ccea.org.uk · Web view1.2.12 demonstrate knowledge and understanding that metallic bonding results from the attraction between the positive ions in a regular lattice and the delocalised

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CCEA Planning Framework for GCSE Chemistry

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3

GCSE Chemistry

Contents

Page

Introduction

1

Unit 1.1: Atomic structure

4

Unit 1.2: Bonding

10

Unit 1.3: Structures

16

Unit 1.4: Nanoparticles

24

Unit 1.5: Symbols, formulae and equations

33

Unit 1.6: The Periodic Table

38

Unit 1.7: Quantitative chemistry

48

Unit 1.8: Acid, bases and salts

57

Unit 1.9: Chemical analysis

70

Unit 1.10: Solubility

81

Unit 2.1: Metals and reactivity series

83

Unit 2.2: Redox, rust and iron

93

Unit 2.3: Rates of reaction

98

Unit 2.4: Equilibrium

103

Unit 2.5: Organic Chemistry

106

Unit 2.6: Quantitative Chemistry

123

Unit 2.7: Electrochemistry

129

Unit 2.8: Energy Changes in Chemistry

132

Unit 2.9: Gas Chemistry

134


Introduction

The purpose of this Planning Framework is to support the teaching and learning of GCSE Chemistry. The Planning Framework is based on specification content but should not be used as a replacement for the specification. It provides suggestions for a range of teaching and learning activities which provide opportunities for students to develop their:

· Knowledge and understanding

· Subject specific skills

· The Cross-Curricular Skills

· Thinking Skills and Personal Capabilities

The Planning Framework is not mandatory, prescriptive or exhaustive. Teachers are encouraged to adapt and develop it to best meet the needs of their students.

Subject Skills Assessed through Chemistry:

The following skills are assessed in GCSE Chemistry:

· planning an experiment

· analysis of experimental data

· drawing conclusions from experimental data

· evaluation of practical procedures and experimental data

· problem solving

· practical work including measuring and observations

· mathematical skills including graphical work

Supporting the Development of Statutory Key Stage 4 Cross-Curricular Skills and Thinking Skills and Personal Capabilities

This specification builds on the learning experiences from Key Stage 3 as required for the statutory Northern Ireland Curriculum. It also offers opportunities for students to contribute to the aim and objectives of the Curriculum at Key Stage 4, and to continue to develop the Cross-Curricular Skills and the Thinking Skills and Personal Capabilities. The extent of the development of these skills and capabilities will be dependent on the teaching and learning methodology used.

Cross-Curricular Skills at Key Stage 4

Candidates will have the opportunity to develop Using Mathematics through the use of scientific notation including chemical symbols and formulae and in calculation work in quantitative chemistry, use of different units including those involved in nanomaterials and graphical work in solubility and rates of reaction. Communication is assessed through different topics including risks and benefits of nanoparticles, discussion of structure and bonding in materials as well as the preparation of potable water.

Thinking Skills and Personal Capabilities at Key Stage 4

Although not statutory at Key Stage 4 this specification also allows opportunities for further development of the Thinking Skills and Personal Capabilities of Managing Information and Creativity.

Candidates will have the opportunity to plan work, set personal goals and targets to meet deadlines, monitor, review and evaluate progress as well effectively manage their time. They will also work with others to participate in effective teams, listening actively to others and influence group decisions in discussions for example in the reactions of acids with predicted observations. Candidates can identify and analyse relationships in terms of solubility and propose justified explanations in terms of observations for known cation and anion tests. They can analyse and evaluate perspectives and explore unfamiliar views without prejudice as with the benefits and risks of nanoparticles. Candidates may also weigh up options and justify decisions as well as applying and evaluating a range of approaches including salt preparation and methods of separating mixtures based on the properties of the components of the mixture.

Key Stage 4 Statutory Skills and Personal Capabilities

Communication SkillsComm - T&L (Talking & Listening) W (Writing) R (Reading)

Using Mathematics UM

Using ICT UICT

Problem solving PS

Working with Others WO

Self-Management SM

Assessment For Learning

When reference is made to past paper questions, teachers will recognise opportunities for formative assessment activities.

Key Features

The Planning Framework:

· Includes suggestions for a range of teaching and learning activities which are aligned to the GCSE Chemistry specification content.

· Highlights opportunities for inquiry-based learning.

· Indicates opportunities to develop subject knowledge and understanding and specific skills.

· Indicates opportunities to develop the Cross-Curricular Skills and Thinking Skills and Personal Capabilities.

· Provides relevant, interesting, motivating and enjoyable teaching and learning activities which will enhance the student’s learning experience.

· There is a range of support available for both

· teachers and students, including specimen papers, mark schemes, planning frameworks, Factfiles and practical manuals. You can download these from our website at www.ccea.org.uk.

144

Planning Framework for GCSE Chemistry

Unit 1

Unit/Option content

Learning Outcomes or Elaboration of Content

Suggestions for Teaching and Learning Activities

Supporting Cross Curricular Skills, Thinking Skills and Personal Capabilities

1.1 Atomic Structure

Students should be able to:

1.1.1

· demonstrate knowledge and understanding of how ideas about the atom changed over time, with reference to:

· the Plum Pudding model;

· Rutherford’s model of a nucleus surrounded by electrons; and

· the discovery of the neutron by Chadwick, leading to today’s model of an atom;

Watch the video (Watch), do the quiz (Think) and further reading (Dig Deeper)

Complete the RSC’s Atomic Detectives Student Worksheets. (Chemists in a social and historical context, Dorothy Warren)

Atomic Structure Timeline: http://atomictimeline.net/index.php

Research the chemists Thompson, Rutherford & Chadwick on the Internet. Create an atomic structure timeline or poster and present to the class

Comm – T&L, W, R

UICT

1.1.2

· describe the structure of an atom as a central positively charged nucleus containing protons and neutrons (most of the mass) surrounded by orbiting electrons in shells;

Explain that we will be using the Periodic Table to help us to visualise what is going on within an atom.

Explain the background to the word atom, from the Greek “atomos”, meaning indivisible. (You may relate the story of John Dalton’s hypothesis using a piece of aluminium cooking foil cut into half and half again and progressing to a molecular model (any element, it does not have to be aluminium) until only one sphere exists. Students are not expected to recall Dalton’s theory, however introducing Dalton’s ideas about the atom and the meaning of the word atom may help the student to visualise the terms ‘Element’, ‘Compound’ and ‘Mixture’)

Unit/Option content




1.1 Atomic Structure (cont.)


1.1.3

· state the relative charges and approximate relative masses of protons, neutrons and electrons;

Students describe the current model of the atom, and provide a summary table of the properties of the sub-atomic particles. (Relative charge, relative mass, where in the atom they are found)

UM, T&L

1.1.4

· define atomic number as the number of protons in an atom;

Students define mass number and atomic number, and demonstrate how these are used to deduce the numbers of protons, neutrons and electrons in an atom

UM

PS

1.1.5

· define mass number as the total number of protons and neutrons in an atom;

1.1.6

· demonstrate knowledge and understanding that an atom as a whole has no electrical charge because the number of protons is equal to the number of electrons;

Discuss the electrical neutrality of atoms in terms of equal numbers of protons and electrons. Refer back throughout to the summary table for the subatomic particles and show how the relative charges links to neutrality and the relative masses ties to mass number

Comm – T&L

Unit/Option content






1.1.7

· calculate the number of protons, neutrons and electrons in an atom or an ion and deduce the charge on an ion or determine the number of subatomic particles given the charge;

Students can now be asked to work out the number of protons, electrons and neutrons in an atom of an element using the atomic number and mass number, identify an atom from the number of protons.

