IB Chemistry HL > UNIT 9 - Redox Processes > Flashcards

UNIT 9 - Redox Processes Flashcards

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1
Q

Define SHE and state its function.

A
  • standard hydrogen electrode
  • the reference point for measuring and comparing the electrode potentials of other half-cells with standard electron potential of zero volts.
  • it performs a redox reaction of protons at a platinized platinum electrode dipped in an acidic solution and pure hydrogen gas bubbled through it.
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2
Q

State the connections between half-cells required for a voltaic cell.

A
  • external electronic circuit connected to the metal electrode of each half-cell through which the electrons will flow.
  • salt bridge containing an aqueous solution of ions maintaining the potential difference and neutralises any build-up of charge.
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3
Q

Define the standard electrode potential of a half-cell.

A

EMF generated when it is connected to the standard hydrogen electrode by an external circuit and a salt bridge measured under the standard conditions.

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3
Q

Define an oxidizing agent.

A

reactant accepting electrons because it brings out the oxidation in the other reactant and it itself becomes reduced in the process.

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3
Q

State the role of a voltaic cell.

A

Conversion of the energy released from a spontaneous exothermic reaction into electrical energy.

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3
Q

State the charge of the anode and cathode in an electrolytic cell.

A

Anode (+)
Cathode (-)

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3
Q

State the full name and define BOD.

A

Biological oxygen demand.

Measures the degree of pollution as it is the amount of oxygen used to decompose the organic matter in a sample of water over a specified time period (usually five day at a specified temp).

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3
Q

State the reason for adding the starch indicator to titrate the iodine redox reaction.

A

The starch serves as an indicator as it forms a deep blue colour by forming a complex with free I2 as the I3- ions embed themselves within the helix of amylose as the I2 is reduced to I- during the reaction as the blue colour disappears marking the equivalence point.

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3
Q

State and describe the method used to calculate BOD.

A

The Wrinkler method.

  1. The dissolved oxygen, O2 in the water is ‘fixed’ by the addition of a maganese (II) salt such as MnSO4. Reaction of this salt with O2 in basic solution causes oxidation of MN(II) to higher oxidation states, such as Mn(IV).

2Mn2+ + O2 + 4OH- -> 2MnO2 + 2H2O

  1. Acidified iodide ions, I- are added to the solution and are oxidized by the Mn(IV) to I2.

MnO2 + 2I- +4H+ -> Mn2+ + I2 + 2H2O

  1. The iodine produced is then titrated with sodium thiosulfate.

2S2O3 2- + I2 -> 2I- + S4O6 2-

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3
Q

Define galvanized iron and state its function.

A

Iron with a layer of zinc deposited on its surface will be protected from corrosion as the zinc will be preferentially oxidized.

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3
Q

Ratio for water and S2O3 2- during the Wrinkler method.

A

1 mol of O2 - 4 mol of S2O3 2-

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4
Q

List the components of an electrolytic cell.

A
  • battery or a DC power source as the source of electric power
  • electrodes immersed in the electrolyte and connected to power supply not touching each other
  • electrodes made from a nonconducting substance - metal or graphite
  • electric wires connect the electrodes to the power supply
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4
Q

State the relationship between the cell potential and spontaneity of a reaction.

A
  • If E-cell is positive, the reaction is spontaneous.
  • If E-cell is negative, the reaction is non-spontaneous BUT the reverse reaction is spontaneous.
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5
Q

Define electrode potential.

A

Charge separation between the metal and its ions in the solution caused by the metal half-cell in which the atoms will form ions by releasing electrons that will make the surface of the metal negatively charged with respect to the solution.

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6
Q

State the alternative name for NaCl (aq).

A

Brine

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7
Q

State when are the electrolytic cell electrodes described as inert.

A

When they do not take part in the redox reactions.

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8
Q

Define reduction.

A

Gain of electrons.

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9
Q

State the factor determining the direction of electron flow and voltage generated by the voltaic cell.

A

The difference in reducing strength of the two metals judged by their relative position in the reactivity series.

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10
Q

Define and state the full name of EMF.

A

the electromotive force

Greatest potential difference that a cell can generate and is measurable only when the cell is not supplying current due to its internal resistance measured in volts.

