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PLEASE REMEMBER, THESE ARE NOT 'STAND ALONE' NOTES, and were designed for my classes for use alongside the Salters resources - Chemical Ideas, Chemical Storylines, Practical Activities-Investigations and the AS-A2 Revision guides all published by Heinemann Secondary Series, to reduce the reading workload and offer a study strategy. From your teacher (not me!), its handy to have the answers to the Chemical Ideas, Storylines Assignments and Activities Questions side by side with the texts and these strategy pages. You haven't time to redo the Q's but a quick read of the Q's and connecting with the official answers is valuable revision - there is too much hit and miss revision from doing past papers in my opinion.

Chemical storylines SS1 "What is steel?"

• Steel is an alloy based on iron.

• An alloy is a mixture of a metal with at least one other element (metal or non-metal) or compound.

• Small differences in composition can have significant effect on the properties of an alloy.

• Too high a % of C in iron makes it too brittle, but a low % C makes a stronger steel.

• Appreciate the versatile nature of steel by changing its composition.

• There is a need for excluding impurities eg P or S which lead to poor quality material.

• The common elements added to iron to make steel, apart from carbon,  are usually d-block metals.

• Note the position of d-block ‘Transition Metals’ in the Periodic Table and don't forget the [green box on alloys]!

• Must know the KMnO₄ titration procedure as an example of a redox titration as well as a basic volumetric analysis method.

• Note:

  1. the endpoint is the first permanent pink, first 'smidge' of MnO₄^- ion in excess, NOT phenolphthalein, this is NOT an acid-base reaction, its a redox self-indicating titration.

  2. be able to work out the redox equation from the given half-cells (sometimes tricky!),

  3. and all the calculations are important, whether its to calculate the molarity of an iron solution or % Fe in table, steel sample or compound.

• extra redox volumetric equation practice for the 'keenies', late 30's to early 70's Q style, went out with 'Spandau Ballet' and JMB syllabus B essay Q's, but now broken down into multi-part structured Q's these days, I'm not complaining and neither should the students!, but the answers and reasoning are a bit compressed!

• Starting with impure iron from blast furnace, the molten iron contains many other elements and the iron is too brittle initially, so there is a need to reduce C and remove others like S and P.

• This is achieved by the Basic Oxygen Steel making process (BOS) which involves many redox reactions. It is a 'batch process' and can't be used as a continuous production line.

• Early removal of S using Mg, and C, P, Si and others oxidised by molecular oxygen before scrap iron/steel introduced.

• After the oxygen blow the basic oxides CaO/MgO are added to form slag salts with the weak acidic oxides of Si and P.

  • The oxides of Mn/Fe also collect in the slag, so some iron is wasted and the Mn might be added in a controlled way later for a particular steel specification.

  • The toxic carbon monoxide formed must be dealt with and not allowed out into the atmosphere, it can be burned as a fuel to harmless carbon dioxide.

• It is important to keep track of temperature and composition (by atomic emission spectroscopy – revise)

• The elements are oxidised in a sequence in exothermic reactions (no extra heat needed), so temperature control is essential to avoid wasting energy and converter lining damage.

• The added scrap iron/steel addition acts as coolant because melting is endothermic.

• The whole process must meet the specification for an individual customer requirement.

• Dissolved oxygen is removed with aluminium and then C, Mn and Si etc. can be re-added to a desired specification, plus any other elements.

• Argon (of light bulb fame) is bubbled through to stir the mixture because it so unreactive and most 'stirrers' will melt, dissolve, and change the composition.

• In the future electric arc furnaces maybe used more to recycle steel. Big carbon electrodes are 'sparked' to melt the scrap iron/steel, lime added to remove impurities as slag. It is possible to use this technology on a small scale to produce 

• Assignments 2 to 6 are good interpretation revision questions.

• Revise in terms of  the theory of emission spectra, its formation and recording it.

• Each element can be qualitatively identified from characteristic frequency/wavelength 'finger print lines' and the intensity of the line forms the basis for quantitative analysis.

