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MECH U3/4 > Redox > Flashcards

Redox Flashcards

1
Q

Define redox

A

A reaction in which electron transfer occurs from the reducing agent to the oxidising agent.

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2
Q

What occurs in a redox reaction

A

the oxidation number of one element will increased (oxidised) and the oxidation number of another element will decrease (be reduced)

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3
Q

Define Oxidation numbers

A

used to determine the change in ‘oxidation’ (change in e-)

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4
Q

What is the oxidation numbers for uncombined elements?

A

0

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5
Q

Give 2 examples of uncombined elements and their oxidation numbers. Mg and O₂

A 
Mg = 0
O2 = 0
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6
Q

What is the oxidation number for a simple ion?

A

the charge on that ion

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7
Q

Give 2 examples of the oxidation number of a simple ion. Cl- and Mg2+

A 
Cl- = -1 
Mg2+ = +2
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8
Q

What is the oxidation number of O in a combined molecule?

A

O = -2

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9
Q

What is the oxidation number for main group metals?

A

an oxidation number equal to the charge on their ions

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10
Q

What are two exceptions for the oxidation numbers of Oxygen?

A

Exceptions:
In compounds with fluorine oxygen has a positive oxidation number (because fluorine is more electronegative than oxygen)
In peroxides, O= -1

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11
Q

What is the oxidation number for hydrogen?

A

H=+1 when it forms compounds with non-metals.

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12
Q

Give 2 examples of the oxidation number for peroxides. H₂O₂ and BaO₂

A

H₂O₂: ON=-2

BaO₂: ON=-2

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13
Q

What is the exception for the oxidation number for hydrogen?

A

In metal hydrides, H=-1

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14
Q

Give an example of the oxidation number for hydrogen in compounds with non-metals. H2O

A

H2O: ON=+1

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15
Q

Give an example of the oxidation number for hydrogen in metal hydrides. NaH and CaH₂

A

NaH: ON=-1, CaH₂: ON=-1

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16
Q

Give an example of the oxidation number for main group metals

A

KCl: ON=+1, MgSO₄: ON=+2

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17
Q

What is the sum of the oxidation number in a neutral compound?

A

0

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18
Q

Give an example of the sum of the oxidation number in a neutral compound. CO₂

A

+4 -2

CO₂: +4 + 2(-2) =0

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19
Q

What is the sum of the oxidation numbers in a polyatomic ion?

A

equal to the charge on the ion

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20
Q

Give two examples of the sum of the oxidation numbers in a polyatomic ion.

A

+6, -2
SO₄⁻²
-3 +1
NH₄⁺

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21
Q

What is the sum of the oxidation number for the most electronegative element?

A

assigned the negative oxidation number

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22
Q

Give an example of the sum of the oxidation number for the most electronegative element.

A

+2 -1

OF₂

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23
Q

What are 5 examples of redox reactions

A

bleaching hair
corrosion of metals
extraction of metals from their ores
combustion of fuels and reactions of batteries that produce electrical energy as well as respiration and photosynthesis

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24
Q

Does oxidation reacts gain or lose oxygen?

A

Gain

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25
Q

Does reduction reacts gain or lose oxygen?

A

lose

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26
Q

Does oxidation reacts gain or lose electrons?

A

lose

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27
Q

Does reduction reacts gain or lose electrons?

A

gain

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28
Q

Does oxidation reacts gain or lose hydrogen?

A

lose

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29
Q

Does reduction reacts gain or lose hydrogen?

A

gain

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30
Q

Does oxidation increase or decrease the oxidation number?

A

increase

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31
Q

Does reduction increase or decrease the oxidation number?

A

decrease

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32
Q

Does oxidation use a reducing or oxidising agent?

A

reducing agents

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33
Q

Does reduction use a reducing or oxidising agent?

A

oxidising agent

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34
Q

Define Electron transfer process

A

oxidation and reduction reactions

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35
Q

What are redox reactions?

A

involve the transfer of electrons from one chemical species to another.

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36
Q

What are the two reactions that are happening simultaneously in a redox reaction?

A

One of the reactants loses electrons in a process called oxidation

One of the reactants gains electrons in a press called reduction

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37
Q

What is an oxidising agent? (4)

A
  • enables or causes another chemical to be oxidised
  • accepts electrons
  • oxidises the reducing agent
  • undergoes reduction
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38
Q

What is a reducing agent? (4)

A
  • enables or causes another chemical to be reduce.
  • loses electrons
  • reduces the oxidising agent
  • undergoes oxidation
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39
Q

What is more likely to act as a reducing agent, metal or non-metal?

A

metals as they tend to lose electrons. , donating electrons to the substance being reduced.

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40
Q

Is the energy required to remove valence electrons alot?

A

It is relatively small amount of energy

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41
Q

What type of metals act as reducing agents? How do they increase in strength?

