redox & electrode potential °‧⭑.ᐟ Flashcards
1
Q
what are reducing agents in redox reactions?
A
electron donors
2
Q
what are oxidising agents in redox reactions?
A
electron acceptors
3
Q
what does oxidation involve in terms of electrons?
A
loss of electrons
4
Q
what is the oxidation reaction of zinc?
A
Zn → Zn²⁺ + 2e⁻
5
Q
what does reduction involve in terms of electrons?
A
gain of electrons
6
Q
what is the reduction reaction of chlorine?
A
Cl₂ + 2e⁻ → 2Cl⁻
7
Q
in the reaction Br₂ +2I⁻ → I₂ + 2Br⁻, which species is oxidised?
A
iodide ion
8
Q
in the reaction Br₂ +2I⁻ → I₂ + 2Br⁻, which species is reduced?
A
bromine
9
Q
how should oxidising and reducing agents be named?
A
full name of substance
10
Q
what is the oxidising agent in the reaction involving Br₂ (aq) + 2e⁻ → 2Br⁻?
A
bromine
11
Q
what is the reducing agent in the reaction involving 2I⁻ (aq) → I₂ (aq) + 2e⁻?
A
iodide ion
12
Q
what is an oxidising agent?
A
species that causes another to oxidise
13
Q
what is a reducing agent?
A
species that causes another to reduce
14
Q
species that causes another to reduce
A
on the left
15
Q
where are electrons placed in an oxidation half equation?
A
on the right
16
Q
how do you balance half equations with varying oxygen amounts?
A
add OH⁻, H⁺ and H₂O
17
Q
what is the half equation for MnO₄⁻ to Mn²⁺ in acidic conditions?
A
MnO₄⁻ + 8H⁺ + 5e⁻ → Mn²⁺ + H₂O
18
Q
what is the full balanced equation for the reduction of MnO₄⁻ and oxidation of C₂O₄²⁻?
A
2MnO₄⁻ + 16H⁺ + 5C₂O₄²⁻ → 2Mn²⁺ + 10CO₂ + 8H₂O
19
Q
what is the half equation for the change of SO₄²⁻ to SO₂
A
SO₄²⁻ + 4H⁺ + 2e⁻ → SO₂ + 2H₂O
20
Q
what is the common exercise in manganate redox titration?
A
redox titration between Fe²⁺ and MnO₄⁻
21
Q
why is the manganate titration considered self-indicating?
A
significant colour change from purple to colourless
22
Q
what is the role of acid in manganate titrations?
A
to supply 8H⁺ ions
23
Q
why should only dilute sulfuric acid be used in manganate titrations?
A
other acids can cause inaccurate titration readings
24
Q
what is the consequence of using nitric acid in manganate titrations?
A
oxidizes Fe²⁺ to Fe³⁺
25
how does the purple colour of manganate affect titration readings?
makes it hard to see the meniscus
26
what indicates the end point of the manganate titration?
first permanent pink colour appears
27
what is the common exercise in thiosulfate redox titration?
redox titration between I₂ and thiosulfate
28
what colour change occurs when starch is added near the end point of thiosulfate titration?
from blue/black to colourless
29
what are the steps in the manganate redox titration process?
- add dilute sulfuric acid to the solution
- titrate with MnO₄⁻ until the first permanent pink colour appears
- record the volume of MnO₄⁻ used
30
what are the steps in the thiosulfate redox titration process?
- add starch indicator near the end point
- titrate with thiosulfate until the iodine fades
- observe the colour change from blue/black to colourless
31
what is the balanced equation for the reaction in manganate titration?
MnO₄⁻ + 8H⁺ + 5Fe²⁺ → Mn²⁺ + 4H₂O + 5Fe³⁺
32
what happens if insufficient sulfuric acid is used?
- solution is not acidic enough
- MnO₂ will be produced instead of Mn²⁺
33
what is the effect of brown MnO₂ in the titration?
masks the colour change
34
why can't weak acids like ethanoic acid be used in manganate titrations?
cannot supply large amount of hydrogen ions needed
35
why is concentrated HCl unsuitable for manganate titrations?
