips of these metals and of rubidium and caesium reacting with water are readily available on the internet.

Health and Safety

� If a member, refer to Hazcard 88.

� The alkali metals are FLAMMABLE and REACT VIOLENTLY with water to produce an explosive gas (hydrogen) and a CORROSIVE liquid (the metal hydroxide solution). The reaction of the metals with water can be EXPLOSIVE. Wear goggles.

� Ensure that the place where the metals are prepared is dry. Erect protective screens between the trough and the students. Students should stand at least 2 m away from the screens and wear eye protection.

continued



� It is tempting to increase the size of the piece of metal used, but this increases the probability of explosion. Pieces of burning metal can fly several metres through the air. Small, similar sized pieces should be used to allow a comparison of the reactions.

� If a large bowl is used, metal samples should be reduced to the size of a rice grain.

� Ensure good ventilation.




1 Sodium and water

1 Write down in the table below what you see when sodium reacts with water.

Observations

What does the sodium look like at the beginning?

What does the sodium look like just after it is put in the water?

How do you know a gas is given off?

What happens when a lighted splint is held near the sodium?

Where did the sodium go?

What happened when indicator was put in the water?

2 Match up and join the beginnings and ends of the sentences shown below using one of the following connecting words:

but if because

When sodium reacts with water, hydrogen is given off …

… it reacts with cold water.

Lithium hydroxide would be formed …

… water contains hydrogen atoms.

Sodium is called a reactive metal …

… lithium is put in water.

Lithium would be less reactive than sodium …

… potassium would be more reactive than sodium.



3 Which of the hazard symbols below should be displayed around a place where sodium is reacted with water? Explain your answer.

2 Patterns in Group 1 reactions

1 Record your observations of the reaction of lithium, sodium and potassium with cold water. Then answer the following questions.

a) What evidence is there that hydrogen is formed in each reaction?

b) What evidence is there that the metal hydroxides are formed in each reaction?

c) Why are the reactions similar?

d) What evidence is there that the reaction gets more violent from lithium to potassium?

2 Look at the equation below for the reaction of lithium with water.

2Li(s) + 2H2O(l) → 2LiOH(aq) + H2(g)

a) What does the 2 in front of the H2O mean?

b) How many atoms of lithium are involved in the reaction?

c) How can you tell that the equation is balanced?

d) What do the letters (aq) and (g) mean after the formulae of the products?

e) Is there anything that the equation does not tell you about the reaction? If so, what is it?

3 Writing balanced equations

1 The balanced equation for the reaction of the Group 1 metals with water can be summarised as follows:

2M(s) + 2H2O(l) → 2MOH(aq) + H2(g)

where M stands for the symbol of the Group 1 metal.

Explain with reference to the electron arrangement why the same equation applies to the reactions of all the Group 1 metals with water.

2 Balance the equations shown below.

a) K(s) + O2(g) → K2O(s) b) K(s) + H2O(l) → KOH(aq) + H2(g)

3 Write balanced chemical equations for the following word equations.

a) rubidium + oxygen → rubidium oxide

b) caesium + water → caesium hydroxide + hydrogen

4 Write a balanced chemical equation for the reaction described below:

When solid sodium is burned in pure oxygen gas, it produces a white solid called sodium peroxide, which has the formula Na2O2.



c4_08 Group 1 – reactions with chlorine

Resources

Student Book pages 116−117 � Interactive Book: Quick starter ‘What salts can be made’ � Homework pack c4_08


Materials for the demonstration; copies of hazard symbols

Learning outcomes C4.1.16 recall that alkali metals react vigorously with chlorine to form colourless, crystalline salts with the formula

MCl

C4.2.14 understand that the chemical properties of an element are determined by its electron arrangement,

illustrated by the electron configurations of the atoms of elements in Groups 1 and 7

Literacy focus: Describing observations.


ICT focus: Accessing video clips of reactions from the internet.


� describe and explain the reactions of Group 1 metals with chlorine.


It is not feasible to show the reactions of all the alkali metals with chlorine, so the similarities and trend in reactivity

have to be based on secondary evidence, i.e. the similarities of the products.


� Tell students that they are going to observe the reaction of sodium with chlorine, a corrosive, toxic gas. Ask them

to discuss the hazards of the demonstration and suggest safety precautions – keep sodium away from water,

handle chlorine in fume cupboard, wear goggles and gloves (see Technician sheet c4_08).

Learning activities worksheet c4_08 Low demand � Demonstrate the reaction of sodium with chlorine gas (see the Technician sheet). If a member,

refer to CLEAPSS Guide L195 page 32. Video clips of the reaction are available on the internet (see CLEAPSS

YouTube account). Ask students to describe what they see and record their observations. They should note the

differences between the product (a white powder) and the reactants. Ask them to explain what has happened. If

students do not offer an explanation themselves, then explain that sodium and chlorine have formed a compound,

sodium chloride, that has different properties from the elements. Ask students to write a word equation for the

reaction, and for the similar reactions of lithium and potassium with chlorine. The worksheet gives further tasks on

this topic, summing up the reactions of the alkali metals.

Teaching and learning notes: Students need to appreciate that a compound has different properties from the

elements from which it is composed, and that reactive elements form stable compounds.

Standard demand � Point out that in the reaction demonstrated, sodium has combined with chlorine to form a

compound with the formula NaCl. Ask students to predict the formulae and properties of other Group 1 chlorides

and to predict the relative vigour of the reactions. Display the equations for the reactions and make sure students

understand how to interpret them. The worksheet gives further tasks, summing up the patterns in Group 1.

Teaching and learning notes: Students should recall the general formula of alkali metal chlorides, but no

discussion of oxidation states/valencies is required at this stage.

High demand � Show samples of lithium chloride, sodium chloride and potassium chloride. Given the symbols and

formulae, ask Higher-tier students to write a balanced chemical equation for sodium reacting with chlorine. Ask

them to write similar equations for other Group 1 elements and to relate the similarities and differences in the

reaction to the electron arrangements of Group 1 elements. The worksheet gives further tasks, summing up the

explanations for the properties of Group 1.

Plenary suggestions Ask students to summarise and revise the properties of the alkali metals, their similarities and trends in the group

(with explanations based on electron arrangements for Higher-tier students). Groups of students could prepare

c4_08 Group 1 – reactions with chlorine continued


questions (and answers) to be posed to other groups. As an alternative, students can produce a mind map (Activity

4 on the worksheet).

Student Book answers Q1 There is a flame. The sodium glows as it reacts. Q2 Potassium chloride Q3 A white solid

Q4 LiCl, KCl Q5 Violent reaction, flame, white smoke Q6 Two

Q7 Similarities – reactions with oxygen, water and chlorine, and the products and their formulae (M2O, MOH, MCl)

Differences – reactions get more violent down the group / metals are more reactive down the group

Q8 a) 2K(s) + Cl2(g) → 2KCl(s) b) 2Rb(s) + Cl2(g) → 2RbCl(s)

Q9 The reactions are similar because each metal has one electron in its outer shell. The reactions get more

violent down the group, because the outer electron is further from the nucleus / in higher shell or level / atoms

are bigger.


Q1 a) Sodium is a silvery metal; and chlorine is a green gas.

b) There is a reaction, with a yellow flame/glow and (off-) white smoke.

c) The product is a white solid.