A

Metals with their smaller number of valence electrons. They increase in reducing agent strength as it becomes more reactive

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42
Q

What type of metals act as oxidising agents? How do they increase in strength?

A

Metals with a larger number of valence electrons. They increase in oxidising agent strength as metals becomes less reactive

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43
Q

The lower the amount of energy required to remove the valence electrons

A

the more readily a metal will act as a reducing agent

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44
Q

As you go down the reactivity series of metals the metal solids

A

become more reactive, meaning metals lower in the series are easier to oxidise and therefore stronger reducing agents

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45
Q

As you go down the reactivity series of metals, metal cations….

A

Become increasingly harder to reduce and are therefore less reactive. Cations higher in the series have a greater attraction for electrons, so they are easier to reduce and are therefore relatively strong oxidising agents

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46
Q

What are electrochemical cells?

A

device in which chemical energy is converted into electrical energy and vice versa

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47
Q

What are galvanic cells?

A

type of electrochemical cell in which chemical energy is converted into electrical energy

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48
Q

What is a battery?

A

combination of cells to obtain a higher potential difference or voltage

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49
Q

What is an external circuit?

A

electrical current that flows through the wire and light globe

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50
Q

What is a spontaneous reaction?

A

What is a spontaneous reaction?

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51
Q

Describe the observations in a galvanic cell (5)

Zinc and Copper(II)

A
  • Redox reactions are occurring
  • Zinc electrode (reducing agent) corrodes as zinc metal forms zin ions in solution
    Zn(s) → Zn²⁺(aq) +2e⁻
  • oxidation of Zn metal releases electrons, which flow through the wire to the copper electrode
  • electrons are accepted by Cu²⁺ ions (oxidising agent) in the solution where ions collide with the copper electrode
    Cu²⁺(aq) + 2e⁻→Cu(s)
  • Cu metal that is formed deposits on the electrode as a dark brown coating
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52
Q

The oxidising agent is also the…

The reducing agent is also the…

A

reaction being reduced

reaction being oxidised

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53
Q

What happens if the reactions come into direct contact with each other?

A

their chemical energy is transformed directly into thermal energy

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54
Q

Why in galvanic cells do half-reactions occur in separate containers?

A

as electrons are transferred through the external circuit so that chemical energy is transformed into electrical energy

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55
Q

How does a galvanic cell operate?

A

designed so that half-reactions occur in two separate compartments,so the oxidising agent and reducing agent do not come into direct contact, electrons can only be transferred through an external circuit connective the negative and positive electrodes.

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56
Q

What are half-cells?

contains? species present form?

A

a half cell contains an electrode in contact with a solution, species present in each half-cell forms a conjugate redox pair (oxidising agent and its corresponding reduced form)

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57
Q

If no metal is present what can be used?

A

inert electrode such as platinum or graphite or a gas electrode, H⁺/H₂ half-cell

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58
Q

Where does oxidation occur in a galvanic half-cell?

A

anode where electrons are released, negative terminal

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59
Q

Where does reduction occur in a galvanic half-cell?

A

cathode where electrons are gained, positive terminal

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60
Q

What is the purpose of the salt bridge? (2)

A
  • contains ions that are free to move so that they can balance charges formed in the two compartments
  • cations move towards cathode and anions move towards anode
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61
Q

What occurs without a salt bridge? (2)

A
  • the solution in one compartment in the galvanic cell would accumulate negative charge and the solution in the other compartment would accumulate positive charge.
  • Accumulation of charge would stop reaction very quickly, hence prevent further reaction
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62
Q

Where is the stronger reducing agent held in?

A

negative electrode (anode)

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63
Q

Where is the stronger oxidising agent held in?

A

postive electrode (cathode)

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64
Q

How does the current flow within the half cells?

A

as one half-cell has a greater tendency to push electrons into the external circuit than the other half-cell

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65
Q

What is potential difference (E)?

A

electromotive force/emf referred to voltage (V), which exists between two half-cells

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66
Q

What standard conditions are cells usually measured under?

pressure, concentration and symbol and temp

A
  • a pressure of 1 bar (100kPa)
  • 1 mol/L concentration of solution
  • E°
  • 25°C
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67
Q

Why is it impossible to measure the potential difference of an isolated half-cell?

A

both oxidation and reduction must take place for potential difference to exist, but you can assign the standard half-cell potential (E°) by connecting it to a standard reference cell (H⁺/H₂) and measuring the voltage produced.

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68
Q

What is the standard reference half-cell used?

A

hydrogen half-cell/standard hydrogen electrode, under standard conditions its E° Value is assigned as 0

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69
Q

Summarise the energy transformation in a secondary cell . (2)

A
  • When a secondary cell discharges it acts as a galvanic cell, converting chemical energy into electrical energy
  • When the cell is recharged, it acts as a electrolytic cell. Electrical energy is transformed into chemical energy in an electrolytic cell.
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70
Q

What happens when the battery is discharging? (3)

A
  • Chemical energy is cinverted into electrical energy
  • acts as a galvanic cell
  • spotaneous reaction
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71
Q

What happens when the cell is recharging? (3)

A
  • Electrical energy is converted into chemical energy
  • acts as electrolytic cell
  • non-sponatneous reaction
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72
Q

Why are car batteries more widely used type of secondary cell?

