Atoms, Bond and Groups SOW 2012Teacher 1

(14 weeks) = 25 lessons

Teacher 2(14 weeks)

= 32 lessons

Relevant assessments(6 hours)

1.2.1 Electrons (7 hours)

1.1.1 Atoms and Reactions1.1.2 Moles and Equations

Quantitative 3 Determining the concentration of a solution of sulphuric acid.

Evaluative 2Identifying a Group 2 carbonate

1.2.2 Bonding and Structure(11 hours)TEST

1.1.3 Acids(18 hours)TEST

Investigating the reactions of some unknown chemicals (qualitative 2)

1.3.1 Periodicity (3 hours)

1.1.4 Redox(4 hours)

1.3.2 Group 2( 3 hours)

1.3.3 Group 7(4 hours)

Qualitative 1 Investigation of solutions and solids

Teacher 1 SOW

Lesson Syllabus content Activities1. Revision of atomic structureRevision of GCSE Electronic Structure

Describe protons, neutrons and electrons in terms of relative charge and relative massDescribe the distribution of mass and charge within an atomDescribe the meaning of atomic (proton) number and mass (nucleon) numberDeduce the numbers of protons, neutrons and electrons in atoms and ions.

Determine protons, neutrons, and electrons in atoms and ions (do via tables)Students draw the electronic arrangements of the first 20 elements in terms of 2.8.1 e.t.c.Relate to period and group number

2. Development of different theories of the structure of the atom.

Modern development of the structure of the atom; the changing accepted view of the structure of the atom; acceptance and rejection of different theories for the structure of the atom from the Greeks, Dalton, Rutherford, Moseley et al up to the Bohr atom.

Show ‘the atom’ DVDs relevant sections – episode 1. Pupils make notes and write an essay for homework.Chemistry in Context has a good section on this.‘quantum theory’ cartoon book is a useful resource.ILPAC 1 p.102 has some notes

3. Evidence of shells

State the number of electrons that can fill the first 4 shells

Pupils can perform flame tests with sodium chloride, potassium chloride, copper chloride, barium chloride and look through spectroscopes to deduce line spectra. From this discuss ‘quantisation of energy’ and the evidence for electron shells i.e. spectra are not continuous. ILPAC 1 p.86 exp.6Use ILPAC resources to derive nth shell gives 2n2 electronsShow relevant parts of Atom DVDAgain ‘quantum theory’ book is useful

4. Orbitals 1 Describe an orbital as a region that can hold up to 2 electrons with opposite spinsDescribe the shapes of s and p orbitalsState the number of orbitals making up s-, p-, and d-sub shells

Mention wave/particle duality. Mathematical equations describing electrons as waves in probability functions have been used to calculate shapes of orbitals.ILPAC resources. Number of electrons in an orbital = 2 (s),6 (p) ,10 (d) ,14 (f)

Describe the relative energies of s-,p-,d-, orbitals for the shells 1,2, 3 and the 4s and 4p orbitals

Again quantum theory book is useful.Shows how spectral lines were split further to give evidence of subshells. Pauli Exclusion principle.Derive energy diagram to show splitting of subshells and explain that the 3d orbital is lower in energy than the 4s

5. Orbitals 2 Deduce the electron arrangements of:(i) atoms given the atomic number, up to Z

=36(ii) ions given the atomic number and ionic

charge, limited to the s and p blocks up to Z =36

Classify the elements into s, p and d blocks

Write electronic configurations like 1s2 2s2 …… e.t.c. for first 36 elementsWrite electrons in boxes for the first 36 elements.Introduce Hund's rule (mutual electron/electron repulsion) and relate to ‘people getting on a bus’Last orbital being filled relates to the block it’s found in in the periodic table.Use the periodic table to count across in blocks to determine electronic configurations (alternative method)

6. Ionisation Energies

Define the terms first ionisation energy successive ionisation energy.Explain that ionisation energies are influenced by nuclear charge, electron shielding and the distance of the outermost electron from the nucleus.Predict from successive ionisation energies of an element:The number of electrons in each shell of an atom.The group of an element.

