topic nine Flashcards
1
Q
define oxidation
A
gain of oxygen
loss of hydrogen
loss of electrons
increase in oxidation state
2
Q
define reduction
A
loss of oxygen
gain of hydrogen
gain of electrons
decrease in oxidation state
3
Q
two half equations should add together to create
A
the original ionic equation
4
Q
do oxidising or reducing agents accept electrons
A
oxidising
5
Q
do oxidising or reducing agents donate electrons
A
reducing
6
Q
If an atom loses control over electrons, it has been waht
A
oxidised
7
Q
If an atom gains control over electrons, it has been what
A
reduced
8
Q
which part comes first in oxidations tate
A
symbol then number
9
Q
what does oxidation state indicate
A
the gain or loss of electron control of an atom during a reaction
10
Q
oxidation states of oxygen, chlorine, nitrogen
A
free elements like these are zero
11
Q
the sum of the oxidation states of all the atoms in a compound mustbe equal to the
A
net charge on the compound
12
Q
alkali metals oxidation states
A
always +1
13
Q
fluorine oxidation states
A
always -1
14
Q
alkaline earth metals oxidation states
A
+2
15
Q
hydrogen oxidation state
A
always +1 unless with certain metal hydrides where it is -1
16
Q
oxygen oxidation states
A
-2 except with perioxides where it is -1, or with fluroine where it is +2
17
Q
ch;loride oxidation state
A
-1 unless with oxygen or fluorine
18
Q
charge on a metal ion is its
A
oxidation state
19
Q
average oxidation state
A
the average of the oxidation states of the same atom in a compound.
20
Q
how are roman numerals used to show the oxidation number
A
Roman numerals are used to show the oxidation number when there is more than one possible oxidation state for an atom in a compound. For example, in copper(I) oxide, the copper has an oxidation number of +1, whereas in copper(II) oxide, it is +2.
21
Q
when does disproportion occur
A
when the same species is oxidised and reduced simultaneously during a reaction to form two different products
22
Q
activity series
A
23
Q
explain the low reactiviy of aluminium
A
the formation of an extremely thin protective layer of aluminium oxide on the surface of the metal prevents the metal underneath from oxidising further.
24
Q
Any metal above hydrogen will react with an acid to produce
A
a salt and hydrogen gas
25
any metals below hydrogen will
not react with dilute acids
26
when can metals displace hydrogen ions from solution to produce hydrogen gas
when the metals are above hydrogen in the reactivity series
27
how does difficulty of extraction correlate with reactivity
the more reactive metals are much harder to extract
Metals above carbon in the activity series, for example, magnesium and aluminium, cannot be extracted from their ores by reduction with carbon. Conversely, metals below carbon in the series, such as iron and zinc, can be extracted by reduction with carbon. The reaction of zinc oxide with carbon is shown below
28
how does the reactivity series correlate with reactions with dilute acid
the more reactive metals produce H2 with decreasing vigour, and the less reactive dont react with dilute acids
29
how does the reactivity series correlate with reactions with air/oxygen
the more reactive metals burn very brightly and vigorously. burn to form an oxide with decreasing vigour.
the middlely reactive elements react slowly to form the oxide
the low reactive elements dont react
30
how does the reactivity series correlate with reactions with water
produce hydrogen with decreasing vigour with cold water. react with steam with decreasing vigour
less reactive elements dont react with cold water or steam
31
displacement reaction
the more reactive metal displaces the ions of the less reactie metal from solution. they are redox reactions which involve the transfer of electrons as metal atoms are oxidised and aqueous metal ions are reduced
32
spectator ion
an ion appearing on both sides of the equation and does not undergo oxidation or reduction
33
how does iron change when rusting
The precipitate of iron(II) hydroxide formed is rapidly oxidised under basic conditions to form red-brown hydrated iron(III) oxide.
34
what is the formula of rust
Fe2O3.xH2O
34
why does the rust form flakes that flake off the surface of the iron
it has a lower density than iron
34
sacrifical protection
blocks of magnesium or zinc are used to protect iron pipes, ships and storage tanks
35
galvanisation
a layer of zinc is coated over all the iron surface to give complete coverage.
36
drawback of tin plating
it is easily scratched to expose the iron surface. In this case the can rusts very rapidly. Once the iron is brought into contact with moist air, it gives sacrificial protection to the less reactive tin.