Cl⁻ ions would be oxidisied to Cl₂ by MnO₄⁻
36
what is the balanced equation for the reaction between thiosulfate and iodine?
2S₂O₃²⁻ (aq) + I₂ (aq) → 2I⁻ (aq) + S₄O₆²⁻(aq)
37
what are the strengths and weaknesses of using different acids in manganate titrations?
strengths
- dilute sulfuric acid provides sufficient H⁺ ions
weaknesses
- other acids can lead to inaccurate readings
- concentrated HCl produces Cl⁻
- nitric acid oxidises Fe²⁺
38
why is heating the conical flask to 60°C beneficial in manganate titrations?
it speeds up the initial reaction
39
how do you calculate the percentage of iron by mass in a nail sample?
% mass = (mass of Fe / mass of nail) × 100
40
what is the role of a salt bridge in an electrochemical cell?
connects the half-cells and conducts charge
41
what is the typical composition of a salt bridge?
filter paper soaked in potassium nitrate
42
why is a high resistance voltmeter used in measuring electrode potential?
prevent current flow and measure maximum potential
43
what happens to the voltage when current flows in an electrochemical cell?
voltage falls to zero as reactants are used up
44
what is the potential difference measured in an electrochemical cell?
difference between two electrodes' potentials
45
what is the standard potential of the hydrogen electrode?
0 volts
46
what is the overall reaction for manganate titration with hydrogen peroxide?
2MnO₄⁻ (aq) + 6H⁺ (aq) + 5H₂O₂ → 5O₂ + 2Mn²⁺ (aq) + 8H₂O
47
how does the reaction between MnO₄⁻ and C₂O₄²⁻ proceed overall?
2MnO₄⁻ (aq) + 16H⁺ (aq) + 5C₂O₄²⁻ (aq) → 10CO₂ (g) + 2Mn²⁺ (aq) + 8H₂O (l)
48
how do you calculate the moles used in a titration?
moles = concentration x volume
49
what is the standard hydrogen electrode (SHE) assigned potential?
0 volts
50
what is the equilibrium reaction for the hydrogen electrode?
H₂ (g) ⇌ 2H⁺ (aq) + 2e⁻
51
why is a platinum wire used in the SHE?
acts as a catalyst
absorbs hydrogen gas
52
what are the conditions for a standard hydrogen electrode?
- hydrogen gas at 100 kPa
- H⁺ concentration at 1 mol dm⁻³
- temperature at 298 K
53
what is a secondary standard in electrochemistry?
a standard calibrated against the SHE
54
why are standard conditions necessary for measuring electrode potentials?
ensure consistent redox equilibrium positions
55
what are the standard conditions for measuring electrode potentials?
1 mol dm⁻³ ions, 298 K and 100 kPa
56
what is the standard electrode potential measured against?
standard hydrogen electrode
57
how is the standard electrode potential quoted in data books?
half equations or cell notation
58
what does a more negative half-cell potential indicate?
it will oxidise and go backwards
59
how do you calculate the cell potential (Ecell) for a spontaneous reaction?
Ecell = E(red) – E(ox)
60
what does a positive Ecell indicate about a reaction?
reaction is feasible
61
what is the half-equation for zinc oxidation?
Zn (s) → Zn²⁺ (aq) + 2e⁻
62
what happens to electrons at the positive electrode during a redox reaction?
electrons are absorbed (gained)
63
what is the trend in electrode potentials for reducing and oxidising agents?
- strong reducing agents: most negative potentials
- strong oxidising agents: most positive potentials
64
how do you determine the strongest oxidizing agent from a series of potentials?
look for the most positive potential
65
how do you determine the strongest reducing agent from a series of potentials?
look for the most negative potential
66
what is the half-equation for lithium reduction?
Li⁺ + e⁻ → Li
67
what is the half-equation for fluorine reduction?
F₂ + 2e⁻ → 2F⁻
68
what does an increasing positive potential indicate about a species?
has a greater tendency to reduce
69
what does an increasing negative potential indicate about a species?
has a greater tendency to oxidise
70
how do you calculate Ecell from two standard electrode potentials?
