Help
Options
Didn't know it?
click below
click below
Knew it?
click below
click below
Don't Know
Remaining cards (0)
Know
retry
shuffle
restart
0:00
Embed Code - If you would like this activity on your web page, copy the script below and paste it into your web page.
Normal Size Small Size show me how
Normal Size Small Size show me how
A2 chemistry 5.3.1
Edexcel chemistry - redox equilibria
| Question | Answer |
|---|---|
| Define oxidation | Gain of O, loss of H, loss of e- (OIL) |
| Define reduction | Loss of O, gain of H, gain of e- (RIG) |
| What is an oxidising agent? | Oxidises another species by removing electrons, agent is simultaneously reduced |
| What is a reducing agent? | Reduces another species by donating electrons, agent is simultaneously oxidised |
| What is disproportionation? | Where an element in a single compound is simultaneously oxidised and reduced |
| How do you balance simple half equations? e.g. reduction of fluorine | Balance number of atoms, then balance charges with electrons. F2 (g) + 2e- -> 2F- (aq) |
| How do you balance complex half equations? e.g. methanol -> methanoic acid? | Balance number of atoms, then O with H2O, H with H+, charges with electrons. CH3OH + H2O -> HCOOH + 4H+ + 4e- |
| What is an oxidation number? | Formal number of electrons lost or gained by an atom of an element in a compound if the compound is thought to be wholly ionic |
| What are the rules for working out oxidation numbers? | uncombined element = 0, monatomic ion = charge on ion, sum in a neutral compound = 0, sum in a polyatomic ion = charge n ion, group 1 = +1, group 2 = +2, F = -1, H = +1(not with metals), O = -2 (not with F, O-, O2-) |
| What is the number of electrons in a half equation equal to? | The change in oxidation number of that element, including stoichiometry |
| What is the increase in oxidation number in a redox reaction equal to? | The decrease in oxidation number of the element being reduced, including stoichiometry |
| What is the oxidation state of carbon in methanol and methanoic acid? | Methanol CH3OH: -2, methanoic acid HCOOH: +2 |
| How many moles of ethanol react with 1 mol of MnO4-? | 1.25 |
| What is a galvanic cell? | Harnesses the spontaneous flow of electrons to convert chemical energy into electrical energy. |
| Cu2+ + Zn <-> Zn2+ + Cu: what occurs at the anode? | Zn is oxidised, as E* is lower. Zn -> Zn2+ +2e-. Electrons are transferred to wire, anode decreases in size, anode has a negative charge relative to the solution, as the solution has a net positive charge (Zn2+ added) |
| What occurs at the cathode? | Cu2+ + 2e- -> Cu, Cu2+ is reduced, E* is higher. Electrons are trasnfered from the wire. Cathode increases in size. Cathode has a positive charge relative to the solution, as the solution has a net negative charge, Cu2+ have been removed from it |
| What is a salt bridge? | Filter paper soaked in inert SATURATED solvent-KNO3, KCl. Movement of ions completes circuit. Ions are attracted to net charge of solutions (anode= positive, so anions are attracted, cathode = negative, so cations are attracted), ions must not react in ha |
| Where do you place the half cell with the higher E*? | RHS |
| What is a standard electrode potential? | The potential difference between a half cell and the standard hydrogen electrode under standard conditions (298K, 1M, 1atm) |
| What is the standard hydrogen electrode? | Reference electrode whose electrode potential is defined as 0V, 2H+ + 2e- -> H2, H2 gas (1atm) bubbled over 1 molar HCl with a platinum electrode |
| Why do you need to use the standard hydrogen electrode? | The electric potential between an electrode and a solution of ions can't be determined; you must couple it to a reference electrode and measure the potential differnce |
| What is a disadvantage of the SHE? What alternatives are there? | Can be difficult to set up: use the Calomel electrode, a secondary standard, mercury based. Hg2Cl2 + 2e- <-> 2Hg + 2Cl-, E*= +0.27V |