This information can be given in a table with blank spaces to fill in. At this stage students should have relatively easy atoms to identify, avoiding isotopes initially. Progress to introducing some examples where the number of protons and electrons is not equal and discuss the electrical charge which will ensue. Introduce the term ‘ion’ and demonstrate how to deduce the charge on an ion, and how to deduce the number of electrons present if you are given the charge on the ion

UM

PS

1.1.8

· write and draw the electronic configuration (structure) of atoms and ions with atomic number 1–20;

Use the idea that the atomic theory has progressed from identifying the particles in an atom to suggesting positions for them. Give the rules for positioning electrons. In the second and subsequent shells inform students the electrons are placed at 12 o’clock, 3 o’clock, 6 o clock and 9 o’clock before being paired in the same order. This method enables drawing an oxygen molecule easier later in the course.

Draw the electronic structure of selected atoms to illustrate the process to the students.

Supply the students with an A3 sized table with the outline of the Periodic table (elements 1-20) drawn. Students draw the electronic structure of the first 20 elements in the blank boxes using their copy of Periodic table as a source of information of atomic numbers

SM

Unit/Option content






These may be accompanied by tabulated information on number of electrons, protons, neurons and/or the electronic structure written in the format 2,8,8,2. The drawings can later be usefully employed in looking at patterns in the Periodic Table.

Using these diagrams of the first 20 elements in the Periodic Table, guide the students to realising the connection between the group number and the number of electrons in the outer shell (you may wish to illustrate another pattern at this stage, the connection between the number of shells in an atom and the period in which the element is placed). Ask the students to equate the lack of reactivity of the noble/inert gases with the pattern of electrons.

Once it is established that stability is gained by securing a full outer shell, explore how this can be achieved by different atoms by losing or gaining the appropriate number of electrons.

Introduce the terms ‘ion’ and illustrate the appropriate chemical nomenclature: use of square brackets around diagrammatic representations of ions, and the superscript charge. Accompany the diagrams with electronic structures written in the format 2,8,8

PS

1.1.9

· recall that atoms have a radius of about 0.1 nm (1 x 10-10 m) and that the nucleus is less than 1/10000 of that of an atom (1 × 10-14);

Students describe the trends, across a period and down a group, of atomic size in the periodic table. State that radii vary around the size of 0.1 nm. Understand that nuclear radii are less than 1/10000 of the radius of an atom and less than 1 × 10-14 m. Be able to convert between units including nm, m and possibly pm and fm

UM

Unit/Option content






1.1.10

· define isotopes as atoms of an element with the same atomic number but a different mass number, indicating a different number of neutrons;

Students may have already noticed that the relative atomic mass of chlorine on their Periodic Table is not the same as the mass number. Ask them to identify two atoms with the same number of protons but a different number of neutrons (using same tabular approach as was used at 1.1.6) which work out to be the same element. This will lead into a discussion about isotopes

Compare isotopes to chocolate button Easter eggs, they look the same on the outside, they taste the same but count the buttons in the centre and often the number of buttons is variable!

Comm – T&L

1.1.11

· interpret data on the number of protons, neutrons and electrons to identify isotopes of an element;

Identify isotopes of elements giving only the numbers of protons and neutrons or the mass number and atomic number only. Use chlorine as an example.

Focus on the two numbers given on the Periodic Table. Point out the key. Use chlorine to illustrate the difference between the mass number and the relative atomic mass

1.1.12

· calculate the relative atomic mass of elements from the mass number and abundances of its isotopes; and

Allow students to perform average calculations from the heights of the students in the class, progress to weighted average calculations from teacher supplied height data from a tally table, eventually supply the data on chlorine and ask students to work out the weighted mean of the mass of chlorine. (Data on 100 atoms given in a tally table format may be the most meaningful way to supply data)

Comm – W

UM

Unit/Option content






1.1.13

· recall that a compound is two or more elements chemically combined.

Students could be asked to draw up mind map linking the ideas they have learnt about atomic structure, sub-atomic particles, electronic structure, atoms, ions and isotopes as a revision summary. This could be an independent activity or a group effort

SM

Resources

Atomic Structure Factfile

Periodic Table as supplied by CCEA in GCSE Data Leaflet

http://atomictimeline.net/index.php

www.rsc.org/learn-chemistry/resource/res00001332/the-atom-detectives

Blank A3 size table with the outline of the Periodic Table (elements 1-20) for students to put in electronic configuration of atoms

Scales and/or measuring tapes for average work


Unit 1

Unit/Option content




1.2 Bonding:

Ionic Bonding


1.2.1

· demonstrate knowledge and understanding that an ion is a charged particle formed when an atom gains or loses electrons and a molecular ion is a charged particle containing more than one atom;

Using the diagrammatic representations of the electronic structures of the first 20 elements in the Periodic Table, guide the students to realising the connection between the group number and the number of electrons in the outer shell (you may wish to illustrate another pattern at this stage, the connection between the number of shells in an atom and the period in which the element is placed). Ask the students to equate the lack of reactivity of the noble/inert gases with the pattern of electrons

Once it is established that stability is gained by securing a full outer shell, explore how this can be achieved by different atoms by losing or gaining the appropriate number of electrons, to attain the electronic structure of the nearest noble gas

Introduce the terms ‘ion’ and illustrate the appropriate chemical nomenclature: use of curly arrows to indicate movement of electrons and use of square brackets around diagrammatic representations of ions. Accompany the diagrams with electronic structures written in the format 2,8,8. When showing the diagrams of ions, use the opportunity to introduce the terms ‘anion’ and ‘cation’ along with their definitions

Further develop the process from the formation of ions into the formation of compounds. Explain that the ionic bond is the attraction between oppositely charged ions, and go through a series of examples with initially Group 1(I) elements bonding with Group 7(VII) elements, and Group 2(II) elements bonding with Group 6(VI) elements. Then progress to Group 1(I) elements bonding with Group 6(VI) elements, and Group 2(II) elements bonding with Group 7(VII) elements

Comm – T&L

1.2.2

· define the terms cation and anion; and explain, using dot and cross diagrams, how ions are formed and how ionic bonding takes place in simple ionic compounds, restricted to elements in Groups 1 (I) and 2 (II) with elements in Groups 6 (VI) and 7 (VII), the ions of which have a noble gas electronic structure;

Unit/Option content




1.2 Bonding:

Ionic Bonding

(cont.)


1.2.3

· demonstrate understanding that:

· ionic bonding involves attraction between oppositely charged ions;

· ionic bonds are strong; and

· substantial energy is required to break ionic bonds;

Discuss the strength of ionic bonds – this aspect will be covered in more detail in Structures

Comm – T&L

1.2.4

· recognise that ionic bonding is typical of metal compounds;

Using the examples given and their Periodic Tables, show that ionic bonding, involving oppositely charged ions, must have ions from opposite sides of the Periodic Table and hence tends to be the bond formed when metals bond with non-metals to form metal compounds

Unit/Option content




1.2 Bonding:

Covalent Bonding


1.2.5

· describe a single covalent bond as a shared pair of electrons;

Refer students to their drawings of the electronic structures of the first 20 elements. Ask them to explain why the elements in Group 8/0 do not generally form compounds. Ask them to reiterate the solution to this dilemma preferred by the elements in Group 1, 2, 6 or 7. Look at the structure of the atoms of these elements in Group 4. How will these atoms procure a full shell?

Explain the theory of sharing electrons to form a covalent bond.

