Redox Flashcards
1
Q
what is a redox equation?
A
Redox reactions involve both an oxidation reaction occurring at the same time as a reduction reaction.
2
Q
what do redox equations involve?
A
Involve a transfer of electrons from one reactant species to another.
3
Q
what is used to help remember oxidation and reduction reactions?
A
OILRIG -
Oxidation
Is
Loss of electrons
Reduction
Is
Gain of electrons
4
Q
what is a reductant?
A
Undergoes oxidation, causes the OTHER REACTANT to undergo reduction, also called reducing agent.
5
Q
what is an oxidant?
A
Undergoes reduction, causes the OTHER REACTANT to undergo oxidation, also called oxidising agent.
6
Q
what are the steps for writing a half equation in acidic conditions?
A
1.Write redox couple separated by an →
2. Balance any atoms/ions (not O or H yet!)
3. Balance O by adding H2O to the other side
4. Balance H by adding H+ (acidic cond.) to the other side.
5. Check total charge on LHS and compare the the RHS
6. Add electrons (negative charge) to the appropriate side such that LHS = RHS. Charge does not need to be zero on both sides.
7
Q
how do you identify a reduction half equation?
A
+ electrons on the LEFT
8
Q
how do you identify a oxidation half equation?
A
+ electrons on the RIGHT
9
Q
what does + electrons on the right really mean?
A
remember that + electrons on the right, really means - electrons on the left. (OIL)
10
Q
what are the steps for combining half equation?
A
- Take note of no. of electrons on each side of the arrow in two half equations. If they are not equal then one/or both half equations need to be multiplied so they become equal.
- Combine two half equations so that ALL species on the LEFT of the arrow end up on the LEFT, and all species on the RIGHT of the arrow end up on the RIGHT.
- Cancel out electrons from both side as well as any other species that appear on both sides of arrow.
11
Q
what conditions are half equations under for this AOS?
A
acidic conditions
12
Q
what are oxidation numbers?
A
Oxidation numbers, sometimes referred to as oxidation states, are a tool we use to help keep track electrons.
13
Q
what are oxidation numbers used for?
A
They are used for both ionic and covalently bonded species, so do not always represent an actual physical/formal charge.
14
Q
what do oxidation numbers help us with?
A
They allow us to help ‘see’ where the transfer of electrons involved in redox reactions is occuring.
15
Q
what reaction has a decrease in oxidation number?
A
reduction reaction
16
Q
what reaction has an increase in oxidation number?
A
oxidation reaction
17
Q
what are the oxidation rules?
A
- The oxidation number for an atom in its elemental form is always zero.
- The oxidation number of a monoatomic ion = charge of the monatomic ion. *
- The oxidation number of all Group 1A metals = +1 (unless elemental).
- The oxidation number of all Group 2A metals = +2 (unless elemental).
- Hydrogen (H) has two possible oxidation numbers:
* +1 when bonded to a non-metal
* -1 when bonded to a metal - Oxygen (O) has two possible oxidation numbers:
* -1 in peroxides (O2 2-)….pretty uncommon
* -2 in all other compounds…most common - The oxidation number of fluorine (F) is always -1.
(Assign non-metal oxidation numbers starting from the element closest to the top right corner of the periodic table.) - The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.
- The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion.
18
Q
what is an electrochemical series?
A
This is a ranking of redox species, based on which one is the best gainer of electrons.
19
Q
what’s on the left side of electrochemical series?
A
oxidants
20
Q
where are the strongest oxidants found in the electrochemical series?
A
The strongest oxidants are found highest in the ranking.
21
Q
what’s on the right side of electrochemical series?
A
reductants
22
Q
where are the strongest reductants found in the electrochemical series?
A
The strongest reductants are found lowest in the ranking.
23
Q
what equation is read left to right in the electrochemical series?
A
reduction reaction
24
Q
what equation is read right to left in the electrochemical series?
A
oxidation reaction
25
what conditions does this electrochemical series occur in?
- solutions are 1 molL-1
- Temp is 25oC
- Pressure is 1atm
26
what might happen when a metal is placed in a aqueous solution containing metal ions?
a displacement reaction may or may not occur
27
what is a spontaneous reaction?
reaction DOES occur
28
what is a non-spontaneous reaction?
reaction DOES NOT occur
29
how can you determine is a reaction is spontaneous or not on the electrochemical series?
Identifying where the metal and metal ions are ranked on the electrochemical series can help make predictions about whether a spontaneous or non-spontaneous reaction will occur.
30
Downward slope between species?
spontaneous
31
Upward slope between species?
non-spontaneous
32
Metal reactivity - electrochemical series?
Reductants on the electrochemical series, are actually an upside down metal reactivity series.
Most reactive metals at the bottom and least reactive at the top.
33
what do spontaneous reactions in a single beaker have?
These spontaneous reactions, occuring in a single beaker, do have a transfer of electrons from reductant to oxidant, however this electron flow/movement occurs at the interface of the metal and solution and is not harness/utilised.