| What types of half cells are there? | Metal in 1 molar solution of its own ions, gas at 1 atm in contact with 1 molar solution of its own ions +Pt electrode, ions in 2 oxidation states that react on the surface of Pt electrode (Fe3+/Fe2+), acidic solution of an oxidising agent +Pt electrode |
| What is the electrochemical series? | Lists all standard electrode potentials as reduction half equations |
| What do you know about the species at the top right corner of the electrochemical series (low to high E*)? | Very stable, weak reducing agent as it is very unlikely to be oxidised, position of equilibrium for half equation lies far to RHS |
| What do you know about the species at the bottom left corner of the electrochemical series? | Very stable, weak oxidising agent as it is very unlikely to be reduced, position of equilibrium for half equation lies far to LHS so E* is very negative |
| How can you tell that F2 is a strong oxidising agent? | E* is very positive, position of equilibrium for 0.5F2 + e- -> F- lies far to the right, F- is very stable, F2 is a strong oxidising agent as it is likely to be reduced to F- |
| Why is the E* for Li+ so low? | Li+ + e- <-> Li. Li+ is stable, position of equilibrium lies far to left, Li+ is a weak oxidising agent as it is unlikely to be reduced, while Li is a strong reducing agent as it is very likely to be oxidised |
| What is an electrochemical cell? | Consists of 2 half cells joined - high resistance voltmeter, salt bridge |
| Define: standard cell potential | The potential difference between 2 electrodes in an electrochemical cell!!!!!!! |
| How do you calculate the standard cell potential? | E* cell = E* anode + E* cathode. At cathode: E* is higher. At anode: E* is lower, so reduction half equation and sign of E* are reversed. At anode, the reduction half equation has the reactant on the RHS |
| How would you reduce a chemical in a galvanic cell? | Couple it to a half cell where E* is less positive, other cell will be oxidised |
| What is the effect of multiplying a reduction half equation by a positive constant and why? | No effect, because the units of E* are V not V/mol |
| What is the relationship between ΔStotal and E*cell? | ΔStotal = n x F x E*cell. n- total number of electrons transferred in redox reaction. F- faraday constant, 96500 C/mol |
| What is the relationship between ΔStotal and k? | ΔStotal = RlnK |
| Therefore, what is the relationship between lnk and E* cell? | lnK = (nFE*cell)/R |
| What can you tell from a reaction with a positive E*cell? | ΔStotal > 0 (directly proportional), thermodynamically feasible and spontaneous, products are thermodynamically stable wrt to reactants |
| Why might a reaction with a positive E* cell not occur? | High Ea - reactants are kinetically stable wrt products, even if thermodynamically feasible Ea may be too high for reaction to occur at RT/ without a catalyst. Non standard conditions - affect the position of equilibrium and values of E* |
| What is a hydrogen fuel cell? | H2 is produced by electrolysis of H2O/ CH4 + H2O. anode: H2 -> 2H+ + 2e-. cathode: 2H+ + 2e- + 0.5O2 -> H2O. used in space shuttles and electric cars |
| What are the advantages and disadvantages of the hydrogen fuel cell? | Advantages: H2O is a renewable resource, only produces H2O not CO2, more efficient than combustion engines, produces electricity directly. Disadvantages: would only be carbon neutral if the energy for electrolysis is produced by renewables, difficult to s |
| What is an ethanol fuel cell? | Ethanol is produced by fermentation of wheat/ sugar cane. Anode: C2H5OH + 3H2O -> 2CO2 + 12H+ + 12e-. Cathode: 2H+ + 2e- + 0.5O2 -> H2O. |
| What are the advantages and disadvantages of the ethanol fuel cell fuel cell? | Advantages: renewable resource, sustainable, releases more energy per unit mass than H2, can operate at low temperatures. Disadvantages: expensive platinum catalyst, not actually carbon neutral (transportation) |
