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Chemistry IB HL > Topic 9: Redox Processes > Flashcards

Topic 9: Redox Processes Flashcards

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1
Q

oxidation

A
  • loss of electrons
  • increase in oxidation no
  • gain of oxygen
  • loss of hydrogen
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2
Q

reduction

A
  • gain of electrons
  • decrease in oxidation no
  • loss of oxygen
  • gain of hydrogen
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3
Q

oxidation state

A

apparent charge of an atom in a molecule or ion.

  • used to measure electron control or possession relative to the atom in pure element
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4
Q

relation between reducing power and reactivity in metals

A

more reactive metals are stronger reductants

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5
Q

relation between oxidizing power and reactivity in non-metals

A

more reactive non-metals are stronger oxidizing agents

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6
Q

redox titration

A

redox reaction between an oxidizing agent and reducing agent

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7
Q

chemically speaking, what happens in a redox titration?

A

electrons are transferred from reductant to oxidant

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8
Q

redox titrations: analysis of iron with manganate (VII)

A

5Fe2+ + MnO4- + 8H+ -> 5Fe3+ + Mn2+ + 4H2O

  • uses KMnO4 in acidic soln. as oxidant
  • oxidises Fe (II) ions to Fe (III)
  • MnO4- is reduced to Mn2+
  • colour change: deep purple to colourless
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9
Q

redox titrations: iodine-thiosulfate reaction

A

2I- (aq) + oxidant -> I2 (aq) + reduced product

  • oxidant reacts with excess iodides to form iodine diatomic molecules
  • oxidants can be KMnO4, KIO3, K2Cr2O7, NaOCl, etc.

2S2O3 2- (aq) + I2 (aq) -> 2I- (aq) + S4O6 2- (aq)

  • the liberated iodine is then titrated with sodium thiosulfate (Na2S2O3)
  • using starch as an indicator (NOT added at the start but during titration
  • initially forms deep blue colour due to starch, but as I2 is reduced to I-, the blue colour disappears
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10
Q

redox titrations: winkler method (what it is, its function, etc)

A
  • calculates dissolved oxygen content of water
  • used to measure degree of pollution
  • as oxygen is used by bacteria in decomposition reactions
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11
Q

biological oxygen demand (BOD)

A
  • amount of oxygen used to decompose organic matter in a sample of water over a specified time period
  • usually 5 days at a specified temp
  • high BOD = high quantity of degradable organic waste = low level of dissolved oxygen
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12
Q

redox titrations: winkler method (process)

A
  1. Dissolved oxygen O2 (g) is fixed by the addition of a manganese (II) salt, such as MnSO4. This causes oxidation of Mn (II) to higher oxidation states:
    2Mn2+ (aq) + O2 (g) + 4OH- (aq) -> 2MnO2 (s) + 2H2O (l)
  2. Acidified iodide ions (I-) are added and are oxidised by Mn (IV) to I2:
    MnO2 (s) + 2I- (aq) + 4H+ (aq) -> Mn2+ (aq) + I2 (aq) + 2H2O (l)
  3. Iodine produced is titrated with sodium thiosulfate:
    2S2O3 2- (aq) + I2 (aq) -> 2I- (aq) + S4O6 2- (aq)
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13
Q

redox titrations: winkler method (ratio of O2 : S2O3 2-)

A

1:4

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14
Q

voltaic cells

A
  • generates electricity from spontaneous redox reactions
  • separates two half-reactions into half-cells, allowing e-s to flow between them through an external circuit

2 connected half-cells = 1 voltaic cell

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15
Q

electrode potential

A

charge separation between the metal and its ions in solution

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16
Q

half cell

A
  • where a half-reaction occurs

- simplest one is made by putting a strip of metal into a solution of its ions

17
Q

relationship between cell equilibrium and reactivity (of metal)