37
winkler method is used to
determine the biochemical oxygen demand or bod
38
what is bod
the amount of dissolved oxygen required to biologically decompose the organic matter in a water sample over a set time period (usually five days). Polluted water with a high BOD without the means of replenishing oxygen will not be able to sustain aquatic life. This can lead to the growth of green algae that feed on excessive nutrients in the water, known as eutrophication
39
when does eutrophication occur
when a body of water beocmes enriched in dissolved nutrients that lead to excessive growth of aquatic platn life such as algae
40
COMPLETE WINKLER METHOD
41
what is a redox titration
one that involves a redox reaction between the analyte and titrant
42
what usually happens in a redox titration
A common experiment using a redox titration is to determine the mass of iron in an iron tablet (iron tablets are prescribed to people who have low iron levels in their blood). In this experiment, an iron tablet is first dissolved in an acidic solution and the iron in the tablet is converted to the iron(II) ion, Fe2+. Next, the solution is titrated with an oxidising agent which is usually acidified potassium manganate(VII) or potassium dichromate(VI).
the iron(II) ion is oxidised to the iron(III) ion:
Fe2+ (aq) → Fe3+ (aq) + e–
the manganate(VII) ion, MnO4– is reduced to the manganese(II) ion:
MnO4– (aq) + 8H+ (aq) + 5e– → Mn2+(aq) + 4H2O (l)
The balanced equation in acidic solution is as follows:
5Fe2+ (aq) + MnO4– (aq) + 8H+ (aq) → Mn2+(aq) + 5Fe3+ (aq) + 4H2O (l)
43
what are voltaic cells aka
galvanic cells or batteries
44
what do voltaic or galvanic cell suse to produced an ee;ctric current
a redox reaction
45
what is a danielll cell
an early example of a simple voltaic cell
46
what does a daniell cell consist of
two half-cells: a zinc electrode in a solution of zinc sulfate, and a copper electrode in a solution of copper(II) sulfate. The two electrodes are connected by an external circuit with a voltmeter to measure the voltage. The voltage produced by a voltaic cell depends on the difference in reactivity between the two metals in the half-cells. The circuit is completed by a salt bridge that allows ions to flow in order to complete the circuit and prevent build-up of electric charge. Electrons flow through the wires of the external circuit in a spontaneous process; no external energy source is required.
47
what does a half cell consist of
an electrode, usually a metal, in a solution of its own ions. For example, a zinc half-cell is composed of a piece of solid zinc metal in a solution of aqueous zinc (Zn2+) ions.
48
how to make a salt bridge
a strip of filter paper soaked in saturated potassium nitrate (KNO3), although more complex salt bridges can be made using a glass tube filled with agar gel.
49
why is potassium nitrate chosen to make a salt bridge with
its constituent ions do not react with the other ions present in the two half-cells, or with the electrodes
50
what happens in the salt beidge
anions (negative ions) migrate from the salt bridge to the anode and cations (positive ions) migrate from the salt bridge to the cathode. This migration of ions prevents the build-up of electric charge in the two half-cells.
51
in a voltaic cell, which mental is oxideised and which is reduced
The more reactive metal is oxidised, and the less reactive metal is reduced
52
what creates a potential difference
The difference in the tendency to either undergo oxidation or reduction creates a potential difference between the two half-cells, also known as the electromotive force (EMF) or cell potential (E⦵cell).
53
what is cell potential
The cell potential (E⦵cell), also known as the electromotive force (EMF), is the difference in the electrode potentials of the half-cells in a voltaic cell. The cell potential of a voltaic cell is measured in volts (V).
54
in a voltaic cell, what happens at the anode and cathode
the anode is where oxidation takes place and is negatively charged. The cathode is where reduction takes place and is positively charged.
55
what does teh voltage of a cell in a voltaic cell epend on
The voltage of a cell depends on the nature of the electrodes and the ions involved, the temperature and the ion concentrations of the solutions.
56
the more reactive metal forms the
anode
57
cell diagram conventiosn have
A phase boundary between a solid and an aqueous solution is represented by a single vertical line.
A salt bridge is represented by a double vertical line.
The anode is usually placed on the left and the cathode on the right.
Two types of ions in the same half-cell are separated by a comma.
Spectator ions are not included in the diagram.
58
explain the electrical conductivity of ionic compounds when molten or dissolved in aqueous solutions
mobile ions move in a particular direction in an electric field
59
what is electrolysis
the process by which a compound is broken down into its constituent elements using electricity.
60
what is an electrolytic cell composed of
a molten or aqueous electrolyte
a battery and two electrodes, the anode and cathode.