| What is a breathalyser + list 3 types? | Measures ethanol concentration in breath to estimate ethanol concentration in blood: dichromate (VI), IR, ethanol fuel cell |
| What is a disadvantage of using breathalysers? | Not as accurate as blood tests as ethanol may be exhaled from the mouth/ nose, giving an overestimate |
| How does a dichromate breathalyser work? | Fixed volume of breath enters vial, ethanol reduces dichromate (VI) to chromium (III), photocells measure the extent of the colour change from orange to green |
| How does an ethanol fuel cell breathalyser work? | Fixed volume of breath enters fuel cell, ethanol is oxidised at anode and O2 is reduced at cathode, size of electric current depends linearly (directly proportional to) amount of ethanol present |
| How does an IR breathalyser work? | Measure IR absorption by breath, in fingerprint region of ethanol/ functional group stretches |
| How do rechargeable batteries work? | Apply an electric current to reverse the spontaneous flow of electrons, cathode -> anode |
| Give an outline of how to estimate the concentration of an oxidising agent | Prepare oxidising agent (MnO4-, Cr2O72-, Cu2+, Cl2, Br2), to 25cm3 add excess KI, titrate I2 formed with Na2S2O3 which reduces it to I- until straw yellow, add starch, blue-black, titrate Na2S2O3 dropwise until colourless, repeat for concordant titres |
| How would you prepare a sample of Cu2+? | Oxidise Cu with HNO3, then neutralise any excess HNO3 with NaHCO3 otherwise it would oxidise the I- produced in the titration back to I2 |
| When should starch be added + why? | Straw yellow. Natural colour change is brown -> colourless but starch makes it easier to see. Too early = insoluble iodine-starch complex makes end point difficult to see, too late = all I2 has already been reduced back to I- so end point has been missed |
| Write the reaction between I2 and Na2S2O3 | I2 + 2S2O32- -> S4O62- + 2I- |
| Describe the generic calculation for working out the % of copper in an alloy | Moles of S2O32- = mean titre x concentration. Moles of I2 = divide by 2. Moles of oxidising agent in 25cm3. Moles of oxidising agent in 250cm3, scale up. Mass = moles x molar mass. % copper = mass of copper/ total mass x 100 |
| What are the sources of error in an iodometric titration? | I2 may evaporate, end point is unclear, point at which starch should be added is subjective and unclear |
| Give an outline of how to estimate the concentration of a reducing agent | Prepare reducing agent (H2O2, I-, Fe2+). add 25cm3 of reducing agent + H2SO4 to a conical flask + indicator if necessary. Titrate with oxidising agent from burette until colour change. repeat for concordant titres |
| Why must H2SO4 be added before the titration? | Creates an acidic environment for oxidising agent e.g. MnO4- => Mn2+ NOT MnO2. HCl would react with oxidising agent (uses it up), Cl2, recorded titre is too large. HNO3 would oxidise reducing agent, recorded titre is too small. |
| Why is there a colour change at the end point? | Reducing agent in excess -> oxidising agent in excess |
| What is the colour change for a titration with manganate? | Colourless -> pale permanent pink (self-indicating) |
| What is the generic calculation for working out % of iron in tablets? | Moles of oxidising agent = mean titre x concentration. Moles of reducing agent in 25cm3. Scale up. Mass = molar mass x moles. % by mass = mass/ total mass x 100 |
| What are the sources of error for this titration? | Reducing agent could be oxidised by O2 in air |
| How do you calculate % error? | Max error in equipment/ value measured x 100 |
| From this, how would you reduce error? | Use larger volumes -> % error is smaller |
| What is the error of a: pipette, burette, 2 dp balance? | 0.05, 0.05 (x2), 0.005 (x2) |
| How do you work out the maximum possible mass of a product? | Mass x mass x % error |
Created by:
11043
Popular Chemistry sets