Unit 3.1 - Redox and Standard Electrode Potential Flashcards
1
Q
What does an oxidation state show?
A
How many electrons the atom has used in bonding
2
Q
What can the method of using oxidation states also be applies to?
A
Covalent substances where complete transfer of electrons does not occur
3
Q
Oxidation state of an uncombined element
A
Zero
4
Q
Oxidation state of a diatomic molecule
A
Zero
(Still an uncombined element)
5
Q
Describe the oxidation states of the elements in a compound of two elements
A
One element has a positive oxidation state
The other has a negative oxidation state
6
Q
Which element has the negative oxidation state in a compound?
A
The more electronegative element
7
Q
What’s the sum of the oxidation states in a compound?
A
Zero
8
Q
What is the oxidation state equal to in ions?
A
The charge on the ion
9
Q
Oxidation states of group 1 elements - Li, Na and K
A
+1
10
Q
Oxidation state of group 2 elements - Mg, Ca, Sr and Ba
A
+2
11
Q
Oxidation state of group 3 elements - Al
A
+3
12
Q
Oxidation state of hydrogen (+exception)
A
+1, except in metal hydrides, where it’s -1
13
Q
Oxidation state of group 7 elements, such as Cl (+exception)
A
-1
Except with oxygen (variable)
14
Q
How do we work out the oxidation state of an individual element?
A
Multiply up however many of the other element there is
15
Q
How do we work out changes in oxidation states of a specific element?
A
Remember to use one of it
16
Q
What do redox reactions involve?
A
Both oxidation and reduction
17
Q
Oxidation
A
Loss of electrons
Increase in oxidation state
18
Q
Reduction
A
Gain of elections
Decrease in oxidation state
19
Q
Alternative way of defining oxidation
A
Gain of oxygen, loss of hydrogen
20
Q
Alternative way of defining reduction
A
Loss of oxygen, gain of hydrogen
21
Q
Name 4 types of reaction that don’t involve redox
A
Precipitation
Acid-base
Acid-carbonate
Thermal decomposition (mainly)
22
Q
Reducing agent
A
A species that reduces another species and is itself oxidised during the redox reaction
23
Q
Oxidising agent
A
A species that oxidises another species and is itself reduced during the redox reaction
24
Q
What do we need to remember with positive oxidation states?
A
Include the positive sign
25
Do we take into account big numbers with oxidation states?
No
26
Electronegativity
The ability of an atom to attract a pair of electrons to itself in a covalent bond
27
What do we do if a question asks us to prove if a reaction is/isn’t redox?
Assign oxidation states
28
Disproportionation reaction
The same element is reduced and oxidised in the same reaction
29
The same element is reduced and oxidised in the same reaction
Disproportionation reaction
30
What is a half equation?
The oxidation half or the reduction half of the redox equation
31
Where are electrons shown in an oxidation half equation?
On the right
32
Where are electrons shown in a reduction half equation?
Electrons on left
33
Why are electrons added to the specific sides that they are in half equations?
To balance the charge
34
What do we do to form redox equations from half equations?
Multiply them out to balance the electrons, and don’t include the electrons in the final equation
35
What type of reaction takes place to generate electricity in a cell (e.g - a battery)?
A redox reaction
36
What is one simple method of generating electricity in a cel?
2 metal strips (copper and zinc) —> electrodes
In a solution of copper (II) sulfate —> electrolyte
Electricity is generated when the two metals are connected externally with a wire, short circuited
37
Electrode
A solid that carries charge
38
Electrolyte
A substance that carries charge
39
What is a more convenient arrangement for generating electricity in a cell than connecting metals externally with a short circuited wire?
Using a reversible cell
40
What does a reversible cell include?
The copper foil is placed in copper (II) sulfate solution
The zinc foil is place in zinc sulfate solution
41
What is the arrangement of zinc foil in zinc sulfate solution known as?
A half cell
42
What makes up a half cell?
A metal placed in a solution containing those metal ions
43
How are half cells connected?
By means of a conducting medium known as a salt bridge
44
Salt bridge
A piece of apparatus that connects the solution in 2 half cells so that the circuit can be complete and the current can flow without the solutions mixing
45
Purpose of a salt bridge
Allow the current to flow without the solutions mixing
46
What does a salt bridge contain?