It is helpful to examine other words with the stem ‘co’ Co-operative co-habit; to emphasise the meaning to share. (Valency the number of electrons in the outer shell of the atom)

Begin with HCl and Cl2. (H2 can be difficult because there is only one electron)

PS

1.2.6

· explain using dot and cross diagrams how covalent bonding occurs in H2, Cl2, HCl, H2O, NH3, CH4 and similar molecules and label lone pairs of electrons;

Illustrate the dot-cross method of drawing covalent compounds. Allow students to use the same method to draw other examples. (Encourage the use of different colours for the electrons from different elements. This is not a prerequisite for dot cross diagrams but does emphasise the point and helps students clarify what they are writing)

Focus on the position of the shared electrons. Students often place shared electrons on one outer shell and close to the outer shell of the second atom rather than on both outer shells

Practise drawing similar molecules; e.g. carbon tetrachloride (tetrachloromethane), nitrogen trifluoride

WO

SM

SM

Unit/Option content




1.2 Bonding:

Covalent Bonding

(cont.)


1.2.7

· draw dot and cross diagrams and indicate the presence of multiple bonds in O2, N2 and CO2;

Point out lone pairs of electrons where they occur in all of the diagrams used

Students progress to drawing the more difficult examples with multiple bonds. Emphasise that the arbitrary spherical nature of the outer shells is not a necessity

Comm – T&L

UM

1.2.8

· recognise covalent bonding as typical of non-metallic elements and compounds;

Draw the attention of the students to the formula of the covalent compounds. Ask students to classify the elements in their previous exercise as metals or non metals and predict a general rule for the formation of covalent compounds. Ask students to come up with two or more examples of covalent compounds

PS

1.2.9

· demonstrate knowledge and understanding that a molecule is two or more atoms covalently bonded and that diatomic means there are two atoms covalently bonded in a molecule;

Emphasise the difference between a drawing of the electronic structure of an atom and a diatomic molecule. (This should be flagged up in the student notes as it is a very common mistake made during examinations). Focusing on H2, F2 and Cl2, refer students to the Periodic Table and point out the diatomic elements. They can see why and how these elements are considered diatomic

Define ‘diatomic molecule’, ‘molecule’, ‘diatomic’

PS

Unit/Option content




1.2 Bonding:

Covalent Bonding

(cont.)


1.2.10

· demonstrate knowledge and understanding that covalent bonds are strong and substantial energy is required to break covalent bonds;

Discuss the strength of covalent bonds – this aspect will be covered in more detail in Structures

Comm – T&L

1.2.11

· demonstrate knowledge and understanding that a covalent bond may be represented by a line; and

Use the covalent compounds that have been drawn out to show how a line can represent a single covalent bone (i.e. a shared pair of electrons)

PS

Unit/Option content




1.2 Bonding:

Metallic Bonding


1.2.12

· demonstrate knowledge and understanding that metallic bonding results from the attraction between the positive ions in a regular lattice and the delocalised electrons.

Students have experience that metal atom and non-metal atom can combine by transferring electrons and a non-metal atom and a non-metal atom can combine by sharing electrons. Pose the question “Do atoms of metal elements combine with atoms of metal elements?”

Illustrate the bonding in metals focusing on

· Regular arrangement/layers;

· of positive ions; and

· in a sea of delocalised electrons.

Students should be familiar with the diagrammatic representation and the language used. Emphasise labelling of diagrams as a general rule

The metallic bond arises from the attraction of the positive ions to the delocalised electrons. Examination questions can ask about the structure of a metal or the bonding in a metal, emphasise the subtle difference in the answer expected by the examiner

Comm – T&L

Resources

Students own completed A3 size table with the outline of the Periodic Table (elements 1 – 20) from section on Atomic Structure. Factfile on Bonding. CCEA Periodic Table. PowerPoints, animations and slide shows are available on YouTube; teachers should choose those appropriate for their classes


Unit 1

Unit/Option content




1.3 Structures:

Ionic Structures


1.3.1

· use the accepted structural model for giant ionic lattices to explain the physical properties of ionic substances such as sodium chloride, including melting point, boiling point and electrical conductivity (diagram of giant ionic lattice is not expected);

Students demonstrate the properties of sodium chloride mentioned in 1.3.1. Include the solubility of sodium chloride in water. Although the students do not have to be able to explain the solubility in terms of structure, the specification expects the students to be familiar with the solubility of ionic compounds.

Student activities

- Electrolysis is not formally examined in the unit.

It is sufficient to demonstrate that sodium chloride conducts as an aqueous solution and state it conducts as a molten liquid and, in a similar fashion, that sugar, which has a simple molecular structure, does not. This is due to the ions which are free to move and carry the charge. Distinguish between the mechanism of conductivity of metals and ionic compounds;

- Explain the energy requirement to break bonds. Equate the energy requirement to melting point. Pose the question: “if the melting point of a substance is high, what does this indicate about the bonds holding the particles together?”; and

PS

UM

1.3.2

· recall that most ionic compounds are soluble in water;

- Although the diagram of a giant ionic lattice is not expected, the use of one, or a model, or a model made by students using marshmallows and toothpicks, can help the students to visualise the ionic crystal.

WO

Unit/Option content




1.3 Structures (cont.)


1.3.3

· use the accepted structural model for molecular covalent structures to explain the physical properties of molecular covalent structures such as iodine and carbon dioxide, including melting point, boiling point and electrical conductivity;

Students demonstrate the sublimation of iodine, using a fume cupboard. Remind the students about iodine’s diatomicity and ask “when the iodine vaporises, what types of particles are moving?”

Students should have sufficient information to attempt to explain why simple molecular substances do not conduct electricity: they have no free electrons; the electrons are shared; they have no ions

Using several single models of any covalent compound, methane, ammonia or water, illustrate to students exactly what breaks when the solid substance is melted. Emphasise the difference between the forces between the molecules and the bonds within the molecule between the atoms

Emphasise the strength of the covalent bond and illustrate this by looking at a model of a covalently bonded substance which has giant covalent structure, e.g. diamond

PS

1.3.4

· demonstrate knowledge and understanding that the intermolecular forces between covalent molecules are weak forces called van der Waals’ forces;

1.3.5

· recall that many covalent molecular substances are insoluble in water;

Unit/Option content




1.3 Structures (cont.): Giant covalent structures


1.3.6

· demonstrate knowledge and understanding of the giant covalent structure of carbon (diamond) and carbon (graphite), and predict and explain their physical properties, including:

· electrical conductivity;

· hardness;

· melting point and boiling point; and

· their uses in cutting tools (diamond), lubricants and pencils (graphite);

Students explain the structures of the various forms of carbon, concentrating on the hexagonal rings, the number of electrons on the outer shell used in bonding, the presences or absences of van der Waals forces and layers. As well as the properties listed in 1.3.6, graphite has numerous other properties. The students may be asked to find out one or two others, for example the use in extractor hoods to capture malodorous molecules or in developing countries to filter milk to make it safe to drink

Students may be interested to note some unusual uses of graphite as a marking agent: in the lake district it is traditional to tie lumps of graphite around a ram’s neck to mark a ewe to indicate mating. Graphite can be used to lubricate zips, which are stuck and placed in the grooves of a lampshade holder to prevent the plastic form from sticking, due to the prolonged effects of the heat of the light bulb. As with all uses, the students should have copies of photographs of the substance in situ

Emphasise that the covalent bonds in diamond are so hard that diamond can only be cut by another diamond, and is used in cutting tools

SM

Unit/Option content




1.3 Structures (cont.): Metallic Structures


1.3.7

· use the accepted structural model for metals to predict and explain their structure and physical properties, including melting point, malleability, ductility and electrical conductivity;