A
  • more reactive metals are stronger reducing agents
  • they have a higher tendency to lose electrons
  • the less reactive the metal, the more to the right the equilibrium position will be
18
Q

voltaic cell connections between half-cells

A
  • external electronic circuit

- salt bridge

19
Q

external electronic circuit

A
  • connected to the metal electrode in each half-cell
  • can have a voltmeter attached to record generated voltage
  • electrons flow from anode to cathode through wire
20
Q

salt bridge

A
  • glass tube/strip of absorptive paper containing an aq soln of ions
  • movement of these ions neutralise build-up of charge and maintains potential difference
  • anions move in salt bridge from cathode to anode, which opposes the flow of e-s in external circuit
  • cations move in salt bridge from anode to cathode
  • solution chosen is typically aq NaNO3 or KNO3
21
Q

what is affected by the difference in reducing strength between electrodes in a voltaic cell?

A
  • direction of electron flow (e-s always flow to the electrode with less reducing power)
  • voltage generated (the greater the reducing power difference, the greater the voltage generated)
22
Q

electromotive force

A
  • potential difference
  • generated due to e- flow between half-cells
  • depends on difference of reducing power
  • also called cell potential/electrode potential
23
Q

standard hydrogen electrode

A

reference standard for measuring reducing power of half-cells

24
Q

standard hydrogen electrode

A
  • modified form of pH electrode
  • platinum is used as the conducting metal in the electrode
  • as platinum is inert and doesn’t ionise
  • it also acts as a catalyst for proton reduction reaction
  • the form of platinum used is platinized platinum (the surface of the metal is covered with very finely divided platinum, or platinum black)
  • this causes the electrode reaction to occur rapidly
  • due to the large surface area helping in the adsorption of hydrogen gas
25
Q

standard conditions for measurement of standard electrode potential

A
  • concentration: 1 mol/dm3
  • pressure: 100 kPa
  • only pure substances
  • temp: 298 K
  • platinum is used as electrode if the half-cell doesn’t include a solid metal
26
Q

standard electrode potential

A
  • denoted by Eϴ
  • potential of a redox system to lose/gain electrons when connected to the SHE (assigned a value of 0 volts) via a salt bridge and high-resistance voltmeter
27
Q

standard reduction potential

A
  • half-cells exist as equilibria
  • they can occur as either oxidation or reduction
  • std. electrode potential is always given for the reduction reaction
  • so standard electrode potential values are sometimes also called standard reduction potentials
  • they don’t depend on total no. of electrons so they’re not affected by stoichiometry
  • the more positive the Eϴ value, the more readily it is reduced
28
Q

determining spontaneity using Eϴ (cell)

A
  • if Eϴ (cell) = positive, the rxn is spontaneous and the reverse is non-spontaneous
  • a voltaic cell always runs in the direction that gives a positive Eϴ (cell) value
29
Q

electrolytic cell

A
  • uses external electrical energy source to bring about a redox rxn that would otherwise be non-spontaneous
  • redox rxns occur at electrodes
  • oxidation at anode and reduction at cathode
  • e- flow from anode to cathode
30
Q

determining products in electrolytic cells

A
  1. Identify all ions in electrolyte and determine which will migrate to which electrode
  2. Write the half-equation for the rxn at each electrode
  3. Balance e-s lost and gained and add the two half-equations
  4. Consider what changes would be observed as a result of the redox processes (e.g. colour change, precipitation of solid, gas discharge, pH change…)
31
Q

electrolysis of molten salt

A
  • only 1 ion migrates to each electrode

- as ionic compounds have v high m.pts, another compound might be added to lower m.pt

32
Q

electrolysis of aq. soln.s

A
  • cathode: H2O can be reduced to H2
  • anode: H2O can be oxidised to O2
  • difficult to predict as there is more than 1 redox rxn possible for each electrode
33
Q

selective discharge

A

the discharge of an ion at the electrode in cases where there is more than 1 possible redox rxn

34
Q

how to predict the products of electrolysis of aq. soln.s

A
  • Eϴ values of the ions
  • relative conc.s of ions in electrolyte
  • nature of electrode

Chemistry IB HL (11 decks)

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