Q2 a) A shiny metal: sodium

A green gas: chlorine

White crystals: sodium chloride

Reacts with water and oxygen: sodium

Dissolves in water without a reaction: sodium chloride

b) Compounds have different properties from the elements that make them up and often very reactive

elements form safe, stable compounds.

Q3 lithium + chlorine → lithium chloride; potassium + chlorine → potassium chloride


Q1 Sodium is a silvery metal, and chlorine is a green gas. There is a reaction, with a yellow flame/glow and (off-)

white smoke. The product is a white solid.

Q2 a) Reactants: lithium, chlorine. Product: lithium chloride.

b) The symbol (s) after the formula stands for solid.

c) Two formula units of lithium chloride, which are formed for every two atoms of lithium and one molecule of

chlorine that react.

Q3 a) Fizzes in water.

b) All the alkali metal chlorides have similar properties and appearance because the metals are in the same

group in the Periodic Table, so have the same outer shell electron arrangement.

Q4 They both burn in oxygen to form a solid oxide. They both react with water to form hydrogen gas and an

alkaline hydroxide. They both react with chlorine to form a white crystalline salt.


Q1 Sodium is a silvery metal, and chlorine is a green gas. There is a reaction, with a yellow flame/glow and (off-)

white smoke. The product sodium chloride is a white solid, just like lithium chloride and potassium chloride.

Q2 a) RbCl b) 2Rb(s) + Cl2(g) → 2RbCl(s) c) They all have one electron in their outer shells.

d) The reaction becomes faster / more violent down the group, because the atoms get larger / have more

electron shells / the outer electron is further from the nucleus.

Q3 Potassium chloride has similar properties to sodium chloride because sodium and potassium are in the same

group / have the same outer shell electron arrangement.

Q4 a) 4K(s) + O2(g) → 2K2O(s) b) 2Rb(s) + 2H2O(l) → 2RbOH(aq) + H2(g) c) 2Cs(s) + Cl2(g) → 2CsCl(s)

Activity 4 (Plenary activity)

The mind map should include the names and symbols of at least the first three members of the group, with their

electron arrangements, a description of the appearance of the metals and trends in properties, a description of the

reactions, the names and properties of the products, along with similarities and trends. Higher-tier students should

include balanced chemical equations and explain the trends in terms of the electron arrangements.




Technician sheet


Demonstration: reacting sodium with chlorine

� gas generator for chlorine (chlorine can be generated by reacting concentrated hydrochloric acid with potassium manganate(VII) – if a member of CLEAPSS, see Recipe 24)

� gas jar and lid


� Bunsen burner

� fire brick

� knife


� sodium

� samples of sodium chloride, potassium chloride and lithium chloride

� goggles and gloves

Method

Collect a gas jar of chlorine gas by upward displacement of air.

Cut a pea-sized piece of sodium and wipe off the oil with a paper towel. Place the sodium in a combustion spoon and hold it in the Bunsen flame until glowing. Plunge the combustion spoon into the chlorine gas.

Wait until the reaction is complete and take care washing out the gas jar. It is useful to have a bowl of cold water to put the combustion spoon in when it is cool. (Note that any sodium left that has not reacted with the chlorine will react violently with the water.)

Notes

� The reaction will produce a smoke. It will be off-white, not white, because residual oil left on the sodium burns to form soot.

� Lithium and potassium react in a similar way to sodium, but it is not recommended that their reaction is demonstrated. Lithium has to be heated to a higher temperature to melt it and can react EXPLOSIVELY, and the reaction of potassium can be more VIOLENT.

� Display samples of lithium chloride, sodium chloride and potassium chloride, and note their similar appearance.

� Video clips of the reaction of these metals with chlorine can be found on the internet.

Health and Safety

� Chlorine gas is TOXIC and CORROSIVE. It should only be made in small quantities in a fume cupboard. If a member, follow the guidelines in the CLEAPSS instructions.

� Sodium is FLAMMABLE and reacts VIOLENTLY with water. Make sure the work area is dry and ensure that sodium does not come into contact with water or skin. The reaction of sodium with chlorine is VIOLENT.

� The experiment should be carried out in a fume cupboard with adequate ventilation. Students should watch from a distance of at least 2 m and wear eye protection.

� If a member, refer to CLEAPSS Guide L195 page 32 and Hazcard 88.




1 Observations and patterns

1 a) Describe the appearance of sodium and chlorine.

................................................................................................................................

................................................................................................................................

b) Write down what you see when a piece of sodium metal is placed in chlorine gas.

................................................................................................................................

................................................................................................................................

................................................................................................................................

c) Describe what the product of the reaction looks like.

................................................................................................................................

2 a) Put ticks in the columns to show which properties apply to sodium, which apply to chlorine and which apply to sodium chloride.

Property Sodium Chlorine Sodium chloride

A shiny metal

A green gas

White crystals

Reacts with water and oxygen

Dissolves in water without a reaction

b) Someone says to you that because sodium chloride contains the dangerous elements sodium and chlorine, then sodium chloride must be dangerous too. How would you reply?

................................................................................................................................

................................................................................................................................

................................................................................................................................

3 What are the word equations for the following reactions?

a) lithium with chlorine

................................................................................................................................

b) potassium with chlorine

................................................................................................................................



2 Formulae and salts

1 Record your observations of the reaction of sodium and chlorine, including the appearance of the reactants and product.

2 a) Name the reactants and the product in the reaction:

2Li(s) + Cl2(g) → 2LiCl(s)

b) How can you tell that the product is a solid?

c) What does 2LiCl stand for?

3 a) An alkali metal reacts with chlorine to form a salt. Which of the following is not a property that you would expect the salt to have?

crystalline solid white in colour fizzes in water soluble in water

b) Explain why you can be sure that your answer can apply to the salt of any alkali metal.

4 Give three examples of reactions that show that sodium and caesium should be put in the same group in the Periodic Table.

3 Equations and trends (Higher tier only)

1 Record your observations for the reaction of sodium with chlorine, and compare the appearance of the product with lithium chloride and potassium chloride.

2 Rubidium (Rb) is an alkali metal that burns violently in chlorine gas (Cl2), producing a salt.

a) What is the formula of the salt formed?

b) Write a balanced chemical equation for the reaction of rubidium with chlorine, including state symbols.

c) Why does rubidium react in a similar way with chlorine as the other members of the alkali metals group?

d) How does the reaction of the alkali metals with chlorine change as you go down the group? Explain the change.

3 Eating too much sodium chloride is said to be bad for health. Why is it possible to replace the sodium chloride with potassium chloride in some brands of low-sodium salt?

4 Complete and balance the following equations for the reactions of the alkali metals.

a) ……K(s) + ……………(g) → ……K2O(s)

b) ……Rb(s) + ……H2O(l) → …………………(aq) + ……H2(g)

c) ……………(s) + ……………(g) → ……CsCl(……)



4 Plenary activity

Design a mind map for the reactions and properties of the alkali metals group. You can use the diagram below as a starting point.