A solution which is a strong electrolyte to conduct electricity but is not reactive
47
Which solution is usually used in a salt bridge?
Potassium chloride or potassium nitrate solution
48
What properties does the solution used in a salt bridge need?
It needs to be a strong electrolyte to be able to conduct electricity, but not be reactive
49
Cathode charge
+
50
Anode charge
-
51
Describe how a redox cell including zinc in zinc sulfate solutions and copper in copper sulfate solution would work
When the cell is connected externally…
Zinc metal electrode is oxidised
Loses electrons
Electrons flow from the zinc electrode to the copper metal and to the copper solution
Copper (II) ions gain electrons
52
What voltage (EMF) does the cell is zinc into copper have’?
1.10 volts
53
What is assumed when saying the voltage (EMF) of a cell?
Standard conditions
Pure metals
54
What does voltage reflect in the redox cell between copper and zinc?
The flow of electrons from the zinc to the copper
55
Equation used to represent a redox cell
[-] oxidation —> II reduction —> [+]
56
What is each cell made up of?
2 half cells
57
Zinc to copper as redox cell equation
[-] Zn (s) I Zn^2+ (aq) II Cu^2+ (aq) I Cu (s) [+]
58
What must each half cell contain?
Both the reactants and products of the half reaction
A metal to allow electrons to flow into or out of the half cell
59
3 types of half cell to learn
Metal/metal ions half cell
A gas in contact with a solution of non-metal ion, with an inert metal electrode
A solution containing ions of a metal in 2 different oxidation states, again using an inert metal electrode
60
Describe metal/metal ions half cells
Metals in contact with meta ions (e.g - Zn (s) with Zn (aq)
Piece of metal to act as the electrode
Solution containing a 1moldm^-3solution of the metal ions to act as the electrolyte
61
Does zinc change colour when involved in a half cell?
No
62
Does copper change colour when involved in a half cell?
Blue solution may lose dolour as the copper ions are reduced
63
In which situation are inert platinum electrodes used?
With non-metals
64
Why must inert platinum electrodes be used for non-metals in half cells?
Non-metals are non-conductors
65
Why is an inert platinum electrode specifically chosen in non-metal containing half cells?
It’s a non-reactive conducting medium
66
What does an inert platinum electrode in a half cell allow to happen?
Allows electrons to flow in or out of the half cell
67
How would we show the presence of an inert platinum electron in a half cell diagram?
Pt (s)
68
In which situation would we include Pt(s) on a half cell diagram?
On any side of the half cell diagram that ends in (aq) or (g) or (l), as they’re not (s)
69
What is an inert platinum electrode typically used for?
A hydrogen electrode (H2/H+ (aq)) or oxygen (O2/OH- (aq)) half cells
70
What happens to the gas in a half cell?
Is bubbled over the inert electrode which is dipping in solution of the ions
71
Do half cells involving gases cause any colour changes?
No
72
In a half cell with a solution containing ions of a metal in two different oxidation states, which substances are platinum electrodes usually used with?
Transition metals
E.g : Fe^2+/Fe^3+ and Mn^2+/MnO4^-
73
Which solutions cause colour changes in half cells containing a solution containing ions of a metal in two different oxidation states, again using an inert metal electrode?
Fe^2+ —> green
Fe^3+ —> orange
Mn^2+ —> colourless
MnO4- —> purple
74
When does a colour change usually take place in a half cell?
When reduction or oxidation takes place in a half cell containing a solution containing ions of a metal in two different oxidation states
75
What does a vertical line do in half cll representations?
Separate substances in different physical states
76
Different physical states
(s), (l), (g), (aq)
77
What do commas represent in half cell representations?
They separate substances in the same physical state
78
What does II represent in a half cell representation?
The salt bridge (the boundary between 2 half cells)
79
In which direction does the electricity flow in a Redox cell?
From the anode to the cathode
80
What happens when a battery is connected to a cell?
The cell reaction is reversed
81
Standard conditions
Concentration —> 1moldm^-3
Pressure —> 1atm
Temperature —> 25/296k
82
Symbol for standard electrode potential
E°
83
What does ° represent in E°?
