Redox

Consider checking out the Edward ATAR 12 Redox Seminar Masterclass for a shortened version of the entire Redox course.

Redox Introduction §

• Reduction & Oxidation

Read question carefully depending if it asks you to explain in reduction or oxidation.

• Reduction
  • gain of electrons
• Oxidation
  • loss of electrons.
• due to conservation of matter, reduction and oxidation is the exchange of electrons, and happens simultaneously.
• OILRIG
  • oxidation is loss, reduction is gain
• example:
  • oxidisation: $Na\to N{a}^{+}+\overline{e}$ (oxidation half equation)
  • reduction: $C{l}_2+2\overline{e}\to 2C{l}^{-}$ (reduction half equation)
  • balance: $2Na\to 2N{a}^{+}+2\overline{e}$
  • combine: $C{l}_2+2Na\to 2NaCl+2\overline{e}$ (full redox equation)
  • Consider the oxidation number
    • Cl2: 0
    • 2Na: 0
    • 2NaCl: +1 -1
    • increase in oxidation number: oxidation
    • decrease in oxidation number: reduction

common types of redox reactions §

• metal metal ion displacement
  • metal transfers ions
• halogen halide ion displacement
  • electrons transferred to a halogen, halides to less reactive halogen
• combustion
• corrosion

Oxidation Numbers §

• increase in oxidation number: oxidation
• decrease in oxidation number: reduction

Tip

Oxidation number (oxidation state): A number assigned to the atom to show how many electrons lost, gained or shared unequally (to form a bond).

• Oxidation number of Na: 0
• Oxidation number of Cl2: 0
• Oxidation number of Na in NaCl: +1
• Oxidation number of Cl in NaCl: -1
• Oxidation number of H in HCl: +1
• Oxidation number of Cl in HCl: -1
• Oxidation number of H in H2O: +1
• Oxidation number of O in H2O: -2
• Oxidation number of F in F2O: -1
• Oxidation number of O in F2O: +2
  • because F is more electronegative than oxygen

elements with fixed oxidation states in their compounds	oxidation state
hydrogen (except in metal hydrides - metal combined with hydrogen & fluorine)	+1 (-1)
oxygen (except in peroxides)	-2 (-1)
group I elements	+1
group II elements	+2
aluminium	+3
halide	-1

• for uncombined elements, the oxidation state is zero
• for atoms in elements (when element combined to itself), the oxidation state is zero.
• for simple ions, the oxidation state is the charge of the ion.
  • e.g. $N{a}^{+}C{l}^{-}$ +1 and -1
• electronegativity: ability to attract bonding pair of electrons hence something peroxide (???)
• halides: ions of halogens

halogens	halides
$F_2$	$F^{-}$
$C{l}_2$	$C{l}^{-}$
$B{r}_2$	$B{r}^{-}$
$I_2$	$I^{-}$

• if theres one other electron added to it it is not a halide!

by convention put oxidation number under element, and the total above element. this is a common mistake !!

Warning

BE SPECIFIC AND BE PARTICULAR. DO NOT BE VAGUE. EXPLAIN WELL AND LOGICALLY. E.G. “IT” NO GOOD 🙅

• in oxidation, the thing being oxidised is called the reducing agent!
  • vice versa
  • H2O something something idk
  • decrease in oxidation number: reducing: oxidation agent!
    • An oxidising agent accepts electrons and is itself reduced.
  • increase in oxidation number: oxidising: reducing agent!
    • A reducing agent donates electrons and is itself oxidised.
  • すごい

Writing Balanced Equations §

Practice §

Copper (I) Oxide

• Cu2O Manganese (IV) Oxide
• MnO2 Iron (III) Sulphate
• Fe2(SO4)3

New Types of Acids: §

• Chloric (I) acid
  • Cl is +1 because O is more electronegative :)
  • $HClO$
• Chloric (V) acid
  • $HCl{O}_3$

Combining Half Equations §

• Combine it to write balanced equations!
• Reactants go together with reactants products go together with reactants.
• cancel everything out.