ALKALI METALS

Names and symbols

Appearance and properties

Group 1

Electron arrangements

Reaction with oxygen

Reaction with chlorine

Reaction with water



c4_09 Group 7 – the halogens

Resources



Equipment and materials for demonstration

Learning outcomes C4. 1.7 understand that a group of elements is a vertical column in the Periodic Table and that the elements in a

group have similar properties

C4.1.20 recall and recognise the chemical symbols for the atoms of the Group 7 elements (also known as the

halogens) chlorine, bromine and iodine

C4.1.21 recall the states of these halogens at room temperature and pressure

C4.1.22 recall the colours of these halogens in their normal physical state at room temperature and as gases

C4.1.23 recall that the halogens consist of diatomic molecules

C4.1.24 use qualitative and quantitative data to identify patterns and make predictions about the properties of the

Group 7 elements (for example melting point, boiling point, formulae of compounds and relative reactivity)

C4.2.13 understand that elements in the same group have the same number of electrons in their outer shell and

how this relates to group number


arrangement, illustrated by the electron configurations of the atoms of elements in Group 7

Literacy focus: Describing the elements and making predictions about the properties of other halogens.

Numeracy focus: Plotting, drawing and interpreting graphs and charts from secondary data; extracting information

from charts and graphs, including patterns in the properties of elements.

ICT focus: Selecting and presenting data in a variety of forms to explore patterns and trends; using an interactive

Periodic Table to explore similarities and differences between elements; using a spreadsheet to display patterns in

chemical data.


� remember the formulae and properties of some elements in Group 7

� use patterns in their behaviour to predict properties of other members of Group 7.

Key vocabulary

halogen �� diatomic molecule


The difference in appearance in the halogens may mask the similarities in the behaviour of the elements.


� Tell students that iodine solution is often used to disinfect wounds, and chlorine is added to water to kill

microbes. Bromine too has been used as a disinfectant, for example in swimming pools. Ask students to discuss

why these elements should each have similar uses. Answers should include that they are in the same group in

the Periodic Table, they react in a similar way, and they have the same outer electron arrangement.

Learning activities worksheet c4_09 Low demand � Show samples of chlorine, bromine and iodine. (Health and Safety: care is needed, because all

the elements are hazardous – see the Technician sheet.) Alternatively, show pictures or video clips of the

elements. Ask students to describe them and, in particular, note their colour and state. Explain that, although the

elements appear very different, they have similarities that we will discover in later lessons and hence they are put

in the same group. Go through the names and symbols of the elements in the group. The worksheet gives further

tasks on this.

Teaching and learning notes: Students need to recall the names and symbols of the elements in the group.

Standard demand � Give the symbols of the three halogens. Point out that the elements appear different but have

similarities, e.g. they all exist as diatomic molecules. Explain what this means and ensure students understand the

formulae of the halogen molecules. Note the difference between a single atom and a molecule. Also explain that

trends in properties can reveal relationships between elements in a group. Display the graph of the melting points

c4_09 Group 7 – the halogens continued


of the elements (see Student Book p. 119) and discuss what it shows. Go over the electron arrangement of F and

Cl and outer shell arrangement of Br, I and At. The worksheet gives tasks on other properties of the halogens.

Show samples of alkali metal halides (see the Technician sheet). Point out that the halogens all form salts with the

alkali metals with formula MX (M = alkali metal, X = halogen).

Teaching and learning notes: Students should recall the formulae of the alkali metal chlorides from the previous

lesson.

High demand � Relate the trends in the properties of the halogens to the increasing size of the atoms / number of

shells. Relate the formulae of the salts to the seven outer electrons. The worksheet provides further tasks on this.

Plenary suggestions Ask students to predict the properties of astatine (or element number 117): crystalline solid, melting point and

boiling point, diatomic molecule, and formula of sodium salt.

Student Book answers Q1 a) iodine b) chlorine c) bromine d) astatine

Q2 It is a molecule of chlorine made up of two chlorine atoms.

Q3 Iodine, astatine

Q4 They increase steadily, but the size of the increases would get smaller further down the group.

Q5 2.8.7

Q6 Five

Q7 The properties of the elements in a group follow a pattern/trend and it is possible to extrapolate/extend the

properties to astatine.


Q1 chlorine: gas; green

bromine: liquid; red-brown

iodine: solid; grey

Q2

1C H L O R I N E

2S A L T

3C L

4B R O M I N E

5G A S

6S E V E N

7I O D I N E


Q1 See answer to Q1 in Activity 1 above.

Q2 a) diagram A b) diagram B c) It has two atoms joined together.

d) Diagram is the same as B but with Br in the circles. e) At2

Q3 a) A suitable and correct line graph or bar chart.

b) The boiling points increase (relatively) steadily. There is a (relatively) smooth line with a positive gradient.

c) Fluorine and chlorine d) 340 °C

Q4 a) LiF b) KBr c) NaI


Q1 a) Sketch showing electrons in three shells with two, eight and seven electrons (2.8.7).

b) It should have seven electrons in its outer shell.

Q2 The atoms get bigger / they have more shells / the outer electrons are further from the nucleus.

Q3 They all have seven electrons in their outer shell.



c4_09 Group 7 – The halogens

Technician sheet


Demonstration – appearance and physical properties of halogens

Samples of the following are required:

� chlorine gas (see the CLEAPSS Recipe 24 for preparing chlorine from the reaction of concentrated hydrochloric acid and potassium manganate(VII))

� bromine

� iodine

� any alkali metal halides that are available, e.g. lithium, sodium and potassium chlorides, bromides and iodides

In addition:

� 2 gas jars and lids


� Bunsen burner

� indicator paper

Method

This demonstration must be carried out in a fume cupboard.

Fill a gas jar with chlorine and cover with a lid. Put a moist piece of indicator paper in the gas jar of chlorine.

Put a drop of bromine in a gas jar and cover with a lid (if a member, refer to the Model Risk Assessment on Hazcard 15A).

Place a few crystals of iodine on a combustion spoon. Hold the spoon briefly in a Bunsen flame.

Notes

Students should observe that:

� chlorine is a green gas

� bromine is a red-brown liquid that evaporates easily, forming a red-brown vapour

� iodine is a grey solid (larger pieces may show that it is crystalline), which turns into a purple gas on warming

� all the alkali metal halides are white crystalline solids.

Health and Safety

� Chlorine, bromine and iodine are CORROSIVE and TOXIC. They should only be handled by trained staff, in a fume cupboard. Students should stay far enough from the fume cupboard that they do not inhale any fumes.



c4_09 Group 7 – the halogens

1 Symbols and appearance

1 Look at samples or pictures of the Group 7 elements chorine, bromine and iodine. Write down what they look like in the table below.

Name of element State at room temperature Appearance (colour)

2 Find the key word (shaded) by filling in the answers to the clues below.

1 The second element in this group of elements

2 The product formed when chlorine reacts with sodium

3 The symbol of chlorine

4 This element is a red-brown liquid

5 The state that chlorine is in at room temperature

6 The group of elements that has fluorine at the top

7 This element is a grey solid at room temperature

? [Keyword clue] Another name for the group of elements

?

1

2

3

4

5

6

7



2 Molecules and properties

1 Make observations of the state and appearance of the halogens at room temperature.

2

a) Which diagram represents an atom of chlorine?

b) Which diagram represents the form that chlorine takes in a gas at room temperature?

c) Why is diagram B called a diatomic molecule?

d) Sketch a diagram to show the arrangement of bromine atoms in liquid bromine.

e) What would be the formula of an astatine (At) molecule?