• Needs revising from GCSE source (eg Metal Extraction). The silica SiO₂ impurities are from 'sand' and 'soil' associated with the iron ore. Sulphur compounds occur in the coal from which the coke was made. Phosphorus and manganese impurities come from other minerals mixed in with the original iron ore.

• All the chemistry is covered in Storylines SS2 BUT how to do an input/output flow diagram in Part is important, could be asked for in the exam given the information or given one to interpret, either way it causes problems.

• Quick read of information as a source of ideas and appreciate the versatility of steel and its various alloy compositions with different properties for different applications.

• CS Fig 17 page 172 sums it up very well. You need to know the redox stuff inside out and apply the ideas of half-cell potentials.

• Rusting = corrosion of iron by an oxidation process which is energetically favourable.

• Rusting is the opposite of its extraction by reduction - Fe redox cycle Fig 16.

• The detailed electrochemistry of rusting

  • half-cell of oxidation of Fe to Fe²⁺  

  • half-cell reduction of O₂ (+ H₂O + e^-) to OH^- with e^- flow through iron

  • iron(II) hydroxide, Fe(OH)₂ is oxidised to hydrated Fe₂O₃  

  • relate the half-cell reactions involved to their E values and calculate cell Emf for overall feasible reactions

• Unfortunately rust flakes off and so it all eventually corrodes away (later xref/contrast ZnO, Al₂O₃, Cr₂O₃ on metal surface, which do not flake away and offer good anti-corrosion properties)

• Apply factors affecting rate of rusting eg the following all speed up the process!

  • decreasing pH, H⁺_(aq) ions remove OH^-_(aq) formed from the reduction of O2(g), 

  • increased concentration of any ions improves the conductivity of the aqueous media, which is part of 'redox circuit',

  • and if the iron is in contact with a 'less reactive' metal (meaning a more +ve half-cell potential), corrosion rates increase, because the iron is preferentially oxidised with the more -ve half-cell potential.

• Rust protection-inhibition ... examples ... are x-ref with assignment 7 on p174.

  • A plastic or paint physical barrier to exclude water and oxygen (air),

  • Either by (i) dipping in molten zinc, or (ii) electrolysis with Zn²⁺_(aq) solution and the iron/steel object as -ve cathode, galvanising with Zn layer which results in the formation of ZnO layer, the redox chemistry is similar to Fe rusting (see Fig 21) but the layer does not flake away giving a protective layer of zinc oxide. Even if scratched, the Zn with a more -ve half-cell potential is preferentially oxidised.

  • Sacrificial corrosion with blocks of Zn or Mg and relate their 'sacrifice' to their more negative half-cell potentials, ie preferentially more favourable oxidation.

  • Stainless steel via Cr addition, forms protective layer of chromium(III) oxide.

• History lesson in food preservation: ‘invention’ of the tin can (tin coated steel) ...

  1. early tin cans suffered from preferential oxidation of Fe due to its more –ve potential, cured by lacquer coating, and the ...

  2. fruit juice problem, carboxylic acids complex with Sn²⁺_(aq) ions, changes Sn_(s)/Sn²⁺_(aq) potential making it more negative than Fe_(s)/Fe²⁺_(aq), so Sn preferentially corrodes, not toxic and contribute to ‘tangy’ taste BUT don’t keep too long as Fe eventually will dissolve too!

CI 9.2 Redox Reactions and electrode potentials

• Analysing redox reaction in terms of two half-reactions, an oxidation and a reduction.

• How the half-cell functions as an oxidising or reducing agent.

• Combining half reactions to produce balanced overall redox equation

  • doc's penta check of redox equations

    1. correct species and state, omitting spectator ions

    2. correct direction of change (ie which is on left and right)

    3. oxidation number change (total increase = total decrease if redox species in the correct ratio)

    4. ion charge (total on left = total on right, handy extra check)

    5. and atom count (usual left = right count, but NOT totally reliable in redox changes, hence other checks 3. and 4. above)

• A simple cell or battery is made from combining two half-cells.