A

Lead acid batteries as they are cheap and reliable, providing high currents and have a long lifetime.

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73
Q

What recharges the car battery?

A

Once the engine starts, an alternator which is run by the engine, provides electrical energy to operate the car’s electrical system and recharges the battery

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74
Q

When is a car battery used?

A

Start the cars’ ongoing and operate the cars’ electrical accessories when the engine is not running

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75
Q

What potential difference does each cell in a car battery have? And overall with all 6 cells connected in a series?

A

2V and 12V

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76
Q

What enables the battery to be recharged?

A

Produce of both electrode regions lead (II) sulphate forms as a solid of the surface

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77
Q

What are fuel cells?

A

continuously produce electricity for as long as is fuel is fed into the cells

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78
Q

What is the one limitation of primary/secondary cells from fuel cells?

A

Limitation of other cells is that they contain relatively small amounts of reactants. But reactants are supplied continuously allowing constant production of electrical energy in fuel cells

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79
Q

What happens when E° value is negative when connected to hydrogen half-cell?

A

it indicates that the electrode is negative and oxidation is occurring

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80
Q

What is the E° value known as when it is connected a hydrogen half-cell?

A

the standard electrode potential and the standard reduction potential

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81
Q

What does the standard electrode potential give a numerical measure of?

A

It gives a numerical measure of the tendency of a half-cell reaction to occur as a reduction reaction

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82
Q

In a galvanic cell, the stronger reducing agent is oxidised

A

so it is in the half-cell with the negative electrode (anode)

83
Q

The stronger oxidising agent is reduced

A

so it is in the half-cell with the positive electrode (cathode)

84
Q

Strong reducing and oxidising agents donate electrons….

A

more readily than weak ones

85
Q

Strong reducing agents have…

A

weak conjugate oxidising agents

86
Q

Strong oxidising agents have….

A

weak conjugate reducing agents

87
Q

What type of reactions occur in cells and batteries?

A

spontaneous redox reactions that give something a source of energy

88
Q

What are the three basic types of galvanic cells?

A

Primary cells
Secondary cells
Fuel cells

89
Q

What are primary cells?

A

cells that disposable and designed not to be recharged

90
Q

What are some types of primary cells?

A

common commercial alkaline cells, such as batteries in a torch or remote control are non rechargeable cells

91
Q

When does a primary cell get used up?

A

When most of the reactants are used up, the cell reaction reaches equilibrium

92
Q

What are secondary cells?

A

which are rechargeable and designed to be reused many times

93
Q

What are some types of secondary cells? (3 types)

A
  • lithium-ion cells and nickel-metal hydride cells (known as accumulators)
  • found in laptops, mobile phones, cameras and portable power tools
  • fundamental to the operation of electric vehicles and solar power energy storage system
94
Q

What reaction occurs when a secondary cell is recharged? (2)

A

the cell reaction must occur in reverse: the products of the reaction are converted back to the original reactants

Electrical energy supplied by the charger is converted into chemical energy in the cell.

95
Q

What is the secondary cell attached to? (2)

A

by connecting the cell to a charger, source of electrical energy which has a potential difference a little greater than the potential difference of the cell.

positive terminal of charger is connected to the cell’s positive electrode and the negative electrode of the charger to the cell’s negative electrode

96
Q

What must be possible for reactants to regenerate?

A

the products formed in the cell during discharge must remain in contact with the electrodes in convertible form

97
Q

What is a key difference between a fuel cell and a primary or secondary cell?

A

reactants are not stored in the fuel cell. They must be continuously supplied from an external source

98
Q

What chemical energy does fuel cells use?

A

chemical energy of hydrogen or other fuels to cleanly and efficiently generate electricity.

99
Q

What are fuel cells used for?

A

a source of power for transport and for emergency back-up power applications

100
Q

What makes fuel cells efficient?

A

fuels cells transform chemical energy directly into electrical energy, enabling efficient use of the energy released by spontaneous redox reactions

101
Q

When does energy loss occur for fuel cells when applied?

A

occur in coal-fired power stations and combustion engines are avoided with a consequential reduction in the volume of greenhouse gases produced.

102
Q

Why do fuel cells use hydrogen?

A

a fuel produces electricity, water, heat and very small amounts of nitrogen dioxide and other emissions

103
Q

What are the two compartments in a fuel cells?

A

one for hydrogen gas and the other for oxygen gas

104
Q

Describe the fuel cell design. (6)

A
  • gas compartments are separated form each other by two porous electrodes and usually a KOH electrolyte solution
  • electrode at the hydrogen compartment is the anode
  • electrode at the oxygen compartment is the cathode
  • oxygen gas inlet on the right
  • hydrogen gas inlet on left
  • water outlet on the bottom where the electrolyte solution is.
105
Q

Explain the 3 step reaction that occurs within a fuel cell.