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11
Q

State the relationship between the electrode potential in a half-cell and its ability to be a cathode or anode.

A

The half-cell with the higher electrode potential is the cathode.
The one with the lower electrode potential is the anode.

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12
Q

State the relationship between the reactivity of a metal and its potency as the reducing agent.

A

More reactive metals loose their electrons more readily and thus are stronger reducing agents.

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13
Q

Define electroplating.

A

Using electrolysis to deposit a layer of metal on top of another metal or other conductive object.

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14
Q

State what process occurs at the anode and what charge does it thus have.

A

oxidation
negative charge

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14
Q

State when is a half-cell given the symbol E (with the little o) and the meaning of its value.

A

When the SHE is connected to another half-cell by an external circuit with a high-resistance voltmeter and a salt bridge and the EMF generated is then know as the standard electrode potential of that half cell.

POSITIVE value of E = greater tendency to be reduced than H+ (electrons flow from the hydrogen half-cell to the other one which is reduced).

NEGATIVE value of E = less tendency to be reduced than H+ (electrons flow from the other half-cell to the hydrogen half-cell which is reduced).

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15
State how are voltaic cells constructed.
As the process relies on separating the two half reactions the process relies on half-cells allowing the electrons to flow between them only through an external circuit.
15
State when does reduction occur in terms of oxidation states.
When there is a decrease in oxidation state of an element.
16
State the reason why platinum is used as the conducting metal in the SHE.
- fairly inert and wont ionize - catalyst for the reaction of proton reduction
17
State the principles of a cell diagram convention.
- single vertical line - phase boundary - double vertical line - salt bridge - aqueous solutions placed next to the salt bridge - anode on the left, cathode on the right, electrons flow from left to right
18
List the purpose of electroplating.
- decorative purposes - corrosion control - sacrificial protection - improvement of function (e.g. chromium improves the wear on steel parts)
19
State the standard conditions used for measuring standard electrode potentials.
- solutions with concentration of 1.0 M - gases at pressure 100 kPa - all substances must be pure - temperature is 298K - platinum used as electrode if the half-cell doesn't include a solid metal
20
State how water can be both oxidized and reduced.
at the CATHODE can be reduced to H2: 2H2O + 2e -> H2 +2OH- at the ANODE can be oxidized to O2: 2H2O -> 4H+ + O2 + 4e
21
State the location of the redox reactions in an electrolytic cell.
At the electrodes as the electric current passes through the electrolyte.
22
State the relationship between the half- cell potential and oxidizing potential of an agent.
The higher the value of half-cell potential the stronger the oxidizing agent and thus the lower the value of standard half-cell potential the stronger the reducing agent.
23
Define an electrode.
A half-cell connected to another half-cell through an external part through which the flow of electrons is enabled.
24
State the features of an electrolytic cell used for electroplating.
- electrolyte containing the metal ions which are to be deposited - the cathode made of the object to be plated - sometimes the anode is made of the same metal which is to be coated because it may be oxidized to replenish the supply of ions in the electrolyte.
25
State the effect of the electrolysis of the molten salts.
The only ions present are those from the compound itself as there is no solvent thus only one ion migrate to each electrode.
26
State the relationship between the reactivity of a non-metal and its potency as the oxidizing agent.
More reactive non-metals are stronger oxidizing agents than less reactive non-metals.
27
State the two types of electrochemical cells.
- voltaic (galvanic) cells - electrolytic cells
28
Distinguish the voltaic cells from electrolytic cells.
Voltaic cells generate electricity from chemical reactions. Electrolytic cells drive chemical reactions using electrical energy.
29
State the factors influencing the amount of product formed in electrolysis.
- current - duration - the charge on the ion
30
State the effect of redox reactions in an electrolytic cell.
Removing the charges of ions (discharging them) and forming products that are electrically neutral.
31
Describe the analysis of iron with manganate (VII) taking into account the colour change.
5Fe2+ + MnO4- + 8H+ -> 5Fe3+ + Mn2+ + 4H20 MnO4- = purple Mn2+ = colourless oxidizing agent: KMnO4 oxidizes Fe2+ ions to Fe3+ MnO4- is reduced to Mn2+