Students explain the properties malleability, ductility in terms of the positive ions being able to move past each other without being repelled, due to the sea of delocalised electrons, the layers of positive ions being able to slide. Explain the electrical conductivity in terms of the sea of delocalised electrons which can move when a potential difference is applied. State that the melting points of metals are generally less than ionic compounds, but greater than those covalent compounds with simple molecular structures. Explain this is due to the attraction of the negative electrons to the positive ions

SM

1.3.8

· demonstrate knowledge and understanding that an alloy is a mixture of two or more elements, at least one of which is a metal, and the resulting mixture has metallic properties;

Discuss the mixing of metals with other materials to form alloys. Emphasise the definition of an alloy, and demonstrate using those alloys which the student will have experience of, for example:

Steel – a mixture of iron and between 0.02 and 2% carbon

Bronze – copper and tin

Brass – copper and zinc

Alloys for use in the aircraft industry – aluminium and magnesium to produce light substances which have extra strength

The physical properties (melting point, boiling point, electrical and thermal conductivity) of the alloy does not significantly differ from the main metal in the alloy but the advantage of the alloy is usually in the engineering properties such as the tensile strength and the sheer strength of the alloy

Comm – T&L

Unit/Option content




1.3 Structures (cont.)


1.3.9

· demonstrate knowledge and understanding that the different sizes of atoms in an alloy distort the layers in the metallic structure, making it more difficult for them to slide over each other and so alloys are harder than pure metals;

As alloys contain atoms of different sizes, which distorts the regular arrangement (metallic structure), it more difficult for the layers to slide over each other, so alloys are harder than the pure metal

Examples include amalgam (mercury), steel (iron), brass (copper, zinc and other constituents), solder (lead and tin)

1.3.10

· recall that gold used as jewellery is usually an alloy with silver, copper and zinc and that the proportion of gold is measured in carats, with 24 carat being pure gold and 18 carat being 75% gold;

Pure gold is very soft and malleable. The process of alloying is used in gold jewellery (mixing with other metals) to change its properties including its hardness, strength and colour. Yellow gold is the most popular in the world and comes in a variety of shades depending on the other metal content. White gold is created by alloying with white metals such as palladium, silver or zinc. In addition it is usually plated with rhodium to create a harder surface with a brighter shine. Rose gold contains copper

The percentage of gold in an alloy is measured in carats.

100% pure gold is 24 carats, 75% gold is 18 carats

Why would gold alloy jewellery be desirable?

Which metals may be used to create different colours of gold jewellery?

Compare standard carats of jewellery in different countries

Comm – T&L

Unit/Option content




1.3 Structures (cont.): Structure and bonding of carbon


· demonstrate knowledge and understanding that carbon can form four covalent bonds;

Recall the structures of graphite and diamond

1.3.12

· demonstrate knowledge and understanding of the structure of graphene (a single atom thick layer of graphite), explain its physical properties, including strength and electrical conductivity, and recall its uses such as those in batteries and solar cells;

Graphene is a single atom thick layer of graphite. It is reportedly about 100 times stronger than the strongest steel. It conducts heat and electricity efficiently and is nearly transparent Graphene’s stability is a result of having its carbon atoms densely packed in a regular hexagonal pattern. Conduction of electricity is due to the availability of free-moving electrons. As graphene is a transparent and flexible conductor, applications may be diverse, including solar cells, light-emitting diodes (LED), touch panels, and smart windows or phones. Wearable device, and home appliance manufacturers may also find applications for this allotrope of carbon. Smart phone products with graphene touch screens are currently on the market. Graphene is getting cheaper; recently scientists at the University of Glasgow have produced graphene at a cost that is 100 times less than the previous methods. This topic lends itself well to a research homework about graphene

UICT

1.3.13

· demonstrate knowledge and understanding of the meaning of the term allotrope as applied to carbon (diamond), carbon (graphite) and graphene; and

This will lead to a discussion of the different forms of the element carbon. An allotrope is a substance which can exist in several physical forms in the same state. The molecules can be built together differently. Initially illustrate this simply with children’s building blocks. The story of the consequences of using tin buttons on the uniforms of Napoleon’s army will lend a historical aspect to the discussion

Comm – T&L

Unit/Option content




1.3 Structures (cont.): Classification of structures


1.3.14

· use given information, to classify the structure of substances as giant ionic lattice, molecular covalent, giant covalent or metallic.

To complete the study of structure, build up a table of comparative information about each different type of structure, using a template similar to that shown in the Resources section below. Students could prepare the summary table shown, in groups and present their work to the class, inviting questions from their peers to consolidate learning

Provide students with tabular data on the melting point, boiling point and electrical conductivity of different types of substances. Ask students to identify types of substances or actual examples, i.e. could be NaCl as it is ionic and ionic compounds have high melting points and only conduct electricity as a liquid or aqueous, or could be iodine as it has a simple molecular structure and is a solid at room temperature. Past paper questions can provide a source of these types of tables

Comm - W

WO

PS

Resources

‘Structures’ Fact File

Models of the structures Past paper questions can be used to show students the level of detail they are expected to know

Animations and PowerPoints are available on YouTube and on the Royal Society of Chemistry’s Learn Chemistry website

Gold Jewellery: World Gold Council website

Bonding

Metallic

Ionic

Covalent

Molecular

Giant

Nature of electrons

Example

Drawing of how bonding is established and/or structure

Melting point

Boiling point

Electrical conductivity


Unit 1

Unit/Option content




1.4 Nanoparticles


1.4.1

· demonstrate knowledge and understanding that nanoparticles are structures that are 1 – 100 nm in size and contain a few hundred atoms;

The teacher introduces the topic with a class discussion to explore prior knowledge and understanding

Nanoscience is the branch of science concerned with the development and production and uses of materials whose basic components are of nanoscale size, i.e. ~1 – 100 nm in size

Nanotechnologies describe the many ways that scientists can now work with the actual molecules and atoms that make up our world. It is a way of making things

We measure things in metres and centimetres, but nano scientists work in nanometres, that’s a billionth of a metre. That is very, very, VERY small!

Explain why small makes a difference

At this nanoscale, things don’t always behave as they do when they are larger. They might be stronger or lighter, or, more reactive or because they are so small, and they can be used in different ways than in their larger form

Scientists working at the nanoscale can make new products or processes, from mobile phones to sunscreens, airplanes to medicines.

The use of the scanning tunnelling microscope allows us to 'see' individual atoms in an atomic or molecular lattice in a way that was inconceivable 100 years ago when the principles of atomic and molecular structure were being discovered.

Comm – T&L

1.4 Nanoparticles (cont.)

So, it is now possible for students to see and investigate nanoscale structures at the atomic-molecular level and this 'feedback' enables you to compare the actual structure with the desired designed structure which eventually you would hope to have the prescribed desirable properties

With CAD (computer aided design) at the molecular level, there doesn't seem to any limit (within the laws and principles of chemistry) as to what structures we can build

Describe the many applications of nanotechnology including the use of semi-conductors that only conduct electricity in specific conditions and allow the design of much very tiny 'devices' normal scale conductors, so the final product can be much smaller, enabling the design and use of faster smaller computers working at the molecular level. It will be/is? possible to make very tiny mechanical devices to perform some task in otherwise inaccessible situations

Nanostructures are material structures assembled from layers or clusters of atoms of nanoscale size i.e. ~1–100 nanometre. By controlling the size and assembling of nanoscale constituents it is possible to alter and control the structure and properties of the final nanostructure. The advantage of these new materials is that they can be designed and built from the atomic level upwards to have specific properties of great use to material scientists, a good example is the ongoing development in the design and use carbon nanotubes. Students research and note samples.