Standard conditions
84
Standard electrode potential
The potential difference between the element in contact with a solution of its ions, concentration 1moldm-3, which is measured against a standard hydrogen electrode under standard conditions, 298K and 1atm
85
Where is the hydrogen electrode always written in half cell representations
As the left hand electrode of a cell
86
What does the sign of the quoted voltage in a standard electrode potential give you and why?
The sign of the other electrode, as the hydrogen electrode is always written as the left-hand electrode of a cell
87
Electrode potential of a half cell
A measure of the tendency of the electrode to attract electrons to itself (i.e - for the cell to be reduced)
88
How are electrode potentials written?
Oxidised state + electrons ——> reduced state
<——
89
Is it possible to measure the electrode potential of an individual half cell?
No
90
How does the electrode potential of a half cell need to be measured?
Against another half cell, usually a standard cell
91
What is usually chosen as the standard cell to measure the electrode potentials of half cells against?
The standard hydrogen electrode
92
What is the standard electrode potential of the hydrogen electrode under standard conditions?
E° = 0.0V
93
Worded description of the standard hydrogen electrode
Hydrogen gas at atmospheric pressure is passed over a platinum electrode in solution which contains hydrogen ions, hydrologic acid, at a concentration of 1moldm3 and a temperature of 298K
94
What type of acid is generally used in a standard hydrogen electrode and why?
Hydrochloric acid
The concentration of H+ ions [H+] is 1moldm-3
95
Why is the platinum electrode used in the standard hydrogen electrode?
It’s inert
It allows current to flow in and out of the half cell
96
What concentration of sulfuric acid should be used if it’s used in a standard hydorgen electrode instead of hydrochloric acid? Why?
0.5moldm-3
We want 1moldm-3 of H+ ions
(using 1moldm-3 would give 2moldm-3 H+ ions —> H2SO4)
97
Hydrogen electrode as a redox cell diagram
Pt I H2 (g) I 2H+ (aq)
98
What does each half cell have that’s unique?
Voltage
99
How are electrode potentials usually written (use hydrogen as an example)?
H+ (aq)/1/2H2 (g)
100
How are half cells connected to eachother? What does this do?
With a high resistance voltmeter
A salt bridge
Completes the circuit
101
What type of voltmeter is used to connect half cells to each other?
High resistance
102
What is required to make contact with the external circuit if the reduced state is gaseous or aqueous?
A platinum electrode
103
What does a more positive E° value mean in terms of being oxidised/reduced?
Greater tendency to be reduced
104
What value of E° do the most powerful oxidising agents have?
The most positive
105
What does a more negative E° value imply in terms of oxidation/reduction?
Greater tendency to be oxidised
106
What value of E°do the most powerful reducing agents have?
The most negative
107
What does the electrochemical series show?
Lists the electrode potentials of common redox changes in order of numerical value
108
What lists the electrode potentials of common redox changes in order of numerical value?
The electrochemical series
109
What tend to be reducing agents?
Metals
110
What’s the most powerful reducing agent?
Lithium
111
What tend to be oxidising agents?
Non-metals
112
What’s the most powerful oxidising agent?
Fluorine
113
What does the LHS of an electrode potential reaction show?
The oxidising agents
114
What does the RHS of an electrode potential reaction show?
Reducing agents
115
In which direction is electron flow?
From the negative electrode to the positive electrode
116
Explain the flow of electrons in the copper and zinc cell
Copper half cell is more positive than the zinc
Zinc becomes the negative electrode ad electrons flow from it to the positive copper electrode
(Electron flow is from the negative electrode to the positive electrode)
117
Which species will be th most dissociated?
The species with the most negative emf
118
Why is electron flow from the negative electrode to the positive electrode?
Species with the most negative emf will be the most dissociated
More atoms lose electrons and go into solution
The electrode will become negative
Species with the most positive emf gains electrons
119
Which species gains electrons
With the most positive emf
120
In which situation would the hydrogen electrode become the positive electrode for electron flow?
If the E° is negative, the electrode potential is more negative than the standard hydrogen half-cell
121
What do all electron flow reactions occur as? Why?