Balancing Redox Equations in Acidic Conditions §

0. Assign Oxidation Numbers
1. Write the skeletons of the oxidation and reduction half-equations.
2. Balance all elements other than H and O
3. Balance the oxygen atoms by adding H2O (where needed)
4. Balance the hydrogen atoms by adding H+ ions (where needed)
5. Balance the charge by adding electrons, e-
6. If the number of electrons lost in the oxidation half-reaction is not equal to the number of electrons gained in the reduction half-reaction, multiply one or both of the half-reactions by a number that will make the number of electrons gained equal to the number of electrons lost.
7. Add the 2 half-reactions as if they were mathematical equations. The electrons will always cancel If the same formulas are found on opposite sides of the half-reactions, you can cancel them. If the same formulas are founded on the same side of both half-reactions, combine them.
8. Check to make sure the atoms and the charges balance.

disproportionation §

• 1 substance is simultaneously oxidised and reduced
• has the general equation: $A+A\to B+C$

displacement reaction of halogens §

• describing appearance of halogens
• ms pilling demonstration

halogens	halides
Cl2 + 2e $\rightleftharpoons$	2Cl-
Br2 + 2e $\rightleftharpoons$	2Br-
I2 + 2e $\rightleftharpoons$	2I-
(oxidising agents)	reducing agent

• if a substance is coloured, it is not due the halides. hence only halogens and colours are listed on page 5 of data booklet.
• halides could be oxidised to halogens
• series of chemical reactions, observe outcomes, and then make predictions

1. Cl2 + Br- -> Cl- + Br2

• ox: $2B{r}^{-}\to B{r}_2+2\overline{e}$
• red: $C{l}_2+2\overline{e}\to 2C{l}^{-}$
• yellow solution added to orange solution forming orange solution (?)
• cell potential calculations: $E_{\text{cell}}=E_{\text{red}}-E_{ox}=1.36-1.08=+0.28V$ positive, therefore feasible.

1. $C{l}_2+2I^{-}\to 2C{l}^{-}+I_2$

• ox:
• red: $C{l}_2+2\overline{e}\to 2C{l}^{-}$

3. $B{r}_2+C{l}^{-}$

• no visible reaction

4. $B{r}_2+2I^{-}\to 2B{r}^{-}+I_2$
5. $I_2+C{l}^{-}$

• no visible reaction

6. $I_2+Br-$

• no visible reaction

Electrochemical Cells §

Galvanic Cells & Daniel Cells §

• what is galvanic cell:
  • a galvanic cell uses two half-reactions in separate containers and makes use of the spontaneous nature of the reactions to generate an electric current. potential chemical energy converted to potential energy.
  • the transfer of electrons in the redox reactions is harnessed through physical separation of the two half-equations.

daniel cells §

• this is a daniel cell
• 
• purpose of salt bridge: joins the two half cells. complete the circuit and ___ ms pilling will give booklet with better description. movement of ions in salt bridge allows flow of charge through the cell and complete the circuit.
• purpose of high-resistance voltmeter: reduce the current because you don’t want dramatic drop in voltage.
• zinc is more reactive, so oxidises and loses electrons that flow blah blah blah blah
• the solid thingy is usually graphite or platinum.

Anode and Cathode §

• An ox & Red cat
• Anode is negative, Cathode is positive (in galvanic cells only - in electrolytic cells anode is positive and cathode is negative!)

Feasibility of a Reaction §

• A reaction is thermodynamically feasible if calculated $E_{\text{cell}}$ is positive.
• choose to change sign or not -> if it is changed it is - oxidised instead.
  • same outcome lol - depends on question on what you need to do, mathematically they are the same formula.
  • +- oxidised depending on if sign is changed.

Sacrificial Anode §

• sacrificial anode: strong reducing agent (more reactive) metal that will be oxidised preferentially attached to iron (or other metals) to protect it (from oxidising)

Ms Pilling on Self Study §

its like any other sport - you dont do training, you dont exercise, you dont get the prize

Cathodic protection §

• ICCP, Impressed Current Cathodic Protection systems are used to provide cathodic protection for pipelines, ship hulls, offshore production platforms, water treatment equipment etc.
• principal advantage of ICCP is greater output capacity compared to sacrificial anode systems - therefore when protection is desired for large, poorly coated and bare structures ICCP is often choice.
• it requires external DC power source that is energised by standard AC current.
• insert cathodic protection diagram
  • protected structure
  • 
• sea water
  • many different ions: sodium ion
    • sodium ion -> metal
    • too much energy, not provided
• enough energy to turn chloride into chlorine! (oxidised) on anode
  • moving to cathode, and chlorine is reduced to chloride
    • if theres enough current, could happen (?)

sacrificial anode CP method vs impressed current CP method §

• sacrific
  • no external source of power
  • spontaneous reactions
  • oxidised more readily than iron
  • cathodic protection
• impressed current
  • metal and anode, less reactive, acting as cathode.
  • non spontaneous reaction.