• The half-cell potential is a measure of the tendency of a species to lose/gain electrons (see discussion on comparing zinc and copper)

• Describe how to set up an electrochemical cell, and relate direction of chemical changes (on electrodes) to the +ve and -ve terminals and the direction of electron flow.

• Measurement of E_cell (cell Emf) with high resistance voltmeter.

• Using two metal_(s)/metal ion_(aq) half-cells to make a complete cell via a 'salt bridge'.

• Using a standard half cell, must know details of H_2(g)/H⁺_(aq) half-cell (Fig 11 p213), arbitrarily given the half-cell potential of 0.00V against which, other half-cells can be measured.

  • ^ø Standard conditions: 298 K, 1 atmosphere pressure, 1 mol dm^-3 concentration of H⁺_(aq) or metal ion in the other half-cell etc.

• In principle, any accurately known half-cell potential can be used in a cell system to obtain an unknown half-cell potential.

• The electrochemical series and electrode potential charts, know how to construct, read and use them.

• Other half-cells, they don’t have to simple metal/ metal ions, all you need is two interchangeable oxidation states eg Cl_2(aq)/Cl^-_(aq) or Mn²⁺_(aq)/MnO₄^-_(aq) etc. but both components of the half-cell must be in the same solution and in contact with a platinum electrode that connects to the rest of the circuit (Fig 12 p213).

• One way of working out E^ø values: E^ø_cell =  E^ø_(red) – E^ø_(ox)*  (amounts to the difference between the half-cell potentials on an electrode potential chart eg Fig 10 p212 or Fig 13 p214).

  • * E^ø_(red) is the most positive or the least negative = the strongest oxidising agent or electron acceptor of the two half-cell systems, and the +ve battery pole, eg Cu²⁺ compared to Zn²⁺, so Cu²⁺_(aq) + 2e^- ==> Cu_(s), rather than reduction of Zn²⁺,

  • and E^ø_(ox) is the least positive or the most negative = the strongest reducing agent or electron donor of the two half-cell potentials, and the -ve battery pole eg Zn compared to Cu, so Zn_(s) - 2e^- ==> Zn²⁺_(aq) happens rather than oxidation of Cu,

  • overall cell redox reaction: Cu²⁺_(aq) + Zn_(s) ==> Cu_(s) +Zn²⁺_(aq)  

  • Calculating the voltage-emf for the copper-zinc cell: 

    • E^ø_(red) = E^ø_{Cu(s)/Cu2+(aq)} = +0.34V, E^ø_(ox) = E^ø_{Zn(s)/Zn2+(aq)} = -0.76V

    • E^ø_cell =  E^ø_(red) – E^ø_(ox)= +0.34V - (-0.76) = + 1.10 V (feasible!)

• Relate E^ø_cell to direction of overall chemical change and feasibility of reaction (leads on into CI 9.3). If you calculate a -ve cell voltage for a given reaction, that is not the way cell reaction goes! please reverse the equation!

Chemical Ideas 9.3 Predicting the direction of redox reactions

• Electrode potentials are a measure of the tendency of a half-cell reaction to accept electrons.

• The more +ve the E_half-cell, the greater the tendency to attract electrons (see end of CI 9.2 above).

• Constructing electrode potential charts and using them to solve problems.

• The E_battery-cell must compute to a value of >= 0.00V for the cell, and any other redox reaction, to be feasible (see end of CI 9.2 above).

• These theoretical calculations can be used for any redox reaction BUT there are limits ...

• You can’t say the reaction will definitely happen because there may be rate limits especially if high activation energy.

• However you can employ a catalyst or raise the reaction temperature, or specific energy input eg light in photosynthesis to get the reaction going!

• Scrap steel is part of BOS process and is cost effective, recycling reduces costs of (i) ore mining extraction, (ii) possibly overseas transport and (iii) blast furnace reduction of ore. These gains are partly offset by the cost of collecting scrap metal.

• In the electric arc process only scrap steel is used and is handy technology to produce small batches of particular steel by carefully controlling what scrap goes in.