Demo Group 1 metals with water and relate enthalpy of reaction with the energy needed to ionise the metal.Plot graphs of successive ionisation energies of different elements. Why do the values increase in general? Use data book to get data.ILPAC 1 p.83Give successive ionisation energy data and ask the pupils to predict which group in the periodic table elements belong to.

7. Consolidation

Review the syllabus for atomic structure Give pupils a brief test and/orGet them to present a famous discovery about the atom (use quantum mechanics cartoon book as an aid)

8. Ionic Bonding

Describe the term ionic bonding as the electrostatic attraction between oppositely charged ionsConstruct dot cross diagrams to describe ionic bonding

Demo a few ionic reactionse.g. magnesium + oxygen zinc + sulphure.t.c.

Define an ionic bond ‘the electrostatic attraction between two oppositely charged ions’Practice drawing ionic bonding diagrams for outer shell electrons only i.e. don’t need to draw inner shells fro metals.

9. Ionic Formulae

Predict ionic charge from the position of an element in the periodic tableState the formulae for the following ions:Nitrate, carbonate, sulphate and ammonium

Show pupils all the ions formed in the periodic table. Emphasise that Group 4 and group 8 don’t form ions.Use CFY resource to get the pupils to draw different ionic formulae. Show them the 2 methods of deriving formulae. Also test them by recall on the formulae and charge of nitrate, carbonate, sulphate and ammonium ions.

10. Covalent Bonding

Describe the term covalent bond as a shared pair of electrons.Construct dot/cross diagrams to describe:

(i) single covalent bonding e.g. as in H2, Cl2, HCl, H2O, NH3, CH4, BF3, SF6

(ii) multiple covalent bonding e.g. as in O2, N2 and CO2

molecules analogous to those specified above.

Quick test on ionic formulaeDefine the term covalent bond and explain that it’s the mutual attraction of the two nuclei on the bonding pair of electrons that holds the two atoms together.Using outer shell electrons only (i.e. group numbers) draw the relevant examples. In SF6 the octet rule is expanded because sulphur has empty d orbitals, which it can use.Get pupils to draw analogous examples e.g. HBr, H2S, PH3, SiCl4, AlCl3, SF6

Then go into double bonded molecules. Introduce the term ‘isoelectronic’ as two molecules containing the same number of electrons e.g. CO is isoelectronic with N2

Set as many as you can for H/W

11. Dative Covalent Bonding

Dative covalent bonding as in NH4+

Analogous examplesDefine a dative covalent bond as a bond in which an electron rich atom donates a lone pair of electrons to from a bond with an electron deficient atom/ionGo through the ammonium example. Demo conc. ammonia reacting to form a white smoke next to a bottle of conc. HClAnalogous examples include NH3BF3 and H3O+, Al2Cl6 and

BeCl2

12/13. Shapes of Molecules

Explain that the shape of a simple molecule is determined by the repulsion between electron pairs surrounding the central atomState that lone pairs of electrons repel more than bonded pairsExplain the shapes of and bond angles in molecules and ions with up to six electron pairs (including lone pairs) surrounding a central atom, e.g. as in:

BF3 (trigonal planar)CH4 and NH4

+ (tetrahedral)SF6 (octahedral)NH3 (pyramid)H2O (non-linear)CO2 (linear)Predict the shapes of analogous molecules and ions.

Explain why a lone pair repels more than a bonding pair (it’s held more closely to the central atom)Go through examples using molymod kits. Students can make different shapes and try to calculate bond angles.Then go through analogous examples e.g.

AlCl3, SiCl4, NF3, PH3

H2S, CS2, H3O+, CH3+, NH2

-, SO2, SO3

14.Electronegativity

Describe the term electronegativity as the ability of an atom to attract the bonding electrons in a covalent bondExplain that a permanent dipole may arise when covalently bonded atoms have different electronegativities, resulting in a polar bond.