Corrosion §

• oxidisation of a metal
• usually occurs in the presence of oxygen and water vapour in the air to form metal oxide.

dry corrosion §

• dry corrosion/oxidation occurs when oxygens or other reactive gases in air reacts with exposed metal without the presence of a liquid. This causes an oxide layer to form over the metal, which damages its surface properties.
  • oxide layer called rust.
  • dry corrosion requires contact between air and metal, so oxide layer forms eventually stops (as it no longer is in contact because the layer prevents it)
    • this stopping is known as passivation.
• oxide layer is porous in some cases, and thus corrosion can continue deep into the material.
  • known as active corrosion.
  • porous is basically having holes
• dry corrosion is sensitive to temperature: reacts faster under application of heat (no way rate of reaction???!)

3 types of dry corrosion §

oxidative corrosion §

• point of contact between metal and o2 gas: o2 gas is reduce and oxidises metal
  • this forms metal oxide.
• metal oxide may be stable: adhere onto non-corroded metal; this prevents more contact between o2 and metal; therefore no more corrosion.
• the metal oxide may be stable: but NOT adhere onto the non-corroded metal (i.e. might flake away): it then fails to prevent direct contact between oxygen gas and metal, and metal corrosion continues

liquid metal corrosion §

• occurs when liquid metal is allowed to flow over solid metal at high temperature.
• one example is flow of sodium metal (used as coolant) over cadmium. cadmium is corroded and sodium ions are reduced.

corrosion by other gases §

Fuel cells §

• h2 and o2 come to make h2o
  • this is foundation of fuel cells.
• spontaneous reaction, produce lots of energy.
• anode is where hydrogen are
• cathode has oxygen
• electrons travels from anode to cathode.
• electrolytes block.
• catalyst helps hydrogen gas to split apart

advantages §

• uses external fuel source, so not reliant on limited store capacity
• high efficiency
• emission-free; zero emissions, only water vapour
• high energy density: a lot of energy in small volume
• renewable: hydrogen can be produced from renewable energy sources, such as wind, solar, hydropower

disadvantages §

• high cost to produce a fuel cell
• production/transportation/distribution/storage problems with hydrogen
• developing technology ? i think you have to expand on this. check practice tests later and come back.

applications of redox: primary/secondary/fuel cells blah bla §

cells and batteries §

• redox: transfer of electrons between 2 substances. separate galvanic cells and the movement of electrons between external circuit to generate electricity.
• electrochemical cell is a store of chemical energy in closed system: closed regards with chemical substance, all reactants and products contained within the battery, but not closed with respect to energy transfer (physics concept woohoo)
• split into primary and secondary cells (rechargeable and non-rechargeable respectively)
• secondary cell can be discharged and recharged =D
• size of battery means more chemicals inside which means it can run for a longer amount of time; it does not affect the voltage output.

a battery is more than one cell connected together in a series arrangement

• the electromotive force (voltage) increases when the cell is connected together in series. voltage = sum of voltages of each cell.
• a battery is a closed system which contains high energy reactants and low energy products in a sealed unit.

primary cell §

• lots and lots of primary cell but only two we learn =D they all have relatively similar patterns.
• one use wonder! non reversible.
• advantages
  • cheap
  • don’t discharge unless power is drawn

zinc-carbon dry cell - leclanche cell §

• 
• produces 1.5V
• carbon is inert (non-reactive) and conductive electrode: hence why it is used. if a metal is used, it would react and that wouldn’t be poggers.
• porous fibre is like filter paper fr no cap.
• allows outside to move between the thing and the slightly red paste; outside is zinc.
• excess zinc is used, so it won’t be entirely consumed and the zinc ions wont leak everywhere.
• half reacitons:
  • ox $Zn(s)\to Z{n}^{2+}+2e^{-}$
  • red: $2Mn{O}_2(s)+2N{H}_4^{+}+2e^{-}\to M{n}_2{O}_3+2N{H}_3+H_{2}O$
  • full $2Mn{O}_2(s)+2N{H}_4^{+}+Zn(s)\to Z{n}^{2+}+M{n}_2{O}_3+2N{H}_3+H_{2}O$
• disadvantages
  • low energy to mass ratio
• advantages