3 The table below shows the boiling points of the halogens.

Element Proton number Boiling point (°C)

fluorine 9 –188

chlorine 17 –35

bromine 35 59

iodine 53 184

astatine 85

a) Plot the data onto a suitable chart.

b) Describe the pattern of the boiling points in Group 7.

c) Which members of Group 7 are gases at room temperature?

d) The boiling point of astatine has not been measured accurately because astatine is so rare. Which of the following may be the boiling point of astatine?

140 °C 240 °C 340 °C 440 °C

4 The formula of sodium chloride is NaCl.

What would you expect the formula of the following compounds to be?

a) Lithium fluoride

b) Potassium bromide

c) Sodium iodide



3 Electrons and shells (Higher tier only)

1 a) Sketch the electron arrangement in the chlorine atom (proton number 17).

b) Some day scientists hope to make enough atoms of the element with proton number 117 to investigate its properties. It should come below astatine in the Periodic Table. Why will the new element belong to the halogens group?

2 Look at the patterns in the physical properties of the Group 7 elements (see Q3 in Activity 2 above). What change occurs in the elements down the group that can explain the patterns?

3 The general formula of the salts of the alkali metals and the halogens is MX, where M stands for the metal and X for the halogen. Why do the halogens form these compounds with the same general formula?



c4_10 Patterns in Group 7

Resources


Files on Teacher Pack CD: c4_10_practical; c4_10_technician

Equipment for demonstration and for class practical

Learning outcomes C4. 1.25 understand that the halogens become less reactive as the group is descended and give examples to show

this

C4.1.26 understand how a trend in reactivity for halogens can be shown by their displacement reactions and by

their reactions with alkali metals and with iron

C4.1.27 state and explain the safety precautions necessary when working with the halogens


arrangement, illustrated by the electron configurations of the atoms of elements in Group 7

Literacy focus: Recording observations.


ICT focus: Using video clips of reactions of halogens; using the internet to research the uses of halogens and their

compounds.


� explain that the halogens become less reactive down the group

� understand the safety precautions needed for working with halogens.

Key vocabulary

corrosive �� toxic �� halide �� displacement reaction


It is easy for students to confuse halogens and halides. Emphasise what the endings of the names mean.


� Fluorine was the last of the common halogens to be separated from its compounds. Ask students to suggest

reasons for this. (It is very reactive and extremely difficult to separate from compounds; and when separated it

will react with most elements and compounds, including glass.) Discuss what precautions need to be taken when

dealing with fluorine. (Look into the story of Henri Moissan and the discovery of fluorine.)

Learning activities practical c4_10 Low demand � Demonstrate the reactions of chlorine, bromine and iodine with sodium and iron (see Technician

sheet c4_10) or show video clips of the reactions. Students should record their observations. Ask students,

individually or in groups, to discuss what the products of the reactions are, what the word equations are, and what

they show about the relative reactivity of the halogens.

Teaching and learning notes: Students should recall the reaction of sodium with chlorine from a previous lesson

and could be asked to predict the products of the reaction of sodium with bromine and iodine.

Standard demand � Give out the practical worksheet, which gives instructions for carrying out displacement

reactions of chlorine, bromine and iodine. Health and Safety: Ensure that the room is well ventilated and draw

students’ attention to the safety procedures.

Students should observe that chlorine will displace both bromine and iodine from their compounds, whereas

bromine will only displace iodine. Ask students to predict the effect of using fluorine or astatine and their

compounds in displacement reactions. Display examples of the balanced chemical equations for displacement

reactions and ask questions to assess whether students can interpret them.

If time allows, ask students to research the uses of halogens and their compounds.

Teaching and learning notes: Students should be able to work out the order of reactivity from the displacement

reactions. Note that this is the same as the order in the Periodic Table.

High demand � Discuss the relationship between electron arrangement and the order of reactivity in the group.

c4_10 Patterns in Group 7 continued


Plenary suggestions Discuss how the uses of halogens and their compounds are related to their properties, i.e. halogens are used as

disinfectants because of their high reactivity and ability to destroy micro-organisms; compounds containing

halogens are generally unreactive (e.g. salt, PVC).

Student Book answers Q1 a) potassium + chlorine → potassium chloride b) iron + bromine → iron bromide

Q2 Halogen vapours are toxic and corrosive and must be cleared out of the air quickly.

Q3 a) Three b) One

Q4 Bromine will displace iodine because iodine is lower in the group. Potassium bromide and iodine are formed.

Q5 NaI

Q6 a) 2Fe(s) + 3Br2(l) → 2FeBr3(s) b) Cl2(g) + 2KI(aq) → 2KI(aq) + I2(aq)

(State symbols were not asked for, so are not essential.)

Q7 No, astatine is at the bottom of the group / has the largest atoms, so will not displace the other halogens.

Practical sheet answers Q1 Good ventilation because the gases are toxic.

Eye protection/goggles because the halogens are corrosive/irritant.

Gloves because the halogens are corrosive/irritant.

Q2 a) Sodium chloride, iron chloride, iron bromide, iron iodide. (Ignore the oxidation state of iron.)

b) sodium + chlorine → sodium chloride

iron + chlorine → iron chloride; iron + bromine → iron bromide; iron + iodine → iron iodide

c) Each of the halogens reacts (exothermically) with iron. The product is a red-brown solid in each case.

d) Chlorine – bromine – iodine

e) Fluorine will react violently with iron, producing a red-brown solid (iron(III) fluoride).

Q3 a) chlorine + bromide (from almost colourless to pale orange)

chlorine + iodide (from almost colourless to yellow/brown)

bromine + iodide (from pale orange to yellow/brown)

b) The halogens can be identified by their colour, and the halide formed is identified by deduction:

sodium chloride + bromine; sodium chloride + iodine; sodium bromide + iodine

c) chlorine + sodium bromide → sodium chloride + bromine

chlorine + sodium iodide → sodium chloride + iodine

bromine + sodium iodide → sodium bromide + iodine

d) The halogen has displaced the halide, so the reactant halogen is more reactive than the halogen formed.

e) chlorine – bromine – iodine. The order is the same as for 2(d).

Q4 a) Fluorine should be above chlorine. Reactivity decreases down the group. Fluorine is at the top of the group.

b) There would be a reaction. Chlorine (green gas/solution) would be formed along with sodium fluoride.

c) fluorine + sodium chloride → sodium fluoride + chlorine

Extension

Q5 For example, disinfectants, polymers (PVC, PTFE), anaesthetics, flavouring (sodium chloride).

(Higher tier only)

Q6 As the size of the atoms / number of electron shells / distance of outer electrons from the nucleus increases,

the reactivity decreases.

Q7 2Na(s) + Cl2(g) → 2NaCl(s) 2Fe(s) + 3Cl2(g) → 2FeCl3(s)

2Fe(s) + 3Br2(l) → 2FeBr3(s) 2Fe(s) + 3I2(s) → 2FeI3(s)

Q8 Cl2(aq) + 2NaBr(aq) → 2NaCl(aq) + Br2(aq) Cl2(aq) + 2NaI(aq) → 2NaCl(aq) + I2(aq)

Br2(aq) + 2NaI(aq) → 2NaBr(aq) + I2(aq)

Q9 F2(aq) + 2NaCl(aq) → 2NaF(aq) + Cl2(aq)




P Displacement reactions of chlorine, bromine and iodine

Objectives

In this activity you will:

� collect and record observations of the reactions of some of the halogens

� determine the order of reactivity of the halogens.