Demonstrate ILPAC experiment ‘Testing Liquids of Polarity (structure and bonding ILPAC book 3 p.46)Define electronegativity. Why did the charged rod bend the water?Give the pupil the table of electronegativities from ILPAC. What are the trends?Give the pupils a variety of molecules and get them to predict whether or not they are polar from the difference in electronegativities. However you also need an asymmetric shape for a molecule to have an overall dipole. For in SiCl4 the bonds are polar but SiCl4 has no overall dipole the shape is completely symmetrical. Then go back to the demo and talk about why different liquids did/didn’t deflect the water.

15. Intermolecular Forces

Describe intermolecular forces based on permanent dipoles, as in hydrogen chloride, and induced dipoles (van der Waals forces) as in the noble gases

Compare the boiling points of the noble gases and explain it’s linked to the number of electrons each atom contains. Then explain the origin of Van Der Waals forces in terms of instantaneous dipole, induced dipole and subsequent attraction between the induced and the instantaneous dipoles.Then demo gas jars of the halogens : chlorine, bromine, iodine to emphasise the point.Can also touch on alkanes and how Van Der Waals forces depend on the surface area of electron charge. E.g. 2-methly propane has a lower boiling point than butane.Molecules with permanent dipoles generally have higher mps, bps because permanent dipoles are stronger than Van der waals forces.Give the pupils a variety of molecules and ask them whether the intermolecular forces is just a van der waals force or if it is permanent dipole/permanent dipole.

16. Hydrogen bonding

Describe hydrogen bonding, including the role of a lone pair, between molecules containing –OH and –NH groups i.e. as in H2O and NH3 and analogous molecules.Describe and explain the anomalous properties of H2O resulting from hydrogen bonding e.g:

(i) the density of ice compared to water(ii) it’s relatively high freezing point and

boiling pointsothers include: high viscosity and high surface tension

Compare H2O’s boiling point with the other hydrides in group 6. Why is it so high? Explain that a hydrogen bond is the ‘electrostatic attraction between a lone pair on an oxygen or nitrogen atom (sometimes fluorine) and an electropositive hydrogen atom (i.e. a hydrogen atom that is also bonded to oxygen/nitrogen)Draw and label the hydrogen bond between water molecules and ammonia molecules. Then give pupils a range of analogous molecules e.g. ethanol, ethylamine, and get them to draw the hydrogen bonds.Also demo that ammonia and ethanol are completely soluble in water and show the hydrogen bonds that are formed.Demo ILPAC experiment ‘the effect of hydrogen bonding on liquid flow’ ILPAC 3 Exp3 p.46

For H/W pupils write an essay on how the physical properties of water are affected by hydrogen bonding.

17. Ionic Lattices and simple molecular lattices

Describe structures as:(i) Giant ionic lattices, with strong ionic

bonding i.e. as in NaCl(ii) Simples molecular lattices, i.e. as in I2

and iceCompare the solubilities in polar/non-polar solvents (water/hexane), melting point and electrical conductivity.

Demo that NaCl conducts electricity in water and that lead bromide conducts when molten (as in Year 11 experiment). Emphasise that this is due to free ions and not free electrons. Demo the high mp of sodium chloride and relate to the giant ionic structure of NaCl.Demo the low mp of iodine and suphur and relate to weak intermolecular bonds. Demo that a sulphur block doesn’t conduct electricity as it has no free electrons. Demo that iodine is more soluble in hexane than in water. (i.e. non polar substances are more soluble in non polar solvents than in polar ones)Then test a number of different liquids solutions for conductivity in order to determine whether the structure is simple covalent/giant ionic.Examples could include:Copper sulphate solution, potassium chloride solution, sulphuric acid, hexane, ethanol, cyclohexane or do ILPAC Exp 2 p.35 ILPAC 3Summarise with a general rule that if a compound contains a metal then the structure is giant ionic.