A

At the anode, H₂ (the fuel) is oxidised by reacting with hydroxide ions :
H₂(g)+2OH⁻(aq)→2H₂O(l) +2e⁻
At the cathode, oxygen gas is reduced:
O₂(g) + 2H₂O(l) + 4e⁻→4OH⁻(aq)
The overall reaction for the reaction is:
2H₂(g)+O₂(g) →2H₂O(l)

106
Q

How much voltage does each fuel cell produce?

A

1V

107
Q

How are higher voltages obtained for a fuel cell? And what is the byproduct?

A

higher voltages are obtained by connecting a number of fuel cells in series to form a battery or fuel cell stack.
byproduct is water and heat is given off

108
Q

At what nature of the electrodes is best for efficiency of a fuel cell?

A

electrodes must both conducting and porous to allow the hydrogen and oxygen to diffuse through them to come into contact with the ions in the electrolyte and to allow the redox half-reactions to occur at their surface

109
Q

What does the size of the current drawn from a fuel cell depend on?

A

size of current that can be drawn from a fuel cell depends on the surface are of the electrodes

110
Q

What do catalysts help with in a fuel cell?

Normally

A

enhancing the rate of reaction and the current that can be produced from a cell

111
Q

What do catalysts help with in a fuel cell? and what catalyst is used?
(anode)

A

in an anode increase the rate of oxidation of fuel gas

platinum metal is used as a catalyst at this electrode

112
Q

What do catalysts help with in a fuel cell? and what catalyst is used?
(cathode)

A

the cathode catalyst, which increases the rate of reduction half-reaction
can be made from a different material, such as nickel power or nanomaterial

113
Q

Why is producing hydrogen not great?

A

if the production of hydrogen is produced using renewable energy the production of hydrogen can result in significant levels of greenhouse gases and other pollutants

114
Q

What are two practical methods of generating hydrogen without producing CO2?

A
  • using electrical energy to convert water to hydrogen, electricity can be generated from renewable sources such as solar-power farms and wind farms
  • collecting biogas from landfill sites and converting the methane in the gas to hydrogen
115
Q

What are is 5 advantage of fuel cells? (convert)

A

convert chemical energy directly to electrical energy.
- it is more efficient than the series of energy conversions that takes place in power stations that burn fossil fuels: chemical energy→mechanical energy→heat energy→mechanical energy→electrical energy

116
Q

What are 5 advantages of fuel cells?

A

Fuel cells convert chemical energy directly to electrical energy more efficiently

Hydrogen fuel cells produce water and heat as byproducts
no greenhouse gases, such as CO2 are released fuel cells can use a variety of fuels

electrodes do not get depleted during the process of generating electricity

no memory effect (as with secondary) efficiency doesn’t reduce over time

fuel cells will generate electricity for as long as the fuel supplied.
Conventional batteries need to be recharged or replaced

117
Q

What are 6 disadvantages of fuel cells?

A

fuel cells require a constant fuel supply

fuel cells are expensive. Still developing, and are not made in large numbers

some types of fuel cells use expensive electrolytes and catalysts

hydrogen used in many fuel cells is mainly sourced from fossil fuels, the process involves energy losses and generates greenhouse gases

The use of fuel cells in transport will require an extensive network of hydrogen filling stations before it can become widespread

significant issues associated with the storage of hydrogen fuel

118
Q

What is dry corrosion? (direct corrosion)

A

Direct reaction with oxygen in the air to form a metal oxide

119
Q

What are some examples of metals that dry corrode?

A

Na is so reactive that it must be stored under oil to prevent contact with oxygen
4Na(s) + O₂(g) → 2Na₂O(s)

While Al forms a tough, impervious coating of aluminium (Al₂O₃), which protects the metal underneath from further contact with oxygen

120
Q

What is wet corrosion?

A

The presence of moisture accelerates the corrosion of iron, formation of rust as flaky, brown-red coating

121
Q

What are two ways wet corrosion is accelerated?

A

the presence of water

impurities, such as salt and acidic pollutants that dissolve in the water

122
Q

When can wet corrosion be reduced?

A

corrosion can be reduced when iron is alloyed with certain other materials or when it has a protective coating

123
Q

What is rust?