You must wear eye protection at all times when carrying out this activity.

The halogens are TOXIC and CORROSIVE. Avoid breathing in fumes of chlorine and bromine.

Avoid getting solutions of chlorine, bromine or iodine on your skin – they may irritate or stain.


6 test tubes • test tube rack • droppers • about 6 cm3 each of all solutions

Halogens: chlorine solution, Cl2(aq); bromine solution, Br2(aq); iodine solution, I2(aq)

Halides: sodium chloride solution, NaCl(aq); sodium bromide solution, NaBr(aq);

sodium iodide solution NaI(aq)

Method

1 Collect the apparatus. Pour about 6 cm3 of each of the halogens into separate test tubes. Pour about 2 cm3 of each of the halides into separate test tubes. Make sure you know which substance is in each test tube. Record the appearance of each of the six materials in a suitable table.

2 Using a dropper, add about 2 cm3 of the chlorine solution to each of the test tubes containing the halide. Shake each mixture gently. Record what you see.

3 Wash out the test tubes in which you mixed the reactants. Pour about 2 cm3 of each of the halides into separate test tubes.

4 Using a dropper, add about 2 cm3 of the bromine solution to each of the test tubes containing the halide. Shake each mixture gently. Record what you see.

5 Wash out the test tubes in which you mixed the reactants. Pour about 2 cm3 of each of the halides into separate test tubes.

6 Using a dropper, add about 2 cm3 of the iodine solution to each of the test tubes containing the halide. Shake each mixture gently. Record what you see.

Results

Record (in a suitable table) the appearance of each solution before mixing, and the appearance after pairs of solutions have been mixed.



Questions

1 What safety precautions are required when working with the halogens? Explain your answers.

2 a) Name the products of the reactions of the halogens that you have seen.

b) Write word equations for these reactions.

c) What are the similarities in the reactions of the halogens with iron?

d) From the reactions of the halogens that you have observed, put chlorine, bromine and iodine in order of reactivity, with the most reactive first.

e) Suggest what you might see if fluorine was reacted with iron.

3 Look at the results from your experiments on the displacement reactions of the halogens and halides.

a) In which pairs of reactants did a chemical reaction take place?

b) What do the colour changes tell you were the products in these cases?

c) Write word equations for the reactions.

d) What does each reaction tell you about the reactivity of the halogens involved in the reaction.

e) Put chlorine, bromine and iodine in order of reactivity. Does your order agree with your answer to Q2 part (d)?

4 a) Where do you think fluorine should come in the order of reactivity of the halogens? Explain your answer.

b) What would you expect to happen if a fluorine solution was added to sodium chloride?

c) Write a word equation for the reaction of fluorine solution, F2(aq), with sodium chloride solution.

Extension

5 Research (in books and on the internet) the uses of the halogens and their compounds.

(Higher tier only)

6 What is the relationship between the electron arrangement of the halogen atoms and the order of reactivity of the halogens?

7 Write balanced chemical equations (with state symbols) for the reactions in your answer to Q2 part (b).

8 Write balanced chemical equations (with state symbols) for the reactions in your answer to Q3 part (c).

9 Write a balanced chemical equation (with state symbols) for the reaction of fluorine solution, F2(aq), with sodium chloride solution as in Q4 part (c).




Technician sheet


Demonstration of reactions of halogens

� sample of chlorine gas (if a member, see CLEAPSS recipe for generation of chlorine by reaction of concentrated hydrochloric acid and potassium manganate(VII))

� bromine

� iodine

� sodium

� iron wool

� 2 gas jars and lids

� white tile and knife

� 2 combustion spoons

� 2 boiling tubes and bungs

� dropper

� clamp stand and clamp

� tongs

� glass rod

� Bunsen burner

Class practical on displacement reactions

Each group of students will need:

� 6 test tubes

� test tube rack

� dropper

� goggles / eye protection for all (gloves and lab coats optional)

and about 6 cm3 each of the following:

� chlorine water (chlorine dissolved in water; if a member of CLEAPSS, see Hazcard 22)

� bromine water (bromine dissolved in water; if a member of CLEAPSS, see Hazcard 15B)

� iodine solution (iodine in potassium iodide solution; if a member of CLEAPSS, see Hazcard 54B)

� sodium or potassium chloride solution (0.1 mol/dm3)

� sodium or potassium bromide solution (0.1 mol/dm3)

� sodium or potassium iodide solution (0.1 mol/dm3)

Method


Wear eye protection. Prepare a gas jar of chlorine (TOXIC) by the large-scale method (if a member of CLEAPSS, see Hazcard 22). The gas does not have to be specially dried. This may be done earlier and stored in a fume cupboard. Place a lid on the gas jar.

Wear eye protection. Cut a piece of sodium (a 6–8 mm cube is sufficient) and remove the oil by squeezing the sodium between pieces of absorbent paper. Place the sodium on the flat side of a brick in a fume cupboard. Direct a non-luminous Bunsen burner flame onto the sodium so that it melts and any remaining oil burns off. When the sodium finally burns



with a small yellow flame, quickly remove the lid and place the gas jar over the burning sodium. If a member of CLEAPSS, refer to Hazcard 88.

Place some iron wool on a combustion spoon. Heat the iron wool in the Bunsen flame until it is glowing, then transfer it quickly to a gas jar of chlorine. If a member of CLEAPSS, refer to Hazcard 22A.

Put a couple of drops of bromine in a boiling tube and clamp it at about 45° to the vertical. Place a piece of iron wool about halfway down the tube. Heat the iron wool gently at first then more strongly until a reaction starts. If a member of CLEAPSS, refer to Hazcard 15A and Handbook section 13, page 1301.

Repeat the procedure in the previous paragraph with a small amount of solid iodine in the boiling tube instead of the bromine.


Full instructions for the class practical are given on Practical sheet c4_10.

Notes


� The reaction of sodium with chlorine may have been demonstrated in previous lessons. The sodium should continue to burn or glow in the chlorine gas, producing a brownish smoke (the brown is caused by the oil remaining on the sodium).

� Iron will glow in the chlorine, provided it is hot enough, producing a reddish-brown smoke of iron(III) chloride.

� Hot iron wool will glow and spark in the bromine, producing reddish brown smoke of iron(III) bromide mixed with bromine vapour.

� Purple iodine vapour will form and react with the heated iron to form a reddish brown smoke, but the reaction will be noticeably less vigorous than with bromine.

� Video clips of these reactions can be found on the internet.


� Chlorine water is colourless unless very concentrated, when it appears very pale green but also has a very powerful chlorine odour. Bromine solution is pale orange (red if more concentrated). Iodine in potassium iodide solution is yellow-brown in colour (dense brown if concentrated). The halides are all colourless.