18. Giant Covalent and Metallic Lattices

Describe the structures of:Giant covalent lattices i.e. as in diamond and graphiteGiant metallic lattices

Compare the properties and structures of diamond and graphite (are some molecular models in lab 24 on top of the shelves)Then define metallic bonding as ‘ the electrostatic attraction between a sea of electrons and a lattice of positive ions’. Then give the metals a variety of metals e.g. sodium, magnesium and aluminium and get them to draw the metallic lattices.

Why is aluminium a better conductor and why does it have a higher mp? Derive the variables upon which the strength of metallic bonding depends e.g. ionic charge, ionic radius and number of free electrons.Also cut up pieces of lithium, sodium and potassium and show them that lithium is the hardest i.e. hardness decreases down the group. Why?Then react the group 1 metals with water and show that lithium doesn’t melt whereas sodium and potassium do. Why does the melting point decrease down Group 1?

19.Interpreting physical data and determining structure type

Describe, interpret and predict physical properties, including melting and boiling points, electrical conductivity and solubility in terms of:

(i) different structures of particles (atoms, molecules, ions and electrons) and the forces between them,

(ii) different types of bonding (ionic bonding, covalent bonding, metallic bonding, hydrogen bonding, other intermolecular interactions)

Deduce the type of structure and bonding from given information

Give the pupils a variety of exam questions on structure and bonding. In particular the ones where hey have to deduce the structure (giant/simple) and bonding (metallic/ionic/covalent) from given data on polarity, electrical conductivity, boiling point, solubility in polar and non-polar solvents.

20. Development of the periodic table (from practical observations)

Development of the periodic table from Dobereiner, Newlands, mendeleev, mosely, seaborgDescribe the periodic table in terms of the arrangement of electrons:

(i) by increasing atomic (proton) number(ii) in periods showing repeating trends in

physical and chemical properties(iii) in groups having similar physical and

chemical properties

Need to find a resource on the development of the periodic table over history Could revise what they know from GCSE about Group1 Demo Group 1 with waterThen to show periodicity could react:NaCl, MgCl2, AlCl3 and PCl5 with water and test the pHs to show pH 7, pH6, pH3, pH1 across a period. The could predict what the pHs of SiCl4 would be.ILPAC experiment ‘investigating the properties of Period 3

Explain that atoms of of elements in a group have similar outer electron configurations, resulting in similar properties

chlorides’ Exp 12 p.124 ILPAC4

21. Variation in 1st ionisation energies across a period and down a group

Describe and explain the variation of the 1st ionisation energies of elements shown by:

(i) a general increase across a period, in terms of increasing nuclear charge

(ii) a decrease down a group in terms of increasing atomic radius and increasing electron shielding outweighing increasing nuclear charge;

Use ILPAC resources in ILPAC 4 p.108. Practice interpreting graphs and providing explanations

22. Variation of physical properties across period 2 and 3

For the elements of Periods 2 and 3(i) describe the variation in electron

configurations, atomic radii, melting points and boiling points

(ii) explain variations in melting points in terms of structure and bonding

Interpret data on electron configurations, atomic radii, first ionisation energies, melting points and boiling points to demonstrate periodicity

ILPAC 4p.109-114 on ‘periodicity’ has some good resources on how the properties of period 2 and 3 elements vary across a period. Emphasis needs to be on interpreting graphs and relating to change in properties from;

Giant MetallicGiant covalent simple molecular/atomic

COULD CONDENSE LESSONS 20 AND 21 INTO 1 LESSON

23.The redox reactions of the Group 2 elements MgBa

Describe the redox reactions of the Group 2 elements Mg_->Ba

(i) with oxygen(ii) with waterExplain the trend in the reactivity of the Group 2 elements down the group due to the increasing ease of forming cations, in terms of atomic size, shielding and nuclear attraction.