alkaline cell §

• uses zinc and mangaese(IV) oxide, but uses potassium hydroxide paste as electrolyte.
• alkaline cells do not need lots of electrolyte, so more space for reactants.
• thus it has higher energy to mass ratio as a result.
• 

secondary cells (rechargeable) §

• secondary cells are a massive part of the modern world
• utilises fairly new (with regards to redox topic)
• we make things go backwards (do, undo, redo, undo blah blah blah)
• process that goes both ways :fire: bars
• rechargeable cells got invented
• chemical energy -> discharge spontaneous energy -> electrical energy -> recharge non-spontaneous reaction -> chemical energy
• 
• once discharged, recharged by applying dc voltage (have to mention dc voltage!)
• we do not use reversible symbols in redox equations; discharge with forward arrow or recharge with a forward arrow

nickel metal hydride cells §

• dont need to memorise diagram
• nickel metal hydride batteries results in high demand for nickel.
• mysterious rare-earth metals: alloy, store hydrogen as hydride
• recharge reverses reaction; in this case me turn the metal alloy back into hydride, so we do half reaction in opposite direction.
• talk about polarity later
  • ms holland will talk about electrolysis first

lead acid battery §

• CAR BATTERIES
• NOT TESLA BATTERIES. REGULAR PETROL-DIESEL
• PURPOSE IS TO TURN THE STARTER MOTOR.
• delivers HIGH VOLTAGE and HIGH CURRENT to turn over the motor getting it starting, reducing the mechanical resistance.
• each cell produces 2V, car battery uses 6 cells in series to output 12V
• advantages
  • cheap reliable
  • very reliable, withstand many discharge-recharge cycles: long lifetime
  • output large current
• disadvantages
  • low energy density
  • safety hazards; corrosive sulphuric acid electrolyte; toxic lead compounds.
    • if broke, waterway can say byebye
  • environmental concern upon end of life.
    • horrible pollutant, lead will troll the environment.
• porous anode, to increase surface area
• cathode made with lead (4) oxide
• all good to go again once recharged.

lithium ion battery §

• lithium exists in cathode and in electrolyte.
• cathode made with lithium cobalt oxide
• when battery is fully charged, we shoved lithiums in graphite
• discharge reactions:
  • anode (-) Lic6 -> Li+ + c6 + e-
  • cathode (+) coO2 + Li+ + e- -> LiCoO2
  • redox: LiC6 + CoO2 -> C6 + LiCoO2
• different materials for two different electrodes results highly energetic favourable electrodes with different materials.
• recharged anode -> cathode, cathode -> anode.
• fire risk
  • banned on flights due to the fire risk they pose
  • 1. short-circuiting
    • porous separator keeps battery electrodes apart; charging for long periods can damage the porous membrane, removing it cause battery to discharge rapidly and generate a lot of heat.
  • 2. overcharging
    • when overcharged, lithium cobalt oxide releases oxygen. react with flammable electrolyte and also cobalt oxide
  • 3. electrolyte breakdown
    • electrolyte catches fire.

Fuel Cells §

• Does not form the definition of electrochemical cell.
• It is not a closed system, but it is instead an open system.
• It is not being used to replace batteries, we use it more often to replace a combustion engine.
• Direct replacement to engine in vehicles.
• Hydrogen Fuel Cell
• Combustion Reaction of Fuel is the only reaction simplified, unless it wants you to state the half equations.
• There are multiple kinds of fuel cells: 3 kinds of Hydrogen Oxygen cells (didn’t catch what she said)