Students should observe a colour change in the following combinations:

� chlorine + bromide (from almost colourless to pale orange)

� chlorine + iodide (from almost colourless to yellow-brown)

� bromine + iodide (from pale orange to yellow-brown)

Health and Safety

� Chlorine, bromine and iodine are TOXIC and CORROSIVE both in pure form and in solution. The demonstrations should be carried out in a fume cupboard and students should keep 2 m away.

� Sodium is extremely REACTIVE with water. Make sure that the work area is dry.

� The class experiment should be carried out in an adequately ventilated laboratory. Students must wear eye protection and lab coats.



c4_11 Ionic compounds

Resources

Student Book pages 122−123 � Interactive Book: Quick starter ‘How do we know sugar isn’t ionic?’; Naked Scientist animation ‘What is ionic bonding?’; Matching pairs ‘Identifying anions’ � Homework pack c4_11


Equipment for demonstration

Learning outcomes C4.3.1 understand that molten compounds of metals with non-metals conduct electricity and that this is evidence

that they are made up of charged particles called ions

C4.3.2 understand that an ion is an atom (or group of atoms) that has gained or lost electrons and so has an

overall charge

C4.3.3 account for the charge on the ions of Group 1 and Group 7 elements by comparing the number and

arrangement of the electrons in the atoms and ions of these elements

C4.3.6 recall that compounds of Group 1 metals with Group 7 elements are ionic

Ideas about Science IaS 3.1 scientific hypotheses, explanations and theories are not simply summaries of the available data – they are

based on data but are distinct from them

IaS 3.2 an explanation cannot simply be deduced from data, but has to be thought up creatively to account for the

data

IaS 3.3 a scientific explanation should account for most (ideally all) of the data already known – it may explain a



ICT focus: Accessing video clips of electrolysis of molten salts, and models of ions and ionic compounds from the

internet.


� explain how Group 1 and 7 elements form ions

� explain how the properties of the compounds of Groups 1 and 7 depend on ions

� explain how ideas are developed to account for the properties of compounds.

Key vocabulary

ion �� ionic compound


It requires imagination to move from concrete experimental data (conduction by molten salts) to the visualisation of

charged particles moving in the liquid. Students will also have difficulty understanding that the loss of a negatively

charged electron leaves a positive charge on the ion (think: minus a minus makes a positive).


� Show pictures of sodium metal, chlorine gas and sodium chloride (see Student Book p. 122) and ask students to

note the properties of each and suggest reasons why sodium chloride is different to the elements.

Learning activities worksheet c4_11 Low demand � Show a sample of sodium chloride (preferably a large crystal or pictures of cubic crystals). Ask

students to recall the properties of sodium chloride (hard, brittle, high melting point, soluble in water). Tell them

that, to explain these properties, we need to know what sodium chloride looks like at a nanometre scale. Because

we cannot do this (directly), we have to use our imaginations and some evidence. Demonstrate that a molten salt

conducts electricity (see Technician sheet c4_11) or show a video from the internet. Note that the salt only

conducts when it is a liquid; this tells us that something is conducting the electricity only when it can move. These

things are called ions and they have charges. Students should attempt the questions in activity 1 on the worksheet.

Teaching and learning notes: If students have not yet studied Module P5, they may be a little uncertain about

what electricity is. Remind them of familiar experiments with charges (rubbing pieces of plastic to pick up bits of

paper, making hair stand on end, sticking balloons to walls) and that electricity is a flow of charge.

Standard demand � Discuss why the fact that molten salts conduct electricity shows that salts are made up of

charged particles called ions. Recall that atoms are made up of protons (+) and electrons (–). Explain what ions are

c4_11 Ionic compounds continued


and explain the formation of positive and negative ions by loss and gain of electrons, respectively. Ask students

whether they think this explanation can explain why salts conduct electricity? Note that the explanation is not the

same as the evidence, and that the scientists who thought up the explanation had to use their imaginations as well

as the evidence they had. Note that a good explanation can make predictions, e.g. that other compounds,

particularly of Groups 1 and 7, are also ionic and will conduct electricity when molten. Confirm that this is true.

Activity 2 on the worksheet works through these ideas.

Teaching and learning notes: Students will need to recall atomic structure from the start of this module and

should learn that Group 1 elements form ions with a 1+ charge while Group 7 form ions with a 1– charge.

High demand � Show students diagrams of the electron arrangement of sodium atoms and chlorine atoms. Ask

pairs or groups of students to explain what occurs in the reaction of sodium with chlorine. Choose a group to

present their response and see if other groups agree or disagree. The Student Book (p. 123) gives the explanation.

Discuss the significance of the electron arrangement of the ions (it is like that of Group 0 elements) and for Higher-

tier students note that this could be used to predict the charges on the ions of elements other than those in Groups

1 and 7 (e.g. magnesium oxide, calcium sulfide).

Plenary suggestions Ask students to state the difference between ions and atoms: ions have charges but atoms do not; ions only occur

in compounds; only in atoms does the number of protons equal the number of electrons. Also ask students to give

the formulae (with charges) of ions in various salts of Group 1 and 7 elements, e.g. sodium iodide, Na+ and I

–.

Student Book answers Q1 It conducts electricity when molten. Q2 Any ionic compound

Q3 Chlorine atoms have gained an electron, so getting an extra negative charge / have one more electron than

the number of protons in the atom.

Q4 K+ and Br

– Q5 Li

+ 2; F

– 2.8

Q6 Yes. Group 0 gases are unreactive, so ionic compounds are unreactive too. Or answers linking the stability of

noble gases to the stability of the electron arrangement of the ions.


Q1 a) The bulb will light up. b) No c) It melts. d) Yes

e) Cathode: grey bits of metal formed. Anode: bubbles of gas formed.

f) Zinc chloride g) Cathode – zinc; anode – chlorine (or alternative answers if other salts are used)

Q2 … particles … current … atoms … negative … ions … move … salt … melted …


Q1 Describing a scientific theory – Lily; describing some evidence – Harry; making a prediction – Dilshan;

suggesting that being creative is necessary in science – Sonya; acting like a scientist – all four.

Q2 a) False. An ionic compound conducts electricity when it is in the liquid state. b) True

c) False. A positively charged ion is an atom that has lost an electron. OR A negatively charged ion is an atom

that has gained an electron.

d) False. Potassium bromide is made up of K+ and Br

– ions. e) True f) True


Q1 a) The outer shell electron in a potassium atom is transferred / moves to the outer shell of a chlorine atom.

b) Correct diagrams with three shells, electron arrangements of 2.8.8 for both ions, with the correct charges

shown outside brackets.

c) They are both the same / They are the same as the noble gases / argon. (Or any other correct statement.)

d) The potassium ion has lost a negative charge, so is left with a positive charge / has 19 protons but only 18

electrons. The chlorine ion has gained an extra negative charge / has 17 protons and 18 electrons.

Q2 a) It explains all the known data and makes predictions that can be tested.

b) Disagree. A theory is more than just a summary. It is an explanation for the data that requires imagination /

creativity from the scientist, e.g. to visualise the charged particles in the ionic compound moving when it is

molten.