Chemical energy → Heat energy
34
what do spontaneous reactions in two beakers have?
Separating this reaction out into two beakers and connecting them with a wire, allows the flow/movement of electrons to be controlled and harness. This allows redox reactions to generate electricity.
Chemical energy → Electrical energy
35
Key components of a Galvanic cell?
- 2 x half cells
- External circuit, including wire and optional light bulb/voltmeter
- Salt bridge connecting half cells
36
what is a half cell?
a metal electrode sitting in its matching metal solution
37
what is a salt bridge?
filter paper soaked in ionic solution (eg. KNO3)
38
what does a galvanic cell have?
a galvanic cell has both halves of each redox couple present from the outset so a spontaneous reaction ALWAYS OCCURS!
39
how do you know what species undergoes the reaction in a galvanic cell?
the stronger oxidant/reductant undergoes the reaction
40
how do you know what electrode is an anode?
Oxidation occurs at Anode (-)
41
how do you know what electrode is an cathode?
Reduction occurs at Cathode (+)
42
why is a salt bridge important?
Salt bridge, sometimes called the internal circuit, plays an important role.
- It completes the electrical circuit (provides charged particles, free to move)
- Balances the changing charge in each half cell.
43
how do you know where the 2Na+/K+ goes in a galvanic cell?
the cell becoming less + over time - the reduction reaction
44
how do you know where the SO4 2-/NO3- goes in a galvanic cell?
the cell becoming more + over time - the oxidation reaction
45
what are some special half cells?
- solution half cell, where a graphite electrode (inert) is used to conduct electricity, when both of the redox couple are in solution
- gas-non metal ion half cell, graphite electrode (inert) and gas chamber used
46
what will happen in a galvanic cell if the voltmeter is placed correctly?
When a voltmeter is placed in the circuit correctly (+ terminal to + electrode) the needle will indicate voltage (electricity) is being produced.
47
how must a voltmeter be placed in a galvanic cell?
- + terminal to + electrode
- - terminal to - electrode
48
do all cells produce the same voltage?
Not all ½ cell combinations produce the same voltage.
49
what cells will produce higher voltage?
Selecting ½ cell combinations that are further apart on the electrochemical series produce greater voltage. This is because you’re using a stronger oxidant and a stronger reductant.
50
are all cells on the electrochemical series?
There may be times when one or both ½ cells used in a galvanic cell do not appear on the electrochemical series.
51
how do correct position of galvanic cells help with electrochemical series?
By correctly constructing the galvanic cell, you can determine where that redox couple should be placed in the electrochemical series.
52
how do you conduct a galvanic cell?
- Ensure the voltmeter is reading a positive voltage
- The -ve terminal of voltmeter is attached to the -ve electrode → ∴ this ½ cell is doing oxidation
- The +ve terminal of voltmeter is attached to the +ve electrode → ∴ this ½ cell is doing reduction
- The reduction redox couple will appear higher in the electrochemical series
53
colour of Cu 2+ (aq) (Copper (II) ion)
/ Cu(s) (Copper) ?
Cu 2+ - blue (can be green)
Cu - pink
54
colour of I2 (s) (Iodine) / I- (aq) (Iodide ion) ?
I2 - brown/black
I- - colourless
55
colour of Zn 2+ (aq) (Zinc ion)
/ Zn(s) (Zinc) ?
Zn 2+ - colourless
Zn - silver/grey
56
colour of Fe 3+ (aq) (Iron (III) ion)
/ Fe 2+ (aq) (Iron (II) ion) ?
Fe 3+ - orange/brown
Fe 2+ - pale green
57
observations of redox?
- colour change of solution
- electrode reduce/increase in size
- solid may start forming (electrode increase)
58
when can spontaneous and non spontaneous reactions occur?
This is true when there is only one species from each couple is present (eg. one species which can act as an oxidant, and one species that can act as a reductant.
59
colour of H+ (aq) (Hydrogen ion)
/ H2 (g) (Hydrogen) ?
H+ - Colourless
H2 - Colourless
60
colour of Br2 (l) (Bromine)
/ Br- (aq) (Bromide ion) ?
Br2 - Orange
Br- - Colourless
61
colour of Fe2+ (aq) (Iron (II) ion)
/ Fe (s) (Iron) ?
Fe2+ - Pale green
Fe - Silver/grey
62
colour of Mg2+ (aq) (Magnesium ion)
/Mg (s) (Magnesium) ?
Mg2+ - Colourless
Mg - Silver/grey
63
colour of O2 (g) (Oxygen) /H2O2 (l) (Hydrogen peroxide) ?
O2 - Colourless
H2O2 - Colourless
64
colour of Cl2 (g) (Chlorine)
/Cl- (aq) (Chloride ion) ?
Cl2 - Pale green
Cl- - Colourless
65
colour of O2 (g) (Oxygen) /O2-(varies) (Oxide ion) ?