32
State the net effect of reduction on the number of oxygen and hydrogen atoms.
loss of oxygen gain of hydrogen
33
State the factors on which the outcome of a redox reaction in an aqueous solution depend on.
- relative cell potential values of the ions - relative concentrations of the ions in the electrolyte - nature of the electrode
34
State the requirement needed for a metal to be able to reduce ions of another metal.
Only a more reactive metal is able to reduce the ions of a less reactive metal.
35
State the factor on which the magnitude of the cell potential depends on.
The difference in the tendencies of the two half-cells to undergo reduction.
36
State and define the SI unit of electric charge (Q).
coulomb (C) amount of charge transported in 1 second by a current of 1 ampere
37
State the relationship between the reactivity of a metal and its electrode potential.
The more reactive a metal the more negative its electrode potential in its half cell.
37
State the mechanism occurring in the electrolytic cell.
1. power source pushes electrons towardas the negative electrode 2. they enter the electrolyte - cathode 3. electrons are released at the positive terminal - the anode - and returned to the source 4. current passed through the electrolyte by the mobile ions migrating to the electrodes 5. chemical reactions occurring at each electrode remove the ions from the solution and enables the process to continue
38
Define an oxidation state.
The value we assign to each atom in a compound which is a measure of the electron control/possession it has relative to an atom in the pure element.
39
State what process occurs at the cathode and what charge does it thus have.
reduction positive charge
40
State the factor on which the electrode potential is dependent on and one that determines the size of it.
Dependent on the reactivity of a metal Position of the equilibrium between the ions formed in the solution and the electrons on the metal strip.
41
State the net effect of oxidation on the number of oxygen and hydrogen atoms.
gain of oxygen loss of hydrogen
42
State the factors that the electrode potentials depend on.
- concentrations of ions - gas pressures - purity of a substance - temperature
43
State the directionality of the flow of electrons through the external electronic circuit.
From the anode to the cathode.
43
State a simple way of making a half-cell.
Putting a strip of metal into solution of its ions.
44
State the formula for quantifying the electric charge.
charge = current x time
44
State the directionality of the flow of electrons through the salt bridge.
From the anode to the cathode.
45
State the environmental importance of calculation the dissolved oxygen content.
The dissolved oxygen content of water is one of the most important indicators of its quality as oxygen is essential of the survival of aquatic life. Level of pollution increases = dissolved oxygen content decreases as it is used by bacteria in decomposition reactions.
45
State the conductivity of electrodes and electrolytes.
- electrodes are electronically conducting - electrolytes are ionically conducting
46
State the redox reaction occurring at the anode of the electrolyte cell.
At the positive electrode (anode): A- -> A + e- anions lose electrons so are oxidized.
46
Define selective discharge.
Discharge of an ion at the electrode when more than one redox reaction is possible at each electrode.
47
State the SI unit of electric current (I).
ampere - amp [A]
47
State the redox reaction occurring at the cathode of the electrolyte cell.
At the negative electrode (cathode): M+ + e- -> M cations gains electrons so are reduced.
48
State the directionality of the electron flow in relation to the electrode potential.
Electrons always flow towards the half-cell with the highest electrode potential.
48
Define oxidation.
loss of electrons
49
State and define the SI unit of potential difference.
volt (V) amount of energy in J that can be delivered by a coulomb of electric charge V=JxC-1 Equal to the difference in electric potential between the two points on a conducting wire.
49
Define a redox reaction.
A chemical reaction in which changes in the oxidation states occur.
50
State when does oxidation occur in terms of oxidation states.
When there is an increase in oxidation state of an element.
50
Define a reducing agent.
reactant suppling electrons because it brings out the reduction in the other reactant and it itself becomes oxidized in the process.
51
Define cell potential (Ecell)
The potential difference that a cell can generate when generating electrical energy.
52
State the underlying mechanism of functioning of an electrolytic cell.
Uses an external source of electrical energy to bring about a redox reaction that would otherwise be non-spontaneous.
53
State the formula for the standard potential of a cell.
E cell = E cathode - E anode
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