Hydrogen Oxygen Cell §

• Electrochemical Energy: Converts between CPE and EE
  • Spontaneous reaction: CPE -> EE (galvanic type redox reaction)
• Most common fuel cells
• Ms Holland likes anode on the left=D
• Pretty efficient, but it is a bit inefficient too, so some energy is converted to heat.
• Structure: Electrolyte (3 types of Hydrogen Oxygen Cells)
  • Acidic Kind:
    • H2SO4 sulphuric acid
      • good for labs
      • not good for industries
    • phosphoric acid
      • good for industries
  • Alkaline:
    • KOH (potassium hydroxide)
      • both lab and industries
  • PEM (Proton Exchange Membrane)
  • minor intricate differences between all 3
• Porous Carbon
  • Tiny Holes
  • +Catalyst (sometimes)
    • changes depending on version
• Anode, Cathode Compartment
• Electrical Energy comes from the movement of electrons.
• Increase voltage by connecting batteries in series.
• Advantages
  • Could replace power stations, main idea is to replace fossil fuel engines and batteries, which are both pollution.
  • Only requires Hydrogen and Oxygen, and does not produce Carbon Dioxide or pollutants.
  • Last longer than batteries, less polluting to dispose of.
• Disadvantages
  • Hydrogen is a gas
    • hard to fuel up, high flammable and not poggers.
  • more space to store than fossil fuels or batteries
  • Explosive when mixed with air
    • will explode in contact with air.
  • storage is dangerous
  • BIGGEST ISSUE: Hydrogen gas requires energy, and its often produced using fossil fuels.
• aqueous acidic hydrogen oxygen fuel cell: has extra thing at anode for depleted fuel and other gases to come out: catalyst provides conductive surface for thing to occur, platinum doesn’t corrode (expensive for jewellery)
• proton exchange membrane works similar to aqueous acidic.
• production of hydrogen
  • hydrogen is a very reactive element
    • reaction of hydrocarbons with steam
    • electrolysis of acidified water
• TRANSPORTING HYDROGEN
  • liquify under pressure
    • very hard tho, you can force it into liquid state with high pressure, but pressure is ridiculously low.
    • expensive method of storage
  • adsorption
    • take place at low pressure and close to room temperature
    • however, materials being used for adsorption does not last forever.
  • absorption
    • absorbed by solid materials, enter spaces in metal alloys lattices, forming hydrides
    • alloys deteriorate over time, and need to be replaced regularly
• hydrogen fuel bus
  • since transporting is difficult, some vehicles makes hydrogen using hydrogen-rich fuels and then using it with fuel cells.
• direct methanol fuel cells
  • 
  • anode: ch3oh + h2o -> 6h+ +6e- + co2
  • cathode: 3/2 o2 + 6h+ + 6e- -> 3h2o
  • redox: ch3oh + 3/2 o2 -> co2 + 2h2o
• making ethanol
  • less toxic more energy dense alcohol than methanol.
  • not ready, as catalysts are not cheap !! very expensive $$ ahhh
  • cheaper, more efficient catalysts capable of fully oxidising ethanol are needed.
• pros cons fuel cells
  • powerpoint slide on ms holland fuel cells ppt, put the slide here LATER AT SOME POINT ! THANKS!!!!
• breath alcohol tester.
  • fuel cell technology used to breath alcohol testers.

Electrolysis of molten compounds §

"im dababy" - ms holland

"probably saw a year 8" - ms holland referring to arvin

• non-spontaneous redox reactions
• similarities and differences aqueous and molten
• industrial application electro-refining, electroplating: big in WACE (and probably in pmod as well.)
• watched GCSE Chemistry Part 1

What is electrolysis?

the process of using an external potential difference to supply electrical energy to enable a non-spontaneous reaction to occur.

• electricity to make non-spontaneous redox reactions occur
• many uses
  • purifying metals (such as copper)
  • plating metals (silver or gold)
  • extracting reactive metals (such as aluminium)
  • making chlorine, hydrogen (not cheapest way, but byproduct of making chlorine) and sodium hydroxide (how most sodium hydroxide is made)
• electrolyte needs to be liquid: it doesn’t have to be solution, and can be molten instead.
  • as long as ions are mobile (free to move), it is O.K.
  • carbon electrodes pretty popular in molten aqueous: graphite
    • as its pretty inexpensive (? not sure if she said this) and it can withstand pretty high temperature.
• battery pushes electrons and forces in non-spontaneous direction.
• “this gives the theoretical minimum potential difference that must be applied for the reaction to occur”
  • in practice, a larger potential difference is required, to overcome energy loss in the system.

electrolysis extracting §

• we dig up aluminium oxide, and turn it into pure aluminium
• aluminium is very useful
• aluminium ore (bauxite) has very high melting point.
• the ore is dissolved in compound called cryolite, which lowers the melting point to 700.
• inert carbon lining steel container.
• aluminium gets reduced, sinking to bottom, oxygen oxidises into gas.
• aluminium oxide -> aluminium + oxygen gas$2A{l}_2{O}_3(l){\to }_{4}Al(l)+3O_2(g)$
• oxygen produced at the anode will react with the carbon, producing co2, anodes are replaced periodically.