We have samples of Mg,Ca and Ba. Pupils can react them with water and observe the reactivity. Use universal indicator/effervescence to judge the extent of a reaction. Use hot water for magnesium and you’ll get a reaction.Then burn the metals in air and compare their reactivity (be careful with barium – emphasise it’s toxic)Sum up reactivity (may be some conflict – cos Mg is likely to be the most reactive with oxygen)Then derive the two redox equations using oxidation numbers and half equations:

Mg -- Mg2+ + 2e-

1/2O2 + 2e- -- O2-

andMg -- Mg2+ + 2e-

2H2O + 2e- -- 2OH- + H2

Emphasise development of observational practical skills throughout this lesson. i.e. use of terms vigorous effervescence, colourless solution, white precipitate (if you get one), flame colours and what the reactants and the products looked like. Pupils use pH scales to infer pH from colours of solutions.

24. Reactions of Group 2 Compounds

Describe the action of water on oxides of elements in Group 2 and the approximate pH of any resulting solutions.Explain the use of Ca(OH)2 in agriculture to neutralise acidic soils and the use of Mg(OH)2 in some indigestion tablets as an antacid.Describe the thermal decomposition of the carbonates of elements in Group 2 and the trend in their ease of decomposition. (no explanation required)

Demo pH of magnesium oxide and calcium oxide, they have a pH of 9/10 because they are only partially soluble in water.Give equation: CaO(s) + H2O(l) - Ca(OH)2 (aq)Show them limewater too and milk of magnesia and emphasise uses in neutralising acidic soils and indigestion remedies respectively.Give the pupils ILPAC 4 Exp2 p.20 practical ‘investigating the thermal decomposition of Group 2 carbonates by heating them and bubbling the gas through limewater. They get more thermally stable as you go down the group. The pupils record the time taken for the limewater to show first signs of forming a white precipitate. Again emphasise observational skills colourless solution to white precipitate. You may want to split this practical up so that different pupils heat different carbonates. (model practical results are available)Give equation:CaCO3(s) -- CaO(s) + CO2(g)

25. Consolidation

Interpret and make predictions from the chemical and physical properties of Group 2 elements and compounds.

Demo heating magnesium nitrate (as in previous ILPAC experiment), calcium nitrate and strontium nitrate. Derive equation. What is the trend in the ease of thermal decomposition?Then split the class into 2 halves. Discuss the likely chemical and physical properties of beryllium and its compounds/ radium and its compoundsi.e.:Of the metal: 1st ionisation energy, melting point, density, electrical conductivity, reaction with water, reaction with oxygen, reaction of the metal with acidsOf the metal compound:Melting point/conductivity of the metal oxide. Reaction of oxide with water, ease of thermal decomposition of carbonate/nitrate, reaction of the metal oxide/hydroxide with acids.Then get the pupils to present their findings to the other half of the class.H/W Write an essay to summarise the likely physical and chemical properties of beryllium and radium and their compounds.

Teacher 2

Lesson Syllabus Activities1. Relative isotopic mass State that carbon-12 is used as the standard in

the measurement of relative atomic massesGeneral discussion about work habits and course structure.Give the pupils some masses of atoms. Ask them to work out the relative isotopic mass compared to hydrogen. Then explain why hydrogen can’t be used as a standard but carbon-12 can. So compare masses of atoms with that of 1/12th of a carbon-12 atom.Define relative isotopic mass.

2. Relative atomic mass Define the terms relative isotopic mass and relative atomic mass based on the carbon-12 scale.Calculate the relative atomic mass of an element given the relative abundances of its isotopes.

Look at the numbers in the periodic table. Why are they not exact whole numbers if protons and neutrons have the same mass?Introduce idea of isotopic abundance and explain why chlorine is 35.5 (it’s 75% chlorine-35 and 25% chlorine 37). Show the maths.Define RAM as being the weighted average of the mass of the naturally occurring isotopes of an element compared to that of 1/12th of a carbon-12 atom.Then set the pupils a number of different problems/exam questions on calculating RAM. Explain that a mass spectrometer is used to do this (but you don’t need to explain how a mass spectrometer works). Can talk about space probes and mars missions if you like. Isotopic abundances and hence RAMs will vary from planet to planet.