Redox reactions
One species loses electrons, one gains electrons
122
How does a species with a negative emf make the electrode negative?
Most negative emf = most dissociates
More atoms lose electrons and go into solution
The electrode will become negative
123
Does the species with the most negative emf gain or lose electrons?
Lose
124
Does the species with the most positive emf gain or lose electrons?
Gains
125
What happens to the species with the most negative emf?
Is oxidised
126
What happens to the species with the most positive emf?
Is reduced
127
What type of agent is the species with the most negative emf?
Reducing
128
What type of agent is the species with the most positive emf?
Oxidising
129
Which rule can help us work out the overall reaction within the cell?
Anti-clockwise rule
130
What does the anti-clockwise rule do?
Helps us work out the overall reaction within the cell
131
How does the anti-clockwise rule work?
1.) write the equations for the 2 half-cells under each other with the most negative at the top
2.) draw arrows starting at the top right hand corner in an anti-clockwise direction, one for each half equation
132
Electrochemical equations
Include the emf
133
Which type of equations are used for the anticlockwise rule?
Electrochemical equations
134
What can be used to determine the voltage of any cell?
Electrode potentials (E°)
135
Equation for working out the voltage of any cell
Ecell (V) = E°(reduction) - E°(oxidation)
Or
Ecell (V) = E°(RHS) - E°(LHS)
136
How come the RHS - LHS method always works when working out cell voltages from E° values?
Oxidation is always shown first
137
How do we write balanced equations for reactions occurring in a cell?
Balance out the electrons so that they cancel out
138
What is the effect of changing concentration on E° values?
For a metal ion/metal electrode, the electrode potential will become less positive when the concentration is decreased, and conversely
139
What is proof that increasing the concentration makes electrode potential become more positive?
According to electrode potential values, manganese (IV) oxide should not be reduced by HCL
However, there is a reaction between them when the acid is concentrated
The E° value for MnO2/Mn2+ becomes more positive and reaction will occur
140
How can we use electrode potentials to predict whether or not a reaction takes place?
Negative E = no reaction
141
What are the stages or working out whether a reaction is possible?
1.) work out what’s been oxidised and what’s been reduced
2.) write out the cell that’s formed
3.) using E° values given, use the equation Ecell = E(reduction) - E(oxidation)
4.) if the cell voltage value is negative, there is no reaction.
142
Why would there be no reaction if the cell voltage value is negative? Give an example
The oxidising agent isn’t reactive enough to displace the other ions from its solution
E.g - iodine isn’t reactive enough to displace chloride ions from its solution and oxidise them
143
What must the cell voltage value be for the reaction to proceed in the direction indicated?
Positive
144
What’s a useful application of electrode potentials?
Use as a guide to the ease or difficulty of extracting a metal from its ore
145
How can we use electrode potentials as a guide to the ease or difficulty of extracting a met from its ore?
The vale of the electrode potential indicate on which side the position of equilibrium lies
146
What do positive electrode potential values imply in terms of position of equilibrium?
Position of equilibrium is to the right
147
What do negative electrode potential values imply in terms of position of equilibrium?
Position of equilibrium is to the left
148
How can we know if a metal occurs naturally on earth?
The value of the electrode potential is positive
149
How can we know if a metal is found combined?
The electrode potential value is negative
150
Method of extracting a combined metal with a SMALL negative value of electrode potential
Carbon reduction (e.g - iron)
151
Method of extracting a combined metal with a LARGE negative value of electrode potential
Electrolysis (e.g - aluminium)
152
Which terminal of the voltmeter musts the hydrogen electrode be connected to?
The positive terminal
153
What happens to a hydrogen half cell when the reading on the voltmeter is negative?
Undergoes reduction
Electrons flow towards the hydrogen half cell
154
What does the anticlockwise rule state?
The half-reactions that occur are those that follow anticlockwise paths
155
How do we deduce the reaction taking part in a cell?
Use the anti-clockwise rule, ensuring that electrons are balanced in the half equations first
156
How do we work out the emf of a reaction
Work out oxidations and reductions
Ecell = E(reduction) - E(oxidation)
157
What is a fuel cell?