PS

Comm – T&L

UICT

PS

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Nanostructures are built up from atomic or molecular precursors and processed via chemical deposition or physical vapour deposition, gas condensation, chemical precipitation, aerosol reactions and biological templating – a wide range of methods of assembling arrays of atoms.

Note that some nanoparticles are created naturally e.g.

· Very finely suspended mineral particles in water - the tiniest of colloidal particles act as nanoparticles.

· During inefficient combustion of organic molecules e.g. fossil fuels or plastics, nanosized particles of soot (mainly carbon) are formed.

· Evaporated seaspray can produce nanoscale salt particles.

Nanomaterial is a general word for any material that has a composition based on nanoparticle units e.g. nanoparticles of silver, carbon nanotubes, inorganic ceramic materials etc.

Nanoparticles are usually in the size range of 1 to 100 nm, described as being of nanoscale

Nanoparticles can be made of elements, organic molecules, inorganic compounds, inorganic cluster compounds or metallic/semi-conductor (maybe ~'semi-metal') particles

Nanoparticles have a high surface area to volume ratio which has a dramatic effect on their properties compared to non-nanoscale forms of the same material

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As a point of comparison, since nanoparticles are in the size range 1 – 100 nm, a human hair is 0.05 to 0.1 mm (50000 – 100000 nm) in diameter, in other words nanoparticles are usually 500 – 100000 times 'thinner' than a human hair!

1 nanometre, 1 nm = 10-9 of a metre (0.000 000 001 m)

Compared to other units:

1 cm = 10-2 m (1 cm = 10000000 nm)

1 mm = 10-3 m (1 millimetre = 1000000 nm)

1 μm = 10-6 m (1 micrometre = 1000 nm)

1 nm = 10-9 m

1 pm = 10-12 m (1 picometre = 0.001 nm)

1 fm = 10-15 m (1 fm = 0.001 pm)

Introduce terms used:

'nano' is a prefix and refers to dimensions-size of 1 – 100 nm. i.e. of nanoscale (1 x 10-9 m to 1 x 10-7 m)

nm is the accepted abbreviation for nanometre (nanometre) – more on the relative size of molecules and bigger particles is given further down in a size comparison data table to put the nanoscale 'scene' in perspective

Scale units for nanometres,

1 nm = 1 x 10-9 m = 1 x 10-6 mm,

so 1 nm is a millionth of a millimetre.

Nanoparticles may contain just a few hundred atoms

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Put nanoparticles in 'size' or 'dimension' perspective; consider the table below of 'materials' – pure elements, pure compounds and other more complex materials etc.

Students create a size comparison table of various 'particles'.

Example of data table of particle sizes

PS

1.4.2

· demonstrate knowledge and understanding of surface area to volume relationships and that as the side of a cube decreases by a factor of 10 the surface area to volume ratio increases by a factor of 10;

Typical nanoparticles are roughly spherical in shape, but the surface area to volume ratio is extremely important.

If you think of a simple cube, if the side is decreased by a factor of 10, the surface to volume ratio increases by 10

Analysis of surface area: volume ratio for selected nm cube sizes for some simple calculations to emphasise this point but for comparison units must be the same

PS

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1.4.3

· demonstrate knowledge and understanding that nanoparticles have properties different from those for the same material in bulk due to their high surface area to volume ratio; and

Explain that nanoparticles behave differently and give rise to properties.

The nano effect...

Scientists are interested in the nanoscale because when we get down to these tiny sizes, many materials start to behave in different ways. They are sometimes much stronger, or conduct more electricity, opaque substances can become transparent, solids become liquids at room temperature or insulators become conductors. This is often down to the change in their surface area when they are used at this tiny scale

Ask the students – What is Surface Area?

The Surface Area is what makes nano really interesting – that is the surface area to volume ratio – and it gets talked about a lot in nanotechnology.

The surface area of an object is the amount of surface it has and the volume is a measure of how much space it takes up. When we talk about the ratio between these two things, we are comparing how much of each quantity it has.

Ask the students – Why is that interesting in a nanoparticle?

The amount of surface area the particle has is larger compared to its volume. This means there are more atoms on the surface of the particle than in the middle of it, and that makes them the most important. Surface atoms act differently to atoms inside a particle, so when there are more surface atoms than inside atoms the way they behave dominates the whole behaviour of the particle

Comm – T&L

Comm – T&L

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The opposite is also true, when the particle is bigger it has a large volume compared to its surface area and the number of atoms inside the particle is much higher than the number of atoms on the outside (the surface) of the particle. What the inside atoms are doing is the most important thing and the behaviour of the particle will be decided by them

Explain what difference it makes:

How surface atoms and inside atoms behave can be very different. This means that when we get a very small piece of material, with comparatively large numbers of surface atoms, the material can act very differently to what we are used to (aluminium nanoparticles explode!). In nanotechnology we are making use of particles with lots of surface atoms and the fact that this makes them behave differently

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1.4.4


· evaluate the benefits of nanoparticles in sun creams, including better coverage and more effective protection from the Sun’s ultraviolet rays, and the risks, such as potential cell damage in the body and harmful effects on the environment.

Example introduction video: A Little Bit of Sunshine, Mr. O

Uses of nanoparticles of titanium(IV) oxide (titanium dioxide, TiO2)

How does titanium dioxide protect us from uv light? What does titanium dioxide do in sun creams? Students can research this question on-line.

In the cosmetics industry the use of nanosized particles in creams etc. is increasing, because of the small particle size, they can be more easily absorbed through the skin, as in moisturisers

Titanium dioxide, TiO2, is a white powder and a good reflector of visible light and most commonly encountered as a brilliant white pigment in paint

However, it also used in the cosmetic and skincare products industry as a pigment, thickener, moisturiser and used in sunscreens as a uv absorber. A research homework could be set requiring students to find out about the role of nanoparticles in uv protection.

Ultraviolet light from the sun impacting on the skin can cause cell damage, e.g. DNA damage, which can lead to skin cancer.

Nanoparticles used in sun creams offer better skin protection than traditional uv sun blocker creams and less is needed to cover the same surface area of skin.

In sun creams, nanoparticles provide better protection from the harmful effects of ultraviolet radiation (uv rays from the sun) and give better coverage of the skin

Titanium dioxide is a good uv light absorber and is effective as many organic molecules used as uv absorbers in many commercial sun creams

UICT

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Some of these organic 'sun blocker' molecules can cause skin irritations on sensitive skin

Can the nanoparticles be absorbed by the skin and cause this irritation or other effects?