3.RMM and RFM Use the terms relative formula mass and Ionic bonding = RFM (mass of smallest

relative molecular mass and calculate values from relative atomic masses.

repeating unit of structure)Covalent bonding = RMMGive them a variety of compounds and then get the students to calculate the RFM/RMM and say ‘which it is’.Make sure you cover brackets and waters of crystallisation.

4. The Avogadro constant The Avogadro constant NA, as the number of particles per mole (6.02 x 1023)Mole (symbol mol) as the unit for the amount of substanceAmount of substance

ILPAC 1 p.5-9Use the formula n = N/L

5. Molar Mass Define and use the term molar mass (units gmol-1) as the mass per mol of a substance

Molar mass is the mass of a substance containing the same number of elementary particles as found in 12g of carbon-12.ILPAC p.9-12Use the triangles moles = mass/molar mass moles = N/Land combine the triangles (Ex 10,11,12 ILPAC)

6. Balancing Equations Construct balanced chemical equations for reactions studied and for unfamiliar reactions given reactants and products.

Go through the syllabus make a list of equations that they need to know on Group 2, Halogens, and acids + bases (perhaps including general types: acid+base, acid+alkali e.t.c.)Pupils practice balancing. Perhaps get molecular models out and use CFY resource as a starter.

7/8 Carry out calculations using mass and moles

Carry out calculations, using amount of substance in mol, involving mass

Use 3 step process:1. Work out moles of what you’re given.2. Look at the equation and ‘get the ratio’3. Work out the mass of what you’ve been asked to from mass=moles x molar mass

This is replicated in the exam questions. ILPAC p.17-19

9. Gas Volumes Carry out calculations, using amount of substance in mol involving gas volume.

Discuss Avogadro’s law : at RTP one mole of gas occupies 24 litres, 24dm3, 24000cm3

Explain why it is invariant and is independent of the substance i.e. in a gas there are no forces between the molecules (hence it is the random motion that determines the volume of a gas).Then calculate gas volumes in just reactions involving gases then calculate gas volumes from reacting masses (using same 3 step method as last lesson but with a final stage volume = moles x 24000)Could show 0.10g of magnesium giving off 100g of hydrogen in a gas syringe when reacted with excess acid.Talk about errors in mass of magnesium(0.005/0.10)x100 = 5%And ‘getting the cork on’ to explain any differences.

10/11 Calculate the relative atomic mass of magnesium by experiment

Do old practical assessment and evaluate it. (important because in the spring term there’s a similar assessment that we’ll do)

Evaluation assessment Evaluative 2 Identifying a Group 2 carbonate12. Solution volume and concentration

Use the terms concentrated and dilute as qualitative descriptions of r the concentration of a solution.Carry out calculations, using amount of substance in mol, involving:Solution, volume and concentration

Use ILPAC p.23-27Moles =cVNeed to have reached ILPAC Ex.26 by the end of the lesson.

13. Preparing a standard Carry out calculations, using amount of ILPAC Exp 2 p.25-27

solution of potassium hydrogen phthalate

substance in mol, involving:Solution, volume and concentration

14/15 Titration of potassium hydrogen phthalate with NaOH

Carry out calculations, using amount of substance in mol, involving:

Solution, volume and concentrationPerform acid-base titrations, and carry out structured titrations

First of all introduce concept of how to calculate an unknown concentration from a titration by running through a simple example say NaOH and HCl. Then demo ‘how to do a perfect titration’ and give them a list of ‘titration do’s and don’ts’.Then pupils do titration and try and work out the concentration of NaOH.