A

hydrate oxide of iron Fe₂O₃·xH₂O.

with 1-3 water molecules associated with the iron oxide

124
Q

What are the 4 steps of wet corrosion? (step 1)

A

Iron is oxidised to form Fe²⁺ ions at one region on the iron surface.
Fe(s)→Fe²⁺(aq) + 2e⁻
at another region on the surface, using the electrons produced by the oxidation process, oxygen is reduced in the presence of water to hydroxide ions.
Overall: O₂(aq) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)

125
Q

What are the 4 steps of wet corrosion? (step 2)

A

Formation of precipitate of iron(II) hydroxide

Fe²⁺(aq)+2OH⁻(aq)→Fe(OH)₂(s)

126
Q

What are the 4 steps of wet corrosion? (step 3)

A

Further oxidation of iron(II) hydroxide occurs in the presence of oxygen and water to produce iron(III) hydroxide, a red-brown precipitate:
4Fe(OH)₂(s) + O₂(aq) + 2H₂O(l) → 4Fe(OH)₃(s)

127
Q

What are the 4 steps of wet corrosion? (step 4)

A

In the air, the iron(III) hydroxide loses water to form hydrated iron(III) oxide (Fe₂O₃∙xH₂O), rust

128
Q

What are the 4 steps of wet corrosion?

A
  1. Iron is oxidised to form Fe²⁺ ions at one region on the iron surface.
    Fe(s)→Fe²⁺(aq) + 2e⁻
    at another region on the surface, using the electrons produced by the oxidation process, oxygen is reduced in the presence of water to hydroxide ions.
    Overall: O₂(aq) + 2H₂O(l) + 4e⁻ → 4OH⁻(aq)
  2. Formation of precipitate of iron(II) hydroxide
    Fe²⁺(aq)+2OH⁻(aq)→Fe(OH)₂(s)

3.Further oxidation of iron(II) hydroxide occurs in the presence of oxygen and water to produce iron(III) hydroxide, a red-brown precipitate:
4Fe(OH)₂(s) + O₂(aq) + 2H₂O(l) → 4Fe(OH)₃(s)

  1. In the air, the iron(III) hydroxide loses water to form hydrated iron(III) oxide (Fe₂O₃∙xH₂O), rust
129
Q

How does the presence of dissolved ions/salts in water increase rate of wet corrosion?

A

Salts are good electrolytes and accelerate the rusting process because they increase the conductivity of moisture on the surface of the iron.

130
Q

What are 2 ways to prevent corrosion?

A

Surface protection

Electrochemical protection

131
Q

What is surface protection?

A

involves covering the surface of the iron to prevent contact with oxygen and moisture.
materials such as paint and plastic can be used for this purpose

132
Q

What can be used for surface protection? Including moving parts(3)

A
  • alloying it with small quantities of metals, such as chromium, nickel, manganese or molybdenum to produce stainless steel. Atoms of the metals used to make the alloy are all bonded into the metallic lattice. It may slightly oxidise in air, but the oxide coating produced is continuous and unreactive. Oxide layer protects the metal from further oxidation.
  • materials such as paint and plastic can be used for this purpose
  • moving parts, such as a bicycle chain, can be coated with oil or grass to both reduce friction and prevent contact with oxygen and moisture
133
Q

What can iron and steel be coated in? (and tin cans)

A

coated with thin layers of less reactive metal in a process known as electroplating.

tin cans are steel plated with tin, tin is much less reactive metal than iron and does not corrode greatly in the atmosphere

134
Q

What are two types of electrochemical protection?

A

Cathodic protection

Sacrificial protection

135
Q

What is cathodic protection?

A

involves the use of low-voltage, direct current power supply to give the iron being protected a negative charge.

136
Q

How does cathodic protection reduce corrosion? (2)

A

for corrosion to occur, iron loses electrons (is oxidised)

act of making iron negative pushes electrons towards the iron, thus reducing the chance of oxidation occurring. Instead the iron becomes the site of reduction reactions and forms the cathode.

137
Q

What is sacrificial protection?

A

involves iron acting as the cathode
A more easily oxidised metal, such as zinc forms the anode and is sacrificed in order to protect the iron from corrosion. More reactive metal loses electrons and forms metal cations in preference to the iron

138
Q

What is galvanised iron? How does it help with corrosion?

A

Iron coated in zinc

When the zinc coating is scratched, the iron is still protected because the zinc loses electrons more readily than the iron. Zinc slowly corrodes, but the iron is protected for years

139
Q

Give an example of how sacrificial protection is used in underground steel pipelines, bridge pillars and steel hulls of ships

A

can be protected by connecting them to blocks or plates of zinc or other more reactive metals such as magnesium.
Because these metals are more easily oxidise than iron, they lose electrons which are transferred to the iron where reduction of oxygen and water occurs
reactive metals, which are called sacrificial anodes are eventually consumed and must therefore be replaced, but this is less expensive than replacing the steel structures that are being protected.

140
Q

What is oxidation?

A

Loss of electrons

141
Q

What is reduction?

A

Gain of electrons

142
Q

What is a disproportionation reaction?