environmental impact §

• starting off with aluminium oxide, has to mine it
  • mine it is bad for environment when you dig a big ass hole :3
  • environmental ecological impact on ecosystems and animals and plants
  • energy consumption of diesel and coal and stuff in the process of mining BAD BAD BAD !!
• processing is expensive so we dont do in australia
  • so we ship it across the world
  • energy consumption from fuel, crude oil, fossil fuels.
• go to china! cheap work yay child labour
  • china cheap energy, cheapest coal, brown coal, low energy density and filthy!!! yucky!!!
  • aluminium gets made there ^.^
  • big cities in china pollutant because brown coal is getting burnt. carbon dioxide hooray!
• aluminium gets shipped back.
  • AHH!! scary
• melting bauxite needs energy (comes from brown coal)
• electrolysis requires energy (requires brown coal)
• massive energy demand and massive environmental impact!! =D
• recycling is good, making new aluminium is not poggers for the trees.
  • recycle !!! - ms holland

Electrolysis of Solutions §

• In theory removing the need to melt, but in practice, adding another compound to melt.
• Aqueous solution is used, allowing the process to take place at lower temperatures.
  • However, presence of water affect outcome of some electrolytic reactions.
• Instead of molten compound, you have a solution, and provides a potential difference.
• thought process
  • which substances are present in the cell
    • which substances can be oxidised
    • which substances can be reduced
  • which substances are the most energetically favourable at each electrode? (process for least energy overall likely most favourable, use SRP to help)
  • which substances are present in higher concentrations? substances present at very low concentrations are less likely to produce bulk of product.
• if there is two possible oxidation reactions, and the provided voltage allows both to occur, the two possible reactions happens simultaneously.
• layer of solid copper forms at cathode =D

Electrorefining: Reactive Electrodes §

Reactive Electrodes §

• you can replace the inert electrodes with reactive electrodes.
• swapping out the electrodes for copper!
• Cu (s) -> Cu (s) if copper is the most reactive thingy. E cell would = 0, and what will happen is that copper will move from the anode to cathode.
• potential difference applied needs to be >= 0

Electrorefining / Electrowinning §

• Impure metal is the anode
• Highly pure metal as cathode
• Anodic Slime
  • can be very valuable because sometimes gold is there :3

Electroplating §

• higher current, higher rate of electroplating
  • this can be increased by increasing concentration!
• chromium forms oxide layer that utilises passivation to protect the metal underneath.
• oxygen and cl2 in solution.
• products of sodium chloride solution produces three very useful products
  • chlorine
  • hydrogen (without use of unsustainable means)
  • seawater; financially viable
• chlorine is expected as product but hydrogen and NaOH is a fun surprising result.
• if ionic compound containing mental has negative SRP< electrolysis of solution of compound will usually produce hydrogen rather than the metal.
• if a metal has negative srp - it will struggle to be reduced (?) you get hydrogen instead sometimes

brine electrolysis: chloralkali process §

• solution becomes basic and you make sodium hydroxide solution!
• most energetically favourable reactions will occur for reduction/oxidation
• ion only membrane : allows ion through, but not gases! this stops the reactions that is spontaneous for electroplating.
• concentration sodium chloride solution
  • dont do competing reactions !!!
• chloride ions being used flow through.
• electrolysis can be used to split h2o into tis elements, hydrogen and oxygen.
  • add something to boost its conductivity; simple solution is add sulphuric acid1
• hydroxides -> oxygen water is an option if there are hydroxides
• +1.23 O2 + 4h -> 2h2o
• electrolysis of water
  • at the negative electrode:
    • 2h+ + 2e- -> h2 (reduction)
  • at the positive electrode:
    • 4oh- -> 2h2o + o2 + 4e (oxidation)

CAP THINGS §

• check resources of lucarelli 7.6-7.8, 8.6-8.8
• read pearson chapter 10, industrial examples very useful in pearson, more familiar, less likely youd get caught by it questions:
• lucarelli set 11 & set 12 (q16 onwards) HAND IT NEXT WEEK
• stawa sets 22 & 23
• DIRECTLY PREPARE YOU!!!!!!!!!!!!!!!!!!! FOR THE CAP!!!!! DUE WEDNESDAY 21ST!!
• ms holland %
  • as you go along, you mark it
  • give it a % for the piece of work
• often examiners start short answer with things to shift modes and let you write something immediately.