Nanoparticles of titanium dioxide are incorporated into 'sun blockers', though the particles have to be specially coated to avoid skin irritation problems from using such titanium dioxide based sunscreens

The interaction of uv light and titanium dioxide in the presence of other molecules can produce highly reactive and harmful free radicals

As well as being a good uv absorber (90% absorbed, 10% scattered), nanosized titanium dioxide particles have another advantage

The TiO2 particles are smaller than the wavelength of light and are therefore too small to see, and when used in sunscreen creams they are transparent to light rather than opaque, so the skin looks a more natural colour (no white creamy marks)

A good example of where particle size matter, and the tiny size of titanium dioxide particles gives it this commercial advantage

Note in passing that nanoparticles of zinc oxide are also used in sunscreens with the similar properties and effects

Students work in small groups, research the topic using ICT, delegate sections of a report to each member, prepare a poster and debate the use of nanoparticles in sun cream for/against including preparing counter arguments

Comm – W, T&L


Unit 1

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1.5 Symbols, formulae and equations


1.5.1

· recognise symbols and names for common elements and recall the diatomic elements;

Using the students’ Periodic Tables, recap the symbols and names for common elements including diatomic elements

Chemical symbols bingo. Prepare a 4 x 4 grid with the symbols of elements. Call out the name, students mark off until they get 4 in a row

Grid with list of symbols. Students write the element name

Pictures of elements. Students write the element name and symbol

Take the students through some common examples such as H2O, NaCl, O2, etc. Explain that the subscript number indicates the number of atoms of that element in the molecule. Note that writing ‘1’ is not required

PS

1.5.2

· interpret chemical formulae by naming the elements and stating the number of each type of atom present;

Show that some formulas have brackets in them. For example, sodium hydroxide is NaOH and magnesium hydroxide is Mg (OH)2. There are two of each atom inside the bracket. Mg(OH)2 contains one magnesium atom, two oxygen atoms and two hydrogen atoms

1.5.3

· write chemical formulae of compounds;

Prepare a worksheet of examples for the students to complete

Students work in pairs to prepare a worksheet for their partner to complete. Set the task with a range of examples of covalent and ionic compounds

Comm – W, UM

WO

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1.5 Symbols, formulae and equations (cont.)


1.5.4

· demonstrate understanding that chemical reactions use up reactants and produce new substances called products;

Explain that in a chemical reaction, the reactants are the substances present when the reaction begins, and the products are the new substances produced as a result of the reaction. In a chemical formula, the reactants are on the left side of the arrow, and the products are on the right. There can be clues to show that a chemical reaction has taken place, for example: can see, smell or feel a change – a new substance is visible, a gas is formed and heat is given off or taken in

PS

1.5.5

· construct word equations to describe the range of reactions covered in this specification;

Demonstrate word equations for the range of reactions in this specification

Use the concept of conservation of mass to explain that all atoms involved in a reaction must be accounted for

E.g. copper + oxygen → copper(II) oxide; Cu + O2 → CuO

Balanced Equation 2Cu + O2 → 2CuO

1.5.6

· recognise that in a chemical reaction no atoms are lost or made but they are rearranged, and as a result we can write balanced symbol equations showing the atoms involved;

Demonstrate how to balance an equation

Students write symbol equations for the word equations in this specification

UM

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1.5.7

· write balanced symbol equations for all reactions covered in this specification and for unfamiliar chemical reactions when the names of the reactants and products are specified;

Prepare word equations of some unfamiliar chemical reactions from A level specification for the students to write balanced symbol equations

PS

1.5.8

· write balanced ionic equations for reactions, to include reactions covered in this specification;

For example: the displacement reactions can be written as ionic equations

Using iron and copper(II) sulfate:

iron + copper(II) sulfate → iron(II) sulfate + copper.

Fe(s) + CuSO4(aq) → FeSO4(aq) + Cu(s)

Copper(II) sulfate and iron sulfate are ionic compounds

When they are dissolved in water the ions become separated by the water molecules

Write the equation showing the ions:

Fe(s) + Cu2+(aq)+ (aq) → Fe2+(aq) + (aq) + Cu(s)

UM

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Going from reactants to products:

iron metal – Fe(s) has become iron ions - Fe2+(aq)

copper(II) ions – Cu2+(aq) have become copper metal - Cu(s)

sulfate ions – (aq) are not changed during the reaction

Sulfate ions are the same on the left and the right side of the arrow. Ions which do not change during the reaction are called spectator ions

Spectator ions can be left out of the equation, giving

Fe(s) + Cu2+(aq) → Fe2+(aq) + Cu(s)

This is the ionic equation for the reaction between iron and copper(II) sulfate. Iron is oxidised and copper is reduced

Links to displacement reactions of metals (KS3) which can be recalled as an introduction to section 2.1

Links to section 1.7 Quantitative Chemistry

UM

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1.5.9.

· write half equations for reactions covered in this specification; and

Links to section 1.6.13 – students write half equations for the formation of a Group 1 (I) ion from its atom

Links to section 1.6.22 – students write half equations for the formation of a halide ion from a halogen molecule or atom

For example, Cl2 + 2e- → 2Cl-.

Ensure half-equations are balanced and explain they are written as reductions by convention

UM

1.5.10

· demonstrate knowledge and understanding that in chemical equations the three states of matter are shown as (s), (l) and (g) with (aq) for aqueous solutions, and include appropriate state symbols in balanced symbol equations for the reactions in this specification.

For example, the precipitation reaction, shown as a word equation and as a balanced symbol equation:

copper(II) sulfate + sodium hydroxide → sodium sulfate + copper(II) hydroxide

CuSO4(aq) + 2NaOH(aq) → Na2SO4(aq) + Cu(OH)2(s)

The copper(II) hydroxide forms a solid (the precipitate) because its state symbol is (s) for solid, rather than (aq) for aqueous (dissolved in water).

The reaction can also be shown by an ionic equation:

Cu2+(aq) + 2OH–(aq) → Cu(OH)2(s)

This only shows the reaction between the ions that produce the precipitate

Resources

Fact file on Symbols, formula and equations.

PowerPoints, animations and slide shows are available on YouTube; teachers should choose those appropriate for their classes


Unit 1 1.6 The Periodic Table

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1.6 The Periodic Table: Basic structure of the Periodic Table


1.6.1

· describe how Mendeleev arranged the elements in the Periodic Table and left gaps for elements that had not been discovered at that time, and how this enabled him to predict properties of undiscovered elements;

Discuss the Periodic Table with students, in terms of the elements, and use the key in the bottom left corner to show how much information is collated here

Who collated all this information into such a handy table? Was it the work of one person or the culmination of the work of many?

Students can do independent research or watch one of the videos available to find out about the development of the Periodic Table

Comm – T&L

PS

1.6.2

· demonstrate knowledge and understanding of how scientific ideas have changed over time in terms of the differences and similarities between Mendeleev’s Periodic Table and the modern Periodic Table;

Card sorting exercises focusing on the features of Mendeleev’s periodic table and the modern periodic table can be carried out to reinforce the differences

Comm – T&L, PS

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1.6 The Periodic Table (cont.)


1.6.3

· describe an element as a substance that consists of only one type of atom and demonstrate understanding that elements cannot be broken down into simpler substances by chemical means;

Look at the Periodic Table as a list of the elements currently known. Ask students to point out elements which they recognise. Show samples of some of the elements they are likely to be familiar with. Emphasise the definition of the term element. Focus on the fact that these substances have only one type of atom and cannot be broken down into simpler substances by chemical means

Comm – T&L

1.6.4

· demonstrate knowledge and understanding that a group is a vertical column in the Periodic Table and a period is a horizontal row;

Refer back to work on electronic structure of atoms to reinforce the terminology of groups and periods, and reinforce the link between position in the periodic table and electronic structure

1.6.5

· identify and recall the position of metals and non-metals in the Periodic Table and distinguish

between them according to their properties, including conduction of heat and electricity, ductility, malleability, melting point and sonority;

Students can code their copies of the Periodic Table using coloured pens

Metal/non-metal division

Gases

Liquids

Diatomic elements/molecules

Names of groups 1, 2, 7 and 8/0 and transition elements

Properties of metals and non-metals, as discussed during the teaching and learning of structures can be reinforced and shown in another context

PS

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1.6.6

· identify elements as solids, liquids and gases (at room temperature and pressure) in the Periodic Table;