A

Process where a substance undergoes oxidation and reduction simultaneously. (Undergoes self-redox)

143
Q

Describe the structure of a battery (7)

A 
Metal cap (+)
Cathode: outer steel case
Powered zince
Mixture of mangahese dioxide and carbon
Potassium hydroxide electrolyte 
Anode: steel or brass 
Metal base (-)
144
Q

Describe the construction of a lead-acid battery

A

Lead plates (-) anode and lead plates impregnated with lead(IV) oxide (+) cathode are seperated by a porous seperator and about 4mol/L of sulfuric acid

145
Q

What is the chemical equations for the corrosion on iron?

A

2Fe(s) + O2(aq) + 2H2O(l) -> Fe2+(aq) + 4OH-(aq)

Fe2+(aq) + 2OH-(aq) -> Fe(OH)2(aq)

146
Q

How is corrosion of iron a redox reaction?

A

As the oxidation number increases from 0 to 2 for the oxidation reaction
And the oxidation number increases from +2 to +3 for the reduction reaction

147
Q

Underwater steel pillars often corrode more rapidly just beneath the surface of the water than above. Suggest a reason.

A

Reduction of oxygen occurs when there is plenty of moisture and the oxygen concentration is high; that is, at or just above the surface of the water. The cathodic region forms here. So the anodic region is close to the cathode and where the oxygen concentration is less; that is, just below the surface. Oxidation or corrosion therefore occurs more rapidly below the surface.

148
Q

What are 6 common oxidising agents?

A

O₂, Cl₂, MnO₄⁻, Cr₂O₇²⁻, ClO⁻, H⁺

149
Q

What are 5 common reducing agents?

A

Zn, C, H₂, Fe²⁺, C₂O₄²⁻

150
Q

Why can hydrogen peroxide solution spontaneously decompose to form water and oxygen?

A

Hydrogen peroxide is both a strong oxidising agent and as a weak reducing agent (it appears in electrochemical series on the left side of one-half reaction and the right side of another ) hydrogen peroxide therefore reactive itself. The reaction is very slow unless a catalyst such as manganese dioxide is used.

151
Q

Why does tin metal added to a solution of tin(II) chloride prevent oxidation of tin(II) ions by oxygen air?

A

Tin(II) ions can be oxidised to Tin(IV) ions by a suitable oxidising agent. Tin metal can reduce Sn(IV) ions so that they re-form Sn(II) ions.

152
Q

What are 5 differences between a fuel cell and a primary cell?

A

Fuel cell produces electrical energy continuously when reactants are provided; primary cells can produce power for only a limited time until their reactants are depleted
Fuel cell uses combustioe fuel and air or oxygen; primary cells can be made from different cimbinarions of two conjugate redox pairs
The reactants in fuel cells are usually gaseous; a range of solids, liquids and gases can be used as the reactants in the primary cells
Electrodes in fuel cells must be porous; electrodes in primary cells need not be porous
Catalysts are used in fuel cells to increase cell efficiency; catalyst are not required in primary cells

153
Q

What are 5 similarities between fuel cells and primary cells?

A

Both cells are designed to convert chemical energy into electrical energy at relatively high efficiencies.
Oxidation and reduction reactions take place in different
Both have an anode, which is negative, at which an oxidation half-reactions occurs
Both have a cathode, which is positive, at which a reduction half-reaction occurs
Both have an electrolyte, which provides ions to balance charges formed at the electrodes; cations flow towards the cathode and anions flow towards the anode.

154
Q

What is electrolysis?

A

Involves passage of electrical energy from, a power supply (such as a battery) through a conducting liquids

155
Q

What is electroplating of copper?

A

a thin surface coating of metal only a fraction of a milllimetre thick, is applied over another metal surface, perform in electrolytic cells.

156
Q

What is the process for electroplating of copper? (4)

A
  • object to be plated, like key, is connected by a wire to the negative terminal of a power supply. Object becomes the negative electrode
  • Rod or sheet of copper metal is connected to positive terminal of the power supply. Metal becomes s the positive electrode in cell
  • Two electrodes are immersed in aqueous solution, such as copper(II) sulphate solution (which contains ions of the metal to be plated). Copper sulphate is an ionic solid composed of Cu2+ and SO42- ions, which dissociated when dissolved in water. Describe as an electrolyte as it conducts electricity
  • During electrolysis, reactions occur at the surface of both electrodes
157
Q

What occurs in the first step in the electroplating of copper?

A

object to be plated, like key, is connected by a wire to the negative terminal of a power supply. Object becomes the negative electrode

158
Q

What occurs in the second step in the process of electroplating of copper?

A

Rod or sheet of copper metal is connected to positive terminal of the power supply. Metal becomes s the positive electrode in cell

159
Q

What occurs the third step of the process of electroplating of copper?

A

Two electrodes are immersed in aqueous solution, such as copper(II) sulphate solution (which contains ions of the metal to be plated). Copper sulphate is an ionic solid composed of Cu2+ and SO42- ions, which dissociated when dissolved in water. Describe as an electrolyte as it conducts electricity

160
Q

What happens in the fourth step in the process of electroplating of copper?