Technician sheet


Demonstration – conductivity of molten salts

� power pack (0–12 V), leads and crocodile clips

� 2 graphite electrodes (fixed in holder or bung about 1 cm apart)

� ammeter or bulb (3 V) in holder

� crucible

� clamp stand and clamp

� Bunsen burner, tripod and pipeclay triangle

� zinc chloride powder

Method

Set up the electrolysis circuit (see the Student Book p. 122) – an ammeter could be used instead of the lamp/bulb. If a member, refer to CLEAPSS Guide L195 page 37 and Handbook section 11.4, page 1125.

Turn on the power and test the circuit (e.g. by shorting out the electrodes briefly with a metal spatula). Turn off the power.

Fill the crucible to within 0.5 cm of the brim with zinc chloride powder. Place the crucible on the triangle on the tripod. Insert the electrodes into the powder, clamp the bung holding the electrodes in position, and turn on the power. Turn up the power progressively to 6 V (there should be no conduction).

Light the Bunsen burner and place it under the crucible.

When the lamp/bulb or ammeter reveals that a current is flowing, allow students to come closer to see what is happening in the crucible (or use a video camera to project pictures of the experiment).

Turn off the Bunsen burner and the power. Lift the electrodes out of the molten zinc chloride. Allow the crucible to cool. (The electrodes could be left in the zinc chloride as it cools to show that the conductivity decreases, but it will have to be re-melted to free them.)

Notes

� Zinc chloride is used as it has a relatively low melting point (the alkali metal halides cannot be melted by a Bunsen burner).

� It will take a minute or so for the zinc chloride to start melting and for electrolysis to commence. Students should be given other tasks to do while they wait, while keeping one eye on the lamp/bulb (or on the projected image if a camera is being used).

� Bubbling (chlorine gas) should be seen at the anode (positive electrode). It could be tested using moist indicator paper, but this is difficult given the proximity to the hot zinc chloride. Do not ask students to smell the gas.

� Crystals of solid zinc may be formed in the cathode (negative electrode) or solid zinc may be found in the crucible when the solidified zinc chloride is dissolved out.

Health and Safety

� Molten zinc chloride is HOT. The electrolysis produces chlorine gas, which is TOXIC. The experiment must be carried out in a fume cupboard. Students should keep away from the fume cupboard except for a brief examination of the crucible carried out with the front of the fume cupboard closed. If a member, refer to CLEAPSS Guide Cl95.




1 Conducting salts

1 Watch the demonstration of the experiment to see if salts conduct electricity.

a) How will you know if the salt completes the electric circuit?

……………………………….………………………………….…..………………………….

……………………………………………………………

b) Does the salt conduct electricity when it is solid?

………………………….

c) What happens to the salt when it is heated?

..………………………………………………

d) Does the liquid salt conduct electricity?

…………………………………………..

e) What do you see at the cathode (–)?

.…………………………………………

What do you see at the anode (+)?

………………………………………….

f) Label the diagram with the name of the salt that was used.

g) Name the substance formed at the cathode (–). …………………………………

Name the substance formed at the anode (+). …………………………………

2 Fill in the missing words in the following sentences. Choose from the words below:

move current negative ions salt melted particles atoms

When any charged ………………… move they produce an electric ………………… A

salt such as sodium chloride is made up of ions. Ions are formed from ………………

that have been given a positive or ………………… charge. In the solid state the

………………… cannot ………………… so an electric current cannot pass through the

………………… If the salt is ………………… the ions can move and the current flows

through the salt.



2 Ionic explanations

1 Read the comments made by the four students. Then put ticks in the table below to identify who is making each type of statement.

Dilshan Harry Lily Sonya None of them

Who is describing a scientific theory?

Who is describing some evidence?

Who is making a prediction?

Who is suggesting that being creative is necessary in science?

Who is acting like a scientist?

2 Are the following sentences true or false? Write corrections for the false statements.

a) An ionic compound conducts electricity when it is in the solid state.

b) Sodium chloride is an ionic compound.

c) A positively charged ion is an atom that has gained an electron.

d) Potassium bromide is made up of K– and Br+ atoms.

e) A salt does not have an overall charge because the number of positive charges is equal to the number of negative charges.

f) Compounds of metals and non-metals are ionic.

Salts, such as sodium chloride, are

made up of charged particles called ions,

which can move when the salt is molten.

A light bulb in an electric circuit containing a

molten salt will light up when the electricity is

turned on.

Other salts of Group 1 and 7 elements

should also conduct electricity when they

are liquid.

You can imagine the charged sodium and

chloride ions moving around in the molten

state and being attracted to the electrodes.

Lily

Sonya Dilshan

Harry



3 Forming ions

1 The electron arrangements in atoms of potassium (proton number 19) and chlorine (17) are as shown below.

a) Explain how the ions in potassium chloride are formed when potassium atoms react with chloride atoms.

b) Sketch diagrams of the electron arrangement of the potassium ions and chloride ions in potassium chloride.

c) What do you notice about the electron arrangements of the potassium and chloride ions?

d) Explain why each of the ions have their particular charges.

2 “Salts such as sodium chloride formed from a metal and non-metal are made up of charged particles called ions, which are atoms that have lost or gained electrons.”

a) Why is the statement above considered to be a good theory?

b) A scientific theory is a summary of the data collected in experiments. Do you agree or disagree with this statement? Explain your answer by referring to the theory of ionic compounds.



c4_12 Understanding ions

Resources

Student Book pages 124−125 � Interactive Book: Quick starter ‘Which formulas are wrong?’; Practical investigations ‘Making salt in space’; Quick starter ‘Predicting the formula of calcium fluoride’ � Homework pack c4_12


Equipment for demonstration

Learning outcomes C4. 3.4 work out the formulae of ionic compounds given the charges on the ions

C4.3.5 work out the charge on one ion given the formula of a salt and the charge on the other ion

C4.3.7 understand that solid ionic compounds form crystals because the ions are arranged in a regular lattice

C4.3.8 describe what happens to the ions when an ionic crystal melts or dissolves in water

C4.3.9 explain that ionic compounds conduct electricity when molten or when dissolved in water because the ions

are charged and they are able to move around independently in the liquid

Literacy focus: Using the vocabulary of ionic compounds – ions, lattice, crystalline, positive, negative, charge, etc.

Numeracy focus: Using ideas of ratios in the context of the formulae of ionic compounds (Higher tier only).

ICT focus: Accessing video clips of electrolysis of solutions and animations of models of ionic lattices and ions in

solutions from the internet.


� explain how ionic compounds form crystalline solids

� explain what happens when ionic compounds melt or dissolve in water

� work out the formulae of ionic compounds and the charges on ions.

Key vocabulary

crystal lattice


Some students may struggle to understand how to balance the charges between ions (required at Higher tier) if

they are not familiar with the concept of lowest common multiples.


� Show students samples or pictures of crystalline ionic compounds (e.g. sodium chloride, copper sulfate, alum,

calcite). Ask students to list the properties of crystalline materials – regular shape, same basic shape (e.g. NaCl

always cubic), shiny, flat surfaces. Ask students to discuss what the crystal shape tells us about the particles

making up the material and report their ideas. Hopefully, they will say that it tells us that the particles are

arranged in regular patterns. If you wish, show X-ray diffraction pictures of crystals (e.g. NaCl) and explain that

the pattern of spots shows that the particles are in a particular pattern and crystallographers can work it out.