16. More titration calculations and evaluation of experiment in lesson 14/15.

Evaluate procedure last lesson.Procedural errors ILPAC p.34Calculate % error in pipette, burette, weighing and volumetric flask to come up with an overall uncertainty in the result (add % errors to get total error).Then do more calculations on titrations.

17/18 Stoichiometry Deduce stoichiometric relationships from calculations.

ILPAC experiment ‘a redox titration’ p.34-35Emphasize consistency of titres. Pupils should be getting to within +/- 0.1ml by now.Then do some other calculations on calculating stoichiometries in equations (text books should be a good place to look).

Practical Assessment Quantitative 3 Determining the concentration of a solution of sulphuric acid.

19. Empirical formula 1 Explain the terms:(i) empirical formula as the simplest

whole number ratio of atoms of each element present in a compound

(ii) molecular formula as the actual number of atoms of each element

Emphasise difference between empirical formula and molecular formula e.g. for benzene C6H6

Then calculate empirical formulae by using moles=mass/molar mass ‘in tables’ and dividing by the ‘smallest number of moles’. Use ILPAC p.20-22. Need to be able to do from both actual

present in a moleculeCalculate empirical and molecular formulae using composition by mass and percentage compositions.

masses and % by mass.Also molecular formula = whole number x empirical formulaIf the molar mass is known then the whole number can be worked out and thus the molecular formula worked out.

20. Empirical formula 2 (from experiment)

Calculate empirical and molecular formulae using composition by mass and percentage compositions.Explain the terms anhydrous, hydrated and waters of crystallisation.Calculate the formula of a hydrated salt from given percentage composition, mass composition or experimental data.

Do past practical exam.

21. Acids and Bases Explain that an acid releases H+ ions in aqueous solution.State the formula of common acids: hydrochloric, sulphuric and nitric acids.State that common bases are metal oxides, metal hydroxides and ammonia.State that an alkali is a soluble base that releases OH- ions in solution.State the common formulae of the common alkalis sodium hydroxide, potassium hydroxide and aqueous ammonia.

Go through key points opposite.Show them a variety of different acidse.g. citric, hydrogen chloride (make from NaCl and conc sulphuric), nitric, sulphuric to emphasise that water needs to be present for an acid to be formed.Then explain a ‘bronsted acid’. Pupils practice equations to show disassociation of acids.Demo CuO, solid NaOH ,Ca(OH)2 and conc NH3

as a few examples of bases. Demo that conc NaOH is soluble whereas CuO isn’t : hence alkalis are a subset of bases (soluble bases).

22/23 Formation of salts Explain that a salt is produced when the H+ of an acid is replaced by a metal ion of NH4+Describe the reactions of an acid with carbonates, bases and alkalis to form a salt.Explain that a base readily accepts H+ ions from

Pupils perform a variety of test tube tests and derive general equations:Acid+metal carbonate -Acid+metal hydroxide -Acid +metal oxide --

an acid; e.g. OH- forming H2O; NH3 forming NH4

+Acid + ammonium hydroxide-Pupils should be able to also write equations using ionic equations.

24. Preparation of copper (II) sulphate and ammonium sulfate

Prepare ammonium sulphate using 1M ammonia in a burette and 1M sulphuric acid (using a pipette) and screened methyl orange indicator.Prepare copper sulphate in usual way from copper oxide and 2M sulphuric acid.

25/26 Rules for assigning oxidation numbers and naming compounds

Apply rules for assigning oxidation number to atoms in elements, compounds and ionsUse Roman numeral to indicate the magnitude of the oxidation state of an element, when a name may be ambiguous e.g. nitrate (III) or nitrate (V)Write formulae using oxidation numbers

Go through rules for assigning oxidation numbers. Relate to periodic table.Then move onto naming compounds e.g. nitrates, sulphates, chlorates e.t.c.Can show that heating chlorates produces a chloride + oxygen and do the jelly baby as a pre-curser to next lesson:2KClO3 -- 2KCl + 3O2

Show it’s a redox reaction.Higher oxidation states of N and Cl when combined with oxygen are generally unstable! Can go into rocket fuels/explosives e.t.c.