A

During electrolysis, reactions occur at the surface of both electrodes

161
Q

What happens at the negative electrode during the process of electroplating of copper?

A

copper ions are attracted to negative electrode (object to be plated), where they accept electrons and converted too copper metal:
Cu2+ + 2e- -> Cu(s)
Coating of copper forms on object, as reduction of copper ions occurs

Cathode

162
Q

What happens at the positive electrode in the process of electroplating of copper? Is it the cathode or anode?

A

positive terminal of power supply withdraws electrons from copper electrode, causing oxidation to occur

Copper metal slowly dissolves as Cu2+ ions are formed
Cu(s) -> Cu2+(aq) + 2e-

Reaction replaces Cu2+ ions in solutions that were consumed by the reaction at the negative electrode
Concentration of copper ions remains constant

Anode

163
Q

What type of electrodes are used in electrolysis of molten sodium chloride?

A

Platinum metal or graphite is used for the electrodes as there are inert

164
Q

What occurs at the negative electrode for the electrolysis of molten sodium chloride?

A

Na+ ions in electrolyte are attracted to negative electrode, accepting electrons and becoming sodium atoms
Na+(l) + e- -> Na(I)

Sodium is less dense than molten sodium chloride therefore floats to the top of cell

165
Q

What occurs at the positive electrode for the electrolysis of molten sodium chloride?

A

Cl ions move to the positive electrode to give up electrons to form Cl atoms

Bubbles of Cl2 gas appear at the electrode
2Cl-(l) -> Cl2(g) + 2e-

Electrons from chloride ions move through electrode to the power supply

166
Q

Why in electrolysis do products need to be kept apart although it occurs in a single container?

A

They will react with each other to reform

167
Q

Are electrolysis reactions spontaneous or non spontaneous?

A

Non spontaneous reaction as electrical energy must be supplied to convert it into chemical energy

168
Q

What are 4 differences between galvanic cells and electrolytic cells? (Electricity)

A

Galvanic cells produce electricity

Electrolytic cells consumes electricity

169
Q

What are 4 differences between galvanic cells and electrolytic cells? (Spontaneous)

A

Galvanic cells are spontaneous reactions

Electrolytic cells are non-spontaneous

170
Q

What are 4 differences between galvanic cells and electrolytic cells? (Conversion)

A

Convert chemical to electrical

Convert electrical energy to chemical

171
Q

What are 4 differences between galvanic cells and electrolytic cells? (Anode and cathode )

A

Anode is negative and cathode is positive

Anode is positive and cathode is negative

172
Q

What are two similarities between galvanic and electrolytic cells?

A

Oxidation occurs at anode and reduction at cathode

Anions flow towards anode and cations flow towards cathode

173
Q

What are two applications of electrolysis?

A

Electroplating and electrorefining

174
Q

What is electroplating?

A

Plating a thin surface coatings of one metal over the surface of another metal

175
Q

Where is the object to be plated connected to in electroplating of tin cans? (Positive or negative terminal)

A

The negative terminal of power supply

176
Q

Where is the electrode of tin metal connected to in electroplating tin cans? (Positive or negative terminal of power)

A

positive terminal of power supply

177
Q

What electrolyte solution is the object immersed in (tin)?

A

Solution containing ions of the metal to be plated

Tin(II) nitrate solution

178
Q

What occurs at the negative electrode in electroplating of tin cans?

A

Tin ions are attracted to cathode, undergo reduction and are converted to tin metal
Sn2+(aq) + 2e- -> Sn(s)

Coating of tin is formed on object

179
Q

What occurs at the positive electrode in electroplating of tin cans?

A

Tin undergoes oxidation
Sn(s) -> Sn2+(aq) + 2e-

As Sn2+ ions are consumed by the reaction at the cathode and produce at the anode, the overall concentration of tin ions remains constant.

180
Q

Object to be plated is at…

A

Cathode (negative)

181
Q

Electrode of metal is at….

A

Anode (positive)

182
Q

What are 6 factors that are critical for achieving a smooth, tightly bonded metal coating that prevents oxygen and water from reaching the base metal:

A 
Voltage
Current
Temperature
Electrode positions
Concentration 
Identity of electrolyte
183
Q

What is blister copper in electrorefining of copper?

A

Copper metal being extracted from its ores by smelting, where copper ore is heated strongly in air to produce impure molten copper metal. As metal solidifies hot gases escape and the surface becomes blistered.
Containign 2% of impurities

184
Q

What is electorefining of copper?