Recap of use of state symbol when writing chemical formulae from section 1.5.10

1.6.7

· demonstrate knowledge and understanding that elements in the same group in the Periodic Table have the same number of electrons in their outer shell and this gives them similar chemical properties; and

Having established that the number of electrons in the outer shell equates to group number, and by extension, elements in the same group will have the same charge on their ions for atoms of elements in Groups 1, 2, 6 and 7, and that the elements in Group 8/0 are chemically inert, it is possible to move smoothly into the concept of elements in the same group having similar reactivity

1.6.8

· recall that elements with similar properties appear in the same group, for example Group 1 (I) and Group 2 (II) are groups of reactive metals, Group 7 (VII) is a group of reactive non-metals and Group 0 is a group of non-reactive non-metals, and locate these groups in the Periodic Table; and recall the names of the groups

Students may already have some experience of the elements and the reactions to be discussed here. Show students samples of lithium, sodium and potassium stored in oil, samples of solid iodine and liquid bromine, coloured photographs of oxygen cylinders/oxygen in gas jars, chlorine in a gas jar

Comm – T&L

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1.6 The Periodic Table (cont.): Group 1 (I)


1.6.9

· demonstrate knowledge that the alkali metals have low density and the first three are less dense than water;

Demonstrate the reactions of lithium, sodium and potassium in water. Emphasise the density of the metals, their storage, and give students the opportunity to view freshly cut samples as they tarnish. Discuss with students the risk assessment of the demonstration so that they understand the risks associated with the use of alkali metals

PS

Comm – T&L

1.6.10

· assess and manage risks associated with storage and use of alkali metals and recall that alkali metals are easily cut, shiny when freshly cut and tarnish rapidly in air;

1.6.11

· demonstrate knowledge and understanding that Group 1 (I) metals react with water to produce hydrogen and a metal hydroxide, and give observations for the reactions;

Students should be encouraged to record their own, independent observations of these reactions, before drawing the class observations into an agreed table of data

WO, SM

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1.6.12

· demonstrate knowledge and understanding that alkali metals have similar chemical properties because when they react an atom loses an electron to form a positive ion with a stable electronic configuration;

Students should be able to follow a sophisticated analysis of the reaction, stating reactants and products, writing chemical formula for reactants and products and attempting to write a balanced chemical equation for the reaction of these alkali metals in water. Focus on the pattern of reactivity and the fact that these elements are classified together due to their similar reactions

SM

1.6.13

· write half equations for the formation of a Group 1 (I) ion from its atom;

In writing the equations for the reactions which take place when the alkali metals react with water, students should notice that the metal changes from an atom to an ion, by losing an electron. Reinforcing the work in ‘Bonding’ we can talk about the relative stability of the ion, and the ionic equation showing the formation of the ion can be introduced

UM

1.6.14

· demonstrate knowledge and understanding of how the trend in reactivity down the group depends on the outer shell of electrons of the atoms;

Discuss the pattern of reactivity down the group, which has been elucidated by the demonstration practical, and tie to the electronic structure

Identify and analyse relationships and patterns: Interpret data to identify and state trends such as the reactivity of the Group 1 metals

PS

Comm – T&L

1.6.15

· demonstrate knowledge and understanding that most Group 1 (I) compounds are white and dissolve in water to give colourless solutions;

Some common salts e.g. NaCl, KCl, can be shown to the class and dissolved in water. Students may want to take photos of these to try and remember them. It may be constructive at this point to emphasise the difference between ‘colourless’ and ‘clear’, perhaps using copper(II) sulfate solution as a clear, but coloured, solution

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1.6 The Periodic Table (cont.): Group 7 (VII)


1.6.16

· recall data about the colour, physical state at room temperature and pressure, diatomicity and toxicity of the elements in Group 7 (VII), interpret given data to establish trends within the group and make predictions based on these trends;

Discuss the pattern of physical characteristics in Group 7, using video clips or PowerPoint slides to emphasise the trend

Ask students to predict properties for other elements based on what they know

Students should prepare a PowerPoint presentation on the properties of Group 7 elements, making comparisons and highlighting trends within the Group

UICT, Comm – T&L

UICT

1.6.17

· recall the observations when solid iodine sublimes on heating and understand the term sublimation;

Demonstrate sublimation of iodine, and discuss how this links to what we know about iodine’s structure (links to ‘structures’). Have students write down their own observations, draw up a class list and

re-demonstrate to reinforce the correct terminology

PS

1.6.18

· describe how to test for chlorine gas (damp universal indicator changes to red and then bleaches white);

Demo production of chlorine gas and show the indicator paper shows acidity before being bleached white

In groups students should discuss and explain trends in reactivity for Group 1 and Group 7 elements in terms of electronic structure.

WO

Comm – T&L

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1.6.19

· investigate the displacement reactions of Group 7 (VII) elements with solutions of other halides to establish the trend in reactivity within the group and make predictions based on this trend;

Use the video clips from the video presented by Leeds University Chemistry Department to show the displacement reactions of the Group 7 elements. Students have experience of the displacement reactions of the metals, a reminder of this may serve to clarify these displacements

These reactions will show the trend in reactivity down the group, and so predictions can be made – ask students to make predictions

PS

1.6.20

· demonstrate knowledge and understanding of how the reactivity down the group depends on the outer shell electrons of the atoms;

By looking at the electronic structures of the elements we can see the influence of electronic structure on reactivity

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1.6.21

· demonstrate knowledge that the halogens have similar chemical properties because when they react an atom gains an electron to form a negative ion with a stable electronic configuration;

Students will notice that the halogen atom changes to a halide ion by gaining an electron. Reinforcing the work in ‘Bonding’ we can talk about the relative stability of the ion, and the ionic equation showing the formation of the ion can be introduced. This should then be extended to equations showing the formation of halide ions from halogen molecules, since halogens are found as diatomic molecules

UM

1.6.22

· write half equations for the formation of a halide ion from a halogen molecule or atom;

1.6.23

· use the concept of electronic configuration to explain the lack of reactivity and the stability of the noble gases;

Using the electronic structure of the noble gases explain that the lack of reactivity is due to the stable nature of a full outer shell

1.6.24

· recall that the noble gases are colourless gases;

Use images of the noble gases, to emphasise this point

1.6.25

· demonstrate knowledge and understanding of the trend in boiling points of the noble gases going down the group;

Students use data on boiling points to demonstrate the trend and link the trend to the increased molecular mass, and increased van der Waals’ forces

UM

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1.6 The Periodic Table (cont.): Transition Metals


1.6.26

· compare the physical properties of the transition metals with Group 1 (I) elements to include melting point, and density and understand that the transition metals are much less reactive with water; and

Explain that the physical properties of transition metals are, they:

· form coloured compounds

· are good conductors of heat and electricity

· can be hammered or bent into shape easily

· have high melting points (but mercury is a liquid at room temperature)

· are usually hard and tough

· have high densities

· are much less reactive with water than Group 1 (I) elements

1.6.27

· demonstrate knowledge that transition elements form ions with different charges (for example iron(II) and iron(III)) and form coloured compounds:

· copper(II) oxide is black;

· copper(II) carbonate is green;

· hydrated copper(II) sulfate is blue; and

· copper salts are usually blue in solution.

Refer back to the colour-coded copy of the Periodic Table and focus this time on the transition metals. Use the information on the GCSE data leaflet to show that the same element can form different ions, namely iron(II) and iron(III), and explain that this phenomenon is common among the transition elements (no need to explain why at this level)

Use samples of transition element compounds to demonstrate the range of colours which exist in the transition metals, and list the coloured compounds they are expected to know. You can give these colours or have the students find them out as a research homework

PS

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Resources

Fact files

Samples of the alkali metals, trough, safety screen, knife, tongs, tissue paper, goggles/face screen, for demo of alkali metals with water.