O2 - Colourless
O2-(varies) - Colour varies
66
colour of MnO4-(aq) / H+ (Permanganate ion)
/ Mn2+(aq) (Manganese ion) ?
MnO4-(aq) / H+ - Purple
Mn2+(aq) - Colourless/pale pink
67
colour of Cr2O72-(aq)/H+ (Dichromate ion) / Cr3+(aq) (Chromium ion) ?
Cr2O72-(aq) / H+ - Orange
Cr3+(aq) - Green
68
colour of Mx+(aq) (Metal ion)
/M(s) (Metal) ?
Mx+(aq) - Most colourless
M(s) - Most silver/grey
69
70
define corrosion
Corrosion is ‘damage’ caused to a metal by a reaction as a result of exposure to chemicals in the environment.
71
what reaction is corrosion?
is an oxidation reaction
72
what does corrosion typically involve?
Corrosion typically involves a metal being converted into its metal ion.
73
what happens when corrosion occurs chemically?
The loss of electrons from a metal is an oxidation half equation.
74
is corrosion wanted?
Corrosion is not desirable, in fact billions of dollars are spent annually trying to prevent it.
75
what does corrosion result in?
Corrosion results in metals losing their lustre/shine, results in breakdown of the metal, affects it strength and structural integrity.
76
what does corrosion cost?
Corrosion is thought to cost Australia $32 billion per year in lost infrastructure.
77
what is dry corrosion?
Corrosion that occurs in the presence of oxygen but no water is called dry corrosion. .
78
what reaction is dry corrosion?
redox reaction -
metal + oxygen → metal oxide
79
dry corrosion as redox?
Cu → Cu2+ + 2e- oxidation
O2 + 4e- → 2O2- reduction
Full/combined: 2Cu + O2 → 2CuO
80
what is the speed of dry corrosion?
This reaction is slow.
81
what is wet corrosion?
Wet corrosion occurs in the presence of both oxygen and water.
82
what is wet corrosion of iron?
The wet corrosion of iron is often known as rusting. Don’t use “rusting” for other metals.
83
what does corrosion require?
Corrosion often requires a ‘stress point’ where protection of the metal has been removed (eg. chipped paint).
84
what is the site where oxidation occur?
Site where oxidation occurs is referred to as the anode.
85
what is the site where reduction occur?
Site where reduction occurs is referred to as the cathode.
86
complex reactions in wet corrosion of iron?
Fe → Fe2+ + 2e- -oxidation
O2 + 2H2O + 4e- → 4OH- -reduction
Then.... Fe2+ + 2OH- → Fe(OH)2(s) (pale green)
Fe(OH)2 is unstable and continues to react with O2 & H2O → Fe(OH)3(s) (orange/brown)
Over time the Iron (III) hydroxide precipitate forms the hydrated Iron (III) oxide compound
Fe2O3.xH2O(s) value of x alters the colour.
87
where does wet corrosion lie on the electrochemical series?
Notice the downward slope from the O2(g) + 2H2O(l) to the Fe(s), so this reaction is spontaneous.
88
what is the speed of wet corrosion?
Reaction is still slow but is a little quicker than dry corrosion.
89
Whole wet corrosion reaction?
4Fe + 302 + X H20 → 2Fe203 X H20
(hydrated iron (III) oxide)
90
are their different types of corrosion prevention?
There are a number of strategies employed to help avoid corrosion. Some methods are simple and cheap, while others are more expensive and complex.
91
what are the prevention methods for corrosion?
- barrier methods
- sacrificial methods (cathodic protection and Galvanising iron with zinc)
92
what are barrier methods?
Barrier methods are the most straightforward and simple. They aim to prevent oxygen and water coming in contact with the metal’s surface. eg:
- Paint
- Lubricants/oils/grease (particular if part is moving)
- Coating in a less reactive metal or metal that is a weaker reductant
93
what are sacrificial methods?
Corrosion prevention can also involve sacrificial methods. These are strategies which sacrifice a more reactive metal (stronger reductant) in order to protect the metal you’re interested in.
94
what is galvanising iron with zinc?
Galvanising iron with zinc helps preserve the iron. Zinc is more likely to undergo oxidation (stronger reductant) than iron. So when oxygen and water are present, zinc will corrode in preference to iron. The zinc is used, knowing that it will corrode and be sacrificed.
95
what is cathodic protection?
A more complex sacrificial method for protecting iron is called cathodic protection. A more reactive metal (stronger reductant) is again used, but in addition to this an electrolytic cell is set up. This involves creating an electric circuit which ‘forces’ electrons onto the metal you are trying to protect (cathode) ensuring reduction occurs. This protects the metal!
Fe2+ + 2e- → Fe Reduction
Electrons are removed/lost from the metal being sacrificed (anode), ensuring oxidation occurs
eg. Mg → Mg2+ + 2e- Oxidation