Learning activities worksheet c4_12 Low demand � Remind students that ionic compounds are made up of positively and negatively charged ions.

Explain that the crystals are formed by ions building up a regular shape called a crystal lattice. It would be useful to

be able to assemble a model of sodium chloride and show how packing the ions together alternately produces a

cubic shape. If such a model is not available, show ready-made models or pictures. There are various models on

the internet, many of which rotate to show the three-dimensional structure. Note that many of the tiny cubes build

up like bricks in a wall to form crystals large enough to see and hold. There are tasks on this topic on the

worksheet.

Teaching and learning notes: Explain that scientists use models (both three-dimensional, two-dimensional and

imaginary) to help develop, explain and investigate theories.

Standard demand � Having explored the lattice structure of ionic compounds (see above), demonstrate the

conduction of electricity by sodium chloride solution (see Technician sheet c4_12 and the Student Book p. 125,

Figure 4). Having made observations, ask students in small groups to provide an explanation for what is

happening. Note that the effect of the water is the same as melting the salt – it makes the ions mobile and free to

be attracted to the electrodes, hence conducting the current. Students should complete activity 2 on the worksheet.

c4_12 Understanding ions continued


Teaching and learning notes: Students will need to recall the demonstration from the previous lesson on the

conductivity of molten salts.

High demand � (Higher tier only) Explain that the charges on ions provide a method of working out the formula of

ionic salts. Note the rule that the positive and negative charges should balance. Demonstrate the calculation for

cases when the charges on the ions are not equal, involving the lowest common multiple of the charges of each ion

(Student Book p. 125 gives one method). Show also how the method can be worked backwards to find the charge

on an unknown ion, given the formula and charge on the other ion. (There are various other ways of showing the

method using cards, model ions with hooks, etc.)

Plenary suggestions Ask students to write a question and its answer, on separate slips of paper, about any aspect of ionic compounds

covered in the last two lessons. Collect in the questions and answers separately, shuffle them and hand a question

and answer to each student. One student then reads out a question. The student with the answer reads that out,

and then reads the next question, and so on. There may be some duplicates but that does not really matter.

Student Book answers Q1 The crystals have a regular shape.

Q2 The ions are arranged in a lattice.

Q3 The ions move out of their fixed lattice positions.

Q4 The ions can move freely in a solution but are fixed in position in the solid.

Q5 a) NaBr b) MgO c) K2O d) Al2S3

Q6 2+


Q1 a) a cube / cubic b) B (crystal lattice)

c) The ions are always arranged in the same lattice shape.

Q2 a modelling activity


Q1 a) Diagram should show a complete circuit when the salt solution conducts, like Figure 4 on p. 125 in the

Student Book, but no ‘heat’, and labelled with ‘sodium chloride / salt solution’ rather than ‘molten salt’.

b) The bulb will light up / the ammeter will register a current.

c) Observations should include: solid salt does not conduct; as water is added, the bulb begins to brighten;

fizzing at both electrodes.

d) When the salt dissolves, the ions become free to move and conduct the current.

e) Both melting and dissolving cause the ions to break out of the crystal lattice and become mobile.

Q2 A is an ionic compound but B is not.


Q1 a) LiBr b) KF c) CaO d) Na2S e) MgCl2

f) AlBr3 g) CaI2 h) Li2O i) KCl j) Al2O3

Q2 a) 1+ b) 2+ c) 2+ d) 3– e) 3+




Technician sheet


Demonstration – conduction by solutions of salts

� samples of crystals of ionic compounds that show a regular shape, e.g. copper sulfate, alum, sodium chloride, calcite (calcium carbonate)

� power pack (0–12 V)

� 2 carbon or carbon-fibre electrodes, preferably fixed in a holder or bung, 1–2 cm apart (alternatively, an electrolysis cell can be used with the connection to the electrodes made beneath the cell)

� leads and crocodile clips

� 3 V bulb in holder or ammeter (0–1 A)

� 100 cm3 beaker (or electrolysis cell)

� wash bottle of water, or a beaker and a dropper

� sodium chloride (fine crystals)

� (optional) other ionic solids, e.g. copper sulfate (fine crystals)

Method

Set up the electrolysis circuit (see Student Book p. 125) and test that the light bulb/ammeter works.

About one-quarter fill the beaker with sodium chloride (sufficient to be able to insert the electrodes into the powder).

Place the electrodes into the solid and turn on the power (6 V).

Add water, a dropper full at a time.

Continue adding water until it can be seen that a current is flowing.

Turn off power before dismantling the apparatus.

Notes

� The dry solid sodium chloride will not conduct.

� As water is added, the bulb will gradually light up as the solution is formed.

� Students should notice fizzing at the electrodes, which is a sign that a chemical change is taking place and that a current is flowing, but do not focus on this in this lesson.

Health and Safety

� Electrolysis of sodium chloride solution produces chlorine gas, which is TOXIC. Carry out the experiment in a fume cupboard or an area of the laboratory with good ventilation.

� Avoid contact between the electrodes, which can cause sparks.




1 Crystals

1

a) What shape is the arrangement of the ions in a sodium chloride crystal?

………………………………………

b) Which phrase is used for the arrangement of the ions in an ionic compound?

A crystal trellis B crystal lattice C crystal ladder D crystal frame

c) Why are sodium chloride crystals always the same shape? .………………..…….

.………………………………………………………………………………………………

2 Make a model of a sodium chloride crystal. Cut around the shape below, and fold along the dotted lines to form a cube. Paste the flaps to stick the structure together.

c4_12 Understanding ions continued


2 Ionic compounds in solution

1 Observe the experiment showing the effect of water on the conduction of electricity by sodium chloride.

a) Draw a diagram of the apparatus used.

b) How will you know if the salt solution is conducting electricity?

c) Describe what you observed during the experiment.

d) What is your explanation for the observations? (Hint: remember what happened when an ionic compound was melted.) You could use your diagram to show what is happening.

e) Explain why dissolving and melting are similar.

2 Substances A and B are both white solids that dissolve in water. A solution of A conducts electricity but a solution of B does not.

What do these observations tell you about substances A and B?

3 Ions and formulae (Higher tier only)

1 Use the table below to work out the formulae of the compounds that follow.

Metal ions Non-metal ions

Name Formula Name Formula

lithium Li+ fluoride F–

sodium Na+ chloride Cl–

potassium K+ bromide Br–

magnesium Mg2+ iodide I–

calcium Ca2+ oxide O2–

aluminium Al3+ sulfide S2–

a) lithium bromide b) potassium fluoride

c) calcium oxide d) sodium sulfide

e) magnesium chloride f) aluminium bromide

g) calcium iodide h) lithium oxide

i) potassium chloride j) aluminium oxide

2 Use the table in question 1 to help you answer the following questions.

a) Rubidium chloride has the formula RbCl. What is the charge on a rubidium ion?

b) The formula of copper oxide is CuO. What is the charge on the copper ion?

c) Barium combines with bromine to form an ionic compound with the formula BaBr2. What is the charge on the barium ion?

d) A compound of magnesium and nitrogen has the formula Mg3N2 and is made up of magnesium and nitride ions. What is the charge ion the nitride ion?

e) The iron oxide used to produce the metal ion is an ionic compound with the formula Fe2O3. What is the charge on the iron ion?

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