A

Blistered copper is purified through a process called electrorefining, uses a large amount amount of electrical energy, so refinery must be located close to a source of abundant cheap power

185
Q

What happens at the positive electrode during electrorefining of copper? (3)

A

anode

Electrons are drawn away from the blister copper anode to positive terminal of power source, copper and impurities (such as nickel and zinc) that are more reactive (stronger reducing agents) than copper, are oxidised
Cu(s() -> Cu2+(aq) + 2e-

Impurities less reactive than copper like silver, gold and platnum are not oxidised and simply fall from the anode, collecting at the bottom

186
Q

What happens at the negative electrode in electrofining of copper? (3)

A

cathode

Copper metal is deposited as electrons from the power source and are accepted by metal ions from solution
Copper metal is the strongest oxidising agent present in solution therefore it is formed
Cu2+(aq) +2e- -> Cu(s)

In copper electrorefining cell, copper metal undergoes oxidation at the anode and copper ions undergo reduction at cathode, yields copper metal of high purity

187
Q

Why do iron objects found at the bottom of the ocean have very few signs of corrosion?

A

As there is very little dissolved O2 at the bottom of the ocean. For corrosion to occur to an appreciable extent-both H2O and oxygen must be present.

188
Q

Why is the anode negative in a galvanic cell but positive in an electrolytic cell?

A

The process of oxidation involves the loss of electrons, and oxidation occurs at the anode that is why the galvanic cell is negative. But in a electrolytic cell, since the positive terminal of power is connected to it the electrode it draws away the electrons from the anode therefore oxidation occurs.

189
Q

Why are two half cells usually used in a galvanic cell whereas the reactants of an electrolysis cell are often placed in a single container?

A

As reactions in galvanic cells are spontaneous, in reactants were in one container and in contact with one another they would be releasing heat energy instead of electrical energy.

While since electrolytic cells are non-spontaneous, both reactions can occur in the same container, however the products should not come into contact or a reaction will occur.

190
Q

The initial voltage found measured in a cell was higher than expected. Give one possible reason for this observation.

A

not measure at 25C

not a 1 mol L solution

191
Q

What does scratched tin cause iron to corrode more rapidly than if the tin were not there?

A

As iron is more reactive than tin. if it is scratched the iron will become the anode and will oxidise quicker than a piece of uncoated iron. The increase in corrosion rate occurs in any situation where iron is in contact with any metal less active than itself

192
Q

Why does the rate of production of electrical current from an electrochemical cell decreases as it operates?

A

rate of reaction decreases as reactants are consumed. Therefore the rate of production of electrical current decreases.

193
Q

During the operation of an electrochemical cell, why is it important that the anode and cathode do not come into contact with each other?

A

Short circuits and so current will not flow. Current will be forced through an external circuit.

194
Q

Why is sulfuric acid and not hydrochloric acid used in the acidification of potassium permanganate solution prior to it being used in redox titration?

A

Chloride ions in HCl will react with the permanganate ions, since the permanganate ions oxidise the chloride ions, there are less (or no) MnO4-(aq) to react with the reducing agent

195
Q

Explain the function of the electrolyte.

A

allow for the movement of ions
maintains electrical neutrality of ions
completes the circuit

196
Q

State 2 reasons why the theoretical (calculated) value was not the same as the actual EMF generate by the fuel cells.

A

fuel cell is not at standard conditions, eg. not at 25ºC

resistance in the wire

197
Q

Explain the chemical principles of an electrolytic cell

A

electrolytic cell utilises an external applied voltage this allows a non-spontaneous redox reaction to occur.

198
Q

What must happen for a cell to go flat (stop providing electrons)? (2)

A
  • Reactants are used up

- chemical equilibrium is reached

199
Q

State the reason for the reactants being kept in seperate half cells.

A

The electrons which are transferred can then pass through an external circuit (rather than being transferred through direct contact)

  • create a current
  • create a potential difference across half cells
200
Q

What is the reaction at the anode for 1.0M solution of NaCl?
Possible anode reactions:
- 2Cl- -> Cl2 + 2e- E=-1.36V
- 2H2O -> O2 + 4H+ + 4e-

A

2H2O -> O2 + 4H+ + 4e-

201
Q

Why is phenolphthalein a good indicator for NaCl?

A
  • NaCl ions are neutral as they do not undergo hydrolysis
202
Q

What is the reaction at the anode for concentrated solution of NaCl?
Possible anode reactions:
- 2Cl- -> Cl2 + 2e- E=-1.36V
- 2H2O -> O2 + 4H+ + 4e- E=-1.23V (E=-0.82V at pH 7)

A

2Cl- -> Cl2 + 2e-

203
Q

What is the reaction at the anode for 1.0M solution of NaCl?
Possible cathode reactions:
- Na* + e- -> Na E= -2.71V
- 2H2O + 2e- -> H2 + OH- E= -0.83V (E= -0.41V at pH 7)

A

2H2O + 2e- -> H2 + OH- E= -0.83V

204
Q

What is the reaction at the anode for concentrated solution of NaCl?
Possible cathode reactions:
- Na* + e- -> Na E= -2.71V
- 2H2O + 2e- -> H2 + OH- E= -0.83V (E= -0.41V at pH 7)

A

2H2O + 2e- -> H2 + OH- E= -0.83V

MECH U3/4 (8 decks)

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