Samples of metals and compounds available in the school chemistry department, images of elements and compounds from internet, PowerPoints and videos available within school or from YouTube

Fume cupboard for sublimation of iodine, preparation of chlorine


Unit 1

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1.7 Quantitative Chemistry: Formula Mass


1.7.1

· recall that the relative atomic mass (Ar) of an atom is the mass of the atom compared with that of the carbon-12 isotope, which has a mass of exactly 12, and demonstrate knowledge and understanding that it is a weighted mean of the mass numbers (linked to 1.1.12);

In the context of the history of the construction of the Periodic Table discuss the only numerical information Mendeleev had. It can be useful, although not required by the specification to inform the students about the original comparison to hydrogen. Remind students of the definition of isotopes and give some examples 14C, 12C, 16O, 18O

Students have already calculated the weighted mean of chlorine isotopes in section 1.1.12. The teaching strategy may be repeated here

Comm – T&L

UM

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1.7 Quantitative Chemistry: Formula Mass

(cont.)


1.7.2

· calculate relative formula mass (Mr) (relative molecular mass) of a compound and percentage of an element, by mass, in a compound

Point out where relative formula mass data can be found on the Periodic Table and explain the key on the bottom left of the centre pages of the data sheet supplied by CCEA. Be specific about the difference between the relative formula mass of a naturally occurring sample of an element and the mass number of an atom of that element

Progress through examples beginning with the formula of simple compounds through to the formula of compounds containing multiple molecular ions, for example from formula of the type NaCl through formula of the type Na2O, CaCl2, Al2O3, to Al2(SO4)3. Students have not yet been introduced to formula of compounds which contain water of crystallisation. Calculating the formula masses of compounds with formula of the type CuSO4.5H2O is covered in section 1.7.9

Students write formula for some compounds before calculating relative formula masses

UM, PS

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1.7 Quantitative Chemistry:

The Mole


1.7.3

· demonstrate knowledge and understanding that chemical amounts are measured in moles and that the mass of one mole of a substance in grams is numerically equal to the relative formula mass;

The concept of the mole is one which many students find difficult although Avogadro’s Number is not required for the purposes of the specification at this point it can be useful to equate the word mole to a number and to compare it to other number words, pair, dozen, score, gross, ream. Progress to suggesting how larger numbers are “counted” by weighing; weighing coins in a bank, weighing paper

Or

Begin the discussion using a balanced chemical equation for the production of sodium chloride from sodium and chlorine. Consider the implications if the equation meant single atoms. Consider a hypothetical salt factory who only wishes to produce the amount of salt their customer orders as excess production means loss of revenue. How is the production manager to know exactly how many atoms/molecules to react? Progress to ‘counting’ by weighing. Give other examples of the use of this technique, weighing coins, weighing paper

Comm – T&L

UM

Comm – T&L

PS

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The Mole (cont.)


1.7.4

· convert the given mass of a substance to the amount of the substance in moles (and vice versa) by using the relative atomic or formula masses;

Allow students to practice using the mathematical expression

Number of moles =

Where Mr = relative formula mass; Note RFM and RAM may be used where appropriate

Graded examples of data from which number of moles, mass or Mr can be calculated

Include examples of finding mass given number of moles and name or formula of substance and Mr given number of moles and mass. Include the diatomic molecules in the examples

UM

1.7.5

· demonstrate knowledge and understanding of the importance of scale in chemistry in terms of calculating moles from masses given in tonnes and kilograms, for example in industrial processes;

Include examples of masses given in tonnes and kg and emphasise the formulae above use mass in grams

Recap that no atoms are lost or made in a chemical reaction and hence it is possible to interpret equations quantitatively

Links to 1.5 Symbols, formulae and equations

After an explanation of the method of calculating the reacting masses of reactants or products, allow students to progress through examples where they are guided through each step to examples where they are given no guidance. Students are always supplied with a balanced chemical equation for the purposes of this exercise. Some students may find a set way of progressing through these calculations useful. Some may find a tabular way of setting out the information useful

What mass of calcium chloride will be formed from the complete reaction of 11.2 g of calcium oxide with excess hydrochloric acid?

UM

SM

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The Mole (cont.)


Students complete the table below;

1.7.6

· calculate the reacting masses of reactants or products, given a balanced symbol equation and using moles and simple ratio, including examples where there is a limiting reactant;

Copy the balanced symbol equation

CaO +

2HCl

CaCl2

H2O

Write the molar ratio

1

2

1

1

Fill in information given in the question

11.2g

Change the information to moles

Moles = 11.2/56

= 0.2 moles

Use the molar ratio to write down the number of moles of all of the reactant and products

0.2

0.4

0.2

0.2

Use the number of moles of the appropriate substance to answer the question

Mass=moles x Mr

= 0.2 x 111

= 22.2g

UM, PS

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1.7 Quantitative Chemistry: Percentage Yield


1.7.7

· calculate the theoretical yield, and hence the percentage yield, of a chemical reaction given the actual yield;

Show students a set of photographs of a student preparing a pure dry sample of a salt, copper(II) sulfate from copper(II) oxide and sulfuric acid is a suitable example. Viewing the photographs of the different stages ask why it is extremely rare to obtain the amount or mass of product calculated from the equation

Comm – T&L

1.7.8

· recognise possible reasons why the percentage yield of a product is less than 100%, including loss of product in separation from the reaction mixture, as a result of side reactions or because the reaction is reversible and may not go to completion;

Allow students to calculate the percentage yield of copper(II) sulfate from data given about this reaction and to calculate percentage yield of products from other examples

UM, PS

Unit/Option content





Calculation of the formula of compounds


1.7.9

· demonstrate knowledge and understanding of the terms empirical formula, molecular formula, hydrated, anhydrous and water of crystallisation;

Emphasise the heating and cooling and reweighing during the thermal decomposition and the necessity of this part of the process in determining the mass of water of crystallisation. Present examples to the students where they are given the mass of the hydrated compound and the mass of the dehydrated compound, progress through examples of the type where masses of the hydrated compound and the weighing boat are given to examples where several masses of the compound as it undergoes thermal decomposition are given

1.7.10

· demonstrate knowledge and understanding that water of crystallisation can be removed by heating to constant mass and any thermal decomposition may be carried out to completion by heating to constant mass;

Students carry out experiment finding the formula of hydrated copper(II) sulfate or hydrated magnesium sulfate

SM

WO

Comm – T&L, R, W

1.7.11

· calculate the relative formula mass of compounds containing water of crystallisation;

Unit/Option content





Calculation of the formula of compounds

(cont.)


1.7.12

· calculate the percentage of water of crystallisation in a compound;

1.7.13

· determine the empirical formulae of simple compounds and determine the moles of water of crystallisation present in a hydrated salt from percentage composition, mass composition or experimental data; and

· determine the mass of water present in hydrated crystals (Prescribed Practical C1).

Students complete graded examples

PS

Resources

Fact file on Quantitative chemistry

PowerPoints, animations and slide shows are available on YouTube; teachers should choose those appropriate for their classes

Graded examples of data from which reacting masses of products and reactants may be calculated

Graded examples of data from which theoretical and percentage yield may be calculated

A series of photographs or video showing the different stages of preparation of copper(II) sulfate crystals from copper(II) oxid

Documents

ccea.org.uk · Web view1.2.12 demonstrate knowledge and understanding that metallic bonding results from the attraction between the positive ions in a regular lattice and the delocalised