• The composition of scrap important, needs to be graded and selected to avoid problems

• When recycling tin cans, you need to remove the tin and other waste.

  • The cans are shredded and paper/residual food removed, mechanical shredding and magnetic separation can be used,

  • and de-tinning is done by reaction with hot NaOH(aq), after which the valuable tin can be recovered by electrolysis of the 'waste solution'.

• A particular scrap case study, need for steel uncontaminated by radioactive isotopes from the nuclear and atomic weapon industries, scrap source from the German ships sunk at Scapa Flow has proved useful (good geography Q and I don't remember the event!).

• A nice read for revision but need to understand the characteristic chemistry of the d-block elements via CI 11.5 and 11.6 as a direct result of their electronic structure AND compare transition metal chemistry with the elements of the s and p blocks.

• The four typical chemical properties are nicely illustrated:

  1. formation of compounds in a variety of oxidation states

  2. catalytic activity of the elements and their compounds

  3. a strong tendency to form complexes

  4. the formation of coloured compounds and ions

• Iron is relatively cheap but many other useful transition metals are expensive due to eg low abundance in Earth's crust, difficulty of extraction, or relatively low commercial demand.

• 3 horizontal rows in periods 4 to 6 each of ten elements.

• The particular electron configuration feature is the filling of the d level.

• Be able to write out the electron configurations of any atom/ion from Sc/Sc³⁺ to Zn/Zn²⁺ in s, p and d notation (Fig 8 page 32 [Ar]????, watch out for Cr and Cu 4s¹ quirks and be able to explain them, or as box diagrams like Fig 20 p260, note that the electrons are unpaired as much as possible in the d orbitals to minimise repulsion)

• Be able to relate electron configurations to ...

  • ease of 3d electron gain/loss in change of oxidation state,

  • maximum oxidation state from Sc to Mn to the total number of outer 3d/4s electrons,

  • the relatively low ionisation energies (x-ref with Gp 1 and 2, with 1 or 2 outer electrons easily lost) before big rise in IE when removing an electron from an inner filled shell,

  • catalytic activity via electron loss/gain - change in oxidation state

• Sc and Zn are not really transition metals (but are in d-block) since they do not show variable oxidation state and coloured compounds because they don't comply with the definition below.

• A transition metal is defined as element which forms at least one ion with a partially, but incompletely, filled sub-shell of d electrons (eg Ti to Cu, the electron configurations of one of the transition metal's ions must be within [Ar]3d¹ to [Ar]3d⁹).

• Typical physical properties (most members of block are quite similar) eg

  • high mpt/bpt, high density, good electrical/heat conductors,

  • and good mechanical properties - hard, durable, high tensile strength,

    • AND using the metallic bonding model to explain the typical physical properties listed above and also their malleability. Main ideas are the strong metallic bond between the ionised metal atoms and the sea of free or 'delocalised' outer d and s electrons, plus the mobility of the electrons,

  • also note the effect of alloying on physical properties

    • layers of atoms can slip accounting for malleability and ductility

    • in alloys added atoms reduce the ‘slip’ increasing strength

• Four important typical chemical properties, (3. and 4. covered in CI 11.6)

  1. variable oxidation states: relate to enthalpies of ionisation, common ones for Ti to Cu, higher values stabilised by forming oxo-cations eg vanadium in (V) [VO₂]⁺_(aq) or oxy-anions eg manganese in [MnO₄]^-_(aq), there are 'vague' stability trends across the block eg the M²⁺ tends to become more stable wrt the M³⁺ oxidation state.

  2. catalytic activity: look for both heterogeneous and homogeneous examples of metals or their compounds.

  3. complex formation: more in Chemical Ideas 11.6

  4. coloured compounds: more in Chemical Ideas 11.6

• Idea of s, p, d and f –sub-shells or electronic energy levels.

• Know the maximum number of electrons in each sub-shell.

• Know the order of filling of the sub-shells from Z=1 to 36 to write out the electron configuration,

  • watch out for the two ‘quirks’ for Cr and Cu (see in CI 11.5)

  • and the order of electron removal when forming positive ions eg for the 3d block of transition metals, you remove the 4s electrons first, before any of the 3d electrons.