27 Redox Reactions Explain that:Metals generally form ions by losing electrons with an increase in oxidation number to form positive ions.Non-metals generally react by gaining electrons with a decrease in oxidation number to form negative ions.Describe the redox reactions of metals with dilute hydrochloric and dilute sulphuric acidsInterpret and make predictions from redox

Pupils react magnesium with 2M hydrochloric, 2M sulphuric and 2M nitric acids. Make observations. Why is the reaction rate faster with sulphuric acid? (it’s dibasic and has a higher conc of H+ ions)Then derive:Mg + 2H+ --- Mg2+ + H2

Emphasise that the acid anion is a spectator.And break into half equations to show redox.Then give a selection of further redox reactions

equations in terms of oxidation numbers and electron loss/gain.

from Group 7, metal displacement reactions and derive redox equations by omitting the spectator ion.

28/29 The Halogens Explain in terms of van der waals forces the trend in the boiling points of chlorine, bromine and iodineDescribe the redox reactions, including ionic equations, of the Group 7 elements with other halide ions, in the presence of an organic solvent, to illustrate the relative reactivity of Group 7 elements.Explain the trend in reactivity of Group 7 elements down the group from the decreasing ease of forming negative ions, in terms of atomic size, shielding and nuclear attraction.

Demo samples of Chlorine, bromine and iodine and explain their volatility in terms of van der waals forces (surface area of electron charge)Then do ILPAC 4 Exp 6 p.62Write equations (including ionic ones)Explain in terms of redoxAnd explain the increasing oxidising power of the halogens as you ascend the group in terms of atomic size, shielding and nuclear attraction.ILPAC 4 p.61

30. Disproportionation Describe the term disproportionation as a reaction in which an element is simultaneously oxidised and reduced, illustrated by:

(i) the reaction of chlorine with water as used in water purification;

(ii) the reaction of chlorine with cold dilute sodium hydroxide , as used to form bleach

(iii) reactions analogous to those specified above

Contrast the benefits of chlorine use in water treatment (killing bacteria) with associated risks (hazards of toxic chlorine gas and possible risks from formation of chlorinated hydrocarbons). Also the use of fluorine in drinking water.

ILPAC 4 exp5 p.59 (could just demo)Derive equation:

Cl2 + 2NaOH --- NaClO + NaCl + H2O

And do ionically – explain the changes in oxidation number of the chlorine. Define disproportionation.

Then do water and chlorine:Cl2 + H2O -- HCl + HClOAgain explain in terms of disproportionation.ClO- is an oxidizing agent, and oxidises the nitrogen compounds in bacteria:ClO- + 2e- + 2H+ --- Cl- + H2O In practice Ca(ClO)2 is usually added in tablet

form to swimming pools rather than keeping gaseous chlorine!Then have discussions about risks and benefits of using chlorine.

31. Tests for halide ions Describe the precipitation reactions, including ionic equations, of the aqueous halide ions with silver ions, followed by aqueous ammonia.Describe the use of the precipitation reactions as a test for different halide ions.

ILPAC 4 Exp 8 p.75 and subsequent ILPAC exercises.Write ionic equation:Ag+(aq) + X- (aq) -- AgX(s)

32. Interpret and make predictions about the likely physical and chemical properties of astantine and fluorine.

Interpret and make predictions from the chemical and physical properties of Group 7 elements and their compounds.

Predict the likely properties of :The element fluorine (mp,colour,oxidising power, reaction with water – it oxidises water!)Sodium fluoride (white ppt with AgNO3 ?)The element astantine (mp, colour, oxidising power, reducing power of At-)Sodium astantide (green ppt with AgNO3)Get pupils to make presentations based on 15mins work with equations.H/W Write an essay on the physical and chemical properties of fluorine/astantine and their compounds.

Assessment Qualitative 1 Investigation of solutions and solids