Ecell = E(red) – E(ox)
71
what is the significance of E° values in redox reactions?
higher E° indicates stronger oxidising agents
lower E° indicates stronger reducing agents
72
what is the difference in voltage behaviour between fuel cells and ordinary cells?
fuel cells maintain constant voltage over time
73
what type of cells are nickel-cadmium cells?
rechargeable cells
74
what is the overall reaction during the discharge of nickel-cadmium cells?
2NiO(OH) + Cd + 2H₂O → 2Ni(OH)₂ + Cd(OH)₂
75
what characterizes non-rechargeable cells?
reactions within them are non-reversible
76
what types of electrochemical cells exist?
- non-rechargeable (irreversible) cells
- rechargeable cells
- fuel cells
77
what is e.m.f?
a measure of how far from equilibrium the cell reaction lies
78
what does it mean if the e.m.f is more positive?
more likely the reaction is to occur
79
what would happen if the concentration of the reactants were to increase?
it would increase the e.m.f
80
what would happen if the concentration of the reactants were to decrease?
it would decrease the e.m.f
81
are most Ecells endothermic or exothermic?
exothermic
82
what would the effect of a temperature rise have upon the Ecell?
decrease in Ecell
83
if the reaction has a high activation energy will it occur?
no
84
if the Ecell positive and it indicates a reaction might occur, what possibility may occur?
reaction will not occur or will occur so slowly that effectively it doesn’t happen
85
why is an uncatalysed reaction very slow?
- reaction needs a collision between two negative ions
- repulsion between the ions is going to hinder this
- high activation energy
86
how can a substance act as a homogenous catalyst?
- electrode potential must lie in between the electrode
potentials of the two reactants
- it can first reduce the
reactant with the more positive electrode potential
- then in the second step oxidise the reactant with the more
negative electrode potential
87
what do both individual stages in the catalysed mechanism involve?
- collision between positive and negative ions
- will have lower activation energies
88
what does using the E values to find a catalyst only show?
catalysts is possible
does not guarantee that rate of reaction will increase
89
what are electrochemical cells?
commercial source of electrical energy
90
what are the three types of cells?
- non-rechargeable
- rechargeable
- fuel cells
91
when are cells non-rechargeable?
reactions that occur with in them are non-reversible
92
what are nickel-cadmium cells used for?
power electrical equipment such as drills and shavers
93
what type of cells are nickel-cadmium cells?
rechargeable
94
what are fuel cells?
uses the energy from the reaction of a fuel with oxygen to create a voltage
95
what is the hydrogen fuel cell overall reaction?
2H₂ + O₂ → 2H₂O
96
what are the hydrogen fuel cell half equations in acidic conditions?
2e⁻ + 2H⁺ → H₂
4e⁻ + 4H⁺ +O₂ → 2H₂O
97
what are the hydrogen fuel cell half equations?
4e⁻ + 4H₂O → 2H₂ + 4OH⁻
4e⁻ + 2H₂O + O₂ → 4OH⁻
98
how can fuel cells will maintain a constant voltage over time?
continuously fed with fresh O₂ and H₂ so maintaining constant
concentration of reactants
99
how do fuel cells differ from normal cells in terms of voltage?
ordinary cells' voltage drops over time as the reactant concentrations drop
100
why can't standard conditions be used in a hydrogen fuel cell?
rate is too slow to produce an appreciable current
101
why are higher temperatures used in hydrogen fuel cells?
increase rate but the reaction is exothermic
applying le chatelier would mean the emf falls
higher pressure can help counteract this
102
advantages of fuel cells
- less pollution and less CO2
(pure hydrogen emits
only water whilst hydrogen-rich fuels produce only
small amounts of air pollutants and CO2)
- greater efficiency
103
disadvantages of hydrogen fuel cells
- storing and transporting hydrogen, in terms of safety,
feasibility of a pressurised liquid and a limited life cycle of a
solid ‘adsorber’ or ‘absorber’
- limited lifetime (requiring regular replacement and
disposal) and high production costs
- use of toxic chemicals in their production