• The idea of atomic orbitals as the space/shape of a particular electronic level or sub-shell.

• The number of orbitals per sub-shell, 1 for s, 3 for p, and 5 for d sub-shell.

• Know how to relate electronic structure to an elements chemistry and position in the Periodic Table (eg s, p and d-blocks or metal/non-metal).

• Oh no, not again! What is a catalyst?, only small amounts needed, homogeneous and heterogeneous catalysts, how do catalysts work?, reaction enthalpy profiles of catalysed and uncatalysed reactions, examples of the types of catalysts, details of hydrogenation reaction, catalyst poisoning.

• What are complexes?

  • They consist of a central metal ion and ligands.

  • Complexes can have an overall negative or positive charge (complex ions), or they may be electrically neutral (neutral complex).

  • Be able to correlate the charge on a complex ion with any charge on a ligand and the oxidation state of the metal.

  • Ligands are electron pair donors, and you can think of it forming a dative covalent bond (but rarely that simple).

  • The coordination number = the number of bonds to the central metal ion (apart from a monodentate ligand, it does NOT equal the number of ligands, see below)

  • The shapes of complexes and examples (co-ordination number of 4 for square planar and tetrahedral, co-ordination number of 6 for octahedral)

  • The naming of complex ions.

  • Why complexes differ in stability? a more strongly bonded ligand will 'push out' a weaker bonded ligand in a ligand exchange reaction

  • The stability expressed as equilibrium constant K_stab 

  • Writing out equilibrium equations and K_stab expressions

  • Interpreting K_stab values eg in terms of relative stability of complex

  • Study examples of monodentate ligands and polydentate ligands

    • polydentate ligands give a chelate ring system

    • bidentate means two bonds formed with the central metal ion per ligand (three ligands would give an octahedral complex of coordination number 6, two might give a square planar or tetrahedral complex of coordination number 4)

    • hexadentate means six bonds formed with the central metal ion per ligand (just one ligand would give an octahedral complex!)

      • and be able to work out oxidation state of the metal in a complex

      • taking into account the charge on ligand, or if its neutral

• The theory of the 3d electronic transition origin of colour

  • The 3d orbitals of central ion split into two groups of quantum levels by the effect of the ligand.

  • Light may be absorbed from transitions between from these ‘split’ lower to upper sub-levels.

  • Examples of visible absorption spectrum of transition metal ions, and the colour you seem is what isn't absorbed.

  • If you change the ligand of a complex, or the ratio of two ligands in the same complex, you change the 'electronic environment', so different colours are seen.

• Geometrical isomerism in transition metal complexes, slightly different electronic situations, so you get slightly different colours.

• Smashing little demo and much more fizzy and colourful than any biology key and lock reaction! Heat it to 80^oC and it nicely froths over the beaker onto the paper towels!

• The Q's e to k are good revision and just remind yourself that the Co²⁺_(aq) catalyst gets oxidised to a green Co³⁺_(aq) complex, does it stuff, gets reduced in process and so returns to its 'pink state', so its another example of 'catalytic cycle'.

• Muse over the answers to Q's a to m. Lots of important general ideas included like K_stab, ligand replacement reaction, shape of complex, different ligands involved eg be familiar with NH₃ OH^- H₂O Cl^- EDTA^4- etc. BUT the reactions of sodium hydroxide, ammonia with copper(II), iron(II) and iron(III) are the most important to learn.

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ALL my KS3 SCIENCE Revision Quizzes (~US K12 grades 6,7,8)

GCSE-IGCSE-KS4 Science-CHEMISTRY notes & quizzes (~US K12 grades 9-10)

Advanced Level CHEMISTRY GCE AS A2 IB notes and quizzes (~US K12 grades 11-12)

All my GCSE-IGCSE Science-CHEMISTRY etc. syllabus help links

 All my GCE-AS-A2-IB AQA, Edexcel, OCR etc. Advanced Level Chemistry syllabus-specification help links