A device for producing electricity from external sources of fuel (on the anode side) and oxidant (on the cathode side
158
Which side produces electricity in a fuel cell?
The anode side
159
What is produced on the cathode side of a fuel cell?
Oxidant
160
What do fuel cells react in the presence of?
An electrolyte and a catalyst
161
What is the electrolyte in a fuel cell usually made from?
Platinum
162
Example of a fuel cell
Hydrogen cars
163
How do fuel cells work?
Generally, the reactants flow in and reaction products flow out while the electrolyte remains in the cell
164
Can fuel cells operate continuously?
Yes, as long as the necessary flows are maintained (e.g - topped up wit hydrogen)
165
Difference between fuel cells and batteries
Fuel cells consume reactant, which must be replenished
Batteries store electrical energy chemically in a closed system
166
How do batteries store electrical energy?
Chemically, in a closed system
167
Describe the electrode of fuel cells
Relatively stable
168
Fuel in a hydrogen cell
Hydrogen
169
Oxidant in a hydrogen cell
Oxygen
170
Catalyst in a hydrogen cell
Platinum
171
Describe the process within a hydrogen cell
1.) the fuel cell passes the fuel over platinum metal which acts as a catalyst, but also as an electrode for the electrochemical system
2.) electrons are removed from the hydrogen atoms at one electrode
3.) the protons (H+) diffuse through a semipermeable membrane to the other electrode where they gain electrons and oxygen molecules to form water molecules
172
Cell diagram of a hydrogen cell
Pt I H2/2H+ II O2/2H2O I Pt
Oxidation ——> reduction
173
Reaction taking place at the anode of a hydrogen cell
H2 ⇌ 2H+ + 2e-
174
Reaction taking place at the cathode of a hydrogen cell
O2 + 4H+ + 4e- ⇌ 2H2O
175
Reaction taking place in the hydrogen cell
2H2 + O2 ⇌ 2H2O
176
Benefits of the hydrogen fuel cell
It offers clean technology as water is the only product - no CO2 = no greenhouse gas emissions
It affords a convenient method storing energy
It is more efficient than the internal combustion engine (36% - 45% for fuel cell to 22% for diesel)
177
Drawbacks of the hydrogen fuel cell
Problems to do with storing gases due to the explosive nature of hydrogen
Energy is lost as the storage cycle is not 100% efficient
Hydrogen doesn’t exist naturally alone - it must be removed from being combined with other elements, or example via electrolysis (using electricity to split H2O to O2 and H2). This is likely to us fossil fuel energy sources, which will cause their own CO2 emissions
178
What must be true so that using the hydrogen fuel cell to generate electricity is green ?
The electricity for electrolysis to obtain the hydrogen must be renewable
179
Emf of a hydrogen fuel cell
+1.23V
180
Why is the emf of the hydrogen fuel cell lower in practice?
Standard conditions may not have been used
Storage cycle isn’t 100% efficient
181
Equation for energy given out
Energy of process x moles
182
Identifying a reducing agent from Ecell values
The most negative
183
Identifying an oxidising agent from Ecell values
The most positive
184
Which element do we always say that an electrode has been made from?
Platinum
185
Explain, using standard electrode potentials, why acid could be used to move iron form copper which is contaminated with iron
Ecell value of reacting iron with acid is positive, so there’s a reaction
Ecell value of reacting copper with acid is negative, so there’s no reaction
186
Which electrode will be the most positively charged?
The one with the most positive electrode potential value
187
Colour change in Cr2O7^2- + 2OH- ——> 2CrO4^2- + H2O
Orange to yellow
188
How to make a salt bridge
By soaking filter paper in potassium nitrate or chloride in solution
189
How do we represent ‘dilute acid’ in an equation?
H+
190
Give the ionic equation for the reaction of iron with dilute acid
Fe (s) + 2H+ (aq) —> Fe2+ (aq)+ H2
191
Chemical test to show that a solution contains iron (II) ions - what is the test, observations and ionic equation
Add NaOH
Green precipitate
Fe2+ (aq) + 3OH- (aq) ——> Fe(OH)3 (s)
192
What does a higher value of electrode potential mean and why?
Higher value = more positive standard electrode potential = stronger oxidising agent