61
what are electrodes usually made of
graphite (carbon) or an unreactive metal such as platinum
62
which terminal is the anode conneced to and which terminal is the cathode connected to
The anode is connected to the positive terminal of the battery and the cathode is connected to the negative terminal.
63
in an electrolytic cell, where does oxidation occur
at the positive anode
64
in an electrolytic cell, where does reduction occur
negative cathode
65
what type of reactiona t a voltaic cell
Involves an exothermic spontaneous redox reaction
66
what type of reaction at an electrolyic cell
endothermic non-spontaneous redox reaction
67
what does a voltaic cell convert into what
Cell converts chemical energy into electrical energy
68
what does an electrolytic cell convert into what
Cell converts electrical energy into chemical energy
69
which electrode is pos or neg in the voltaic cell
The cathode is positive and the anode is negative during discharge
70
which electrode is pos or neg in the electrolytic cell
The cathode is negative and the anode is positive during electrolysis
71
what does a voltaic cell use
two separate aqueous solutions connected by a salt bridge and an external circuit
72
what does an electrolytic ccell electrolye consist of
The electrolyte is a molten liquid (or an aqueous solution)
73
how is electric current conducted in a voltaic cell
Electric current is conducted by the electrons in the external circuit and the movement of ions in the salt bridge
74
how is electric current conducted in an electrolytic cell
Electric current is conducted by the electrons in the external circuit and the movement of ions in the electrolyte
75
When the two half-cells are connected together, the electrons flow from... to...
from the half-cell with the more negative standard electrode potential (E⦵) value to the half-cell with the more positive standard electrode potential value
76
The voltage generated by a voltaic cell depends on
the difference between the values of the standard electrode potentials of the two half-cells.
77
what is the standard electrode potential E⦵
the voltage produced when a half-cell is connected to a standard hydrogen electrode under standard conditions.
78
how is standard electrode potential measured
by connecting it to a standard hydrogen electrode (SHE)
79
conditions of the SHE:
hydrogen gas, H2 (g), at a pressure of 100 kPa and temperature of 298 K
acidic solution with [H+] of 1.0 mol dm–3
inert platinum electrode.
80
A standard hydrogen electrode (SHE).
81
The half-equation for the reaction that takes place at the surface of the platinum electrode is at an SHE is
2H+ (aq) + 2e– → H2 (g) or H2 (g) → 2H+ (aq) + 2e–
82
When E⦵ > 0, this half-cell will undergo
reduction when connected to a SHE. The electron flow will be from the SHE to the half-cell in question, which is the cathode of the cell.
83
When E⦵ < 0, this half-cell will undergo
when connected to a SHE. The electron flow will be from the half-cell to the SHE. This half-cell will be the anode of the cell.
84
which metals in the electrochemical series are stronger reducing agents
metals in the half cells at the top of the series with large negative electrode potential values have the greatest tendency to lose electrons and undergo oxidation and from positive ions in aqueous solution
85
what causes the electromotive force (EMF) or cell potential (E⦵cell)
The difference in the tendency to either undergo oxidation or reduction creates a potential difference between the two half-cells, also known as the electromotive force (EMF) or cell potential (E⦵cell).
86
SHE: When two half-cells are connected, the half-cell with the more negative standard electrode potential value undergoes
oxidation (is the anode of the voltaic cell)
87
When two half-cells are connected, he half-cell with the more positive standard electrode potential value undergoes
reduction (is the cathode of the voltaic cell)
88
what is the equation used to measure E⦵cell=
E⦵cell = E⦵half-cell where reduction occurs − E⦵half-cell where oxidation occurs
89
ΔG⦵ =
– nFE⦵
90
when E⦵ is positive, ΔG⦵ is negative and the reaction is
spontaneous
91
when E⦵ is negative, ΔG⦵ is positive and the reaction is
non spontaneous
92
when E⦵ = 0, ΔG⦵ = 0, the reaction is
at equilibrium
93
what is electrolysis
a process in which an electric current is used to break apart a compound into its constituent elements
94
Electrolysis involves a non-spontaneous reaction with what ecell values and what gvalues
negative E⦵cell value and a positive ΔG⦵ value.
95
where can electrolysis take palce
in eithe rmolten liquids or aqueous solutions of ioni compou ds
96
why is the electrolysis of aqueous solutions of ionic compounds is more complicated than that of molten ionic compounds
because the water itself can be electrolysed.
97
equation for how water can be pxidised at the anode
Oxidation: 2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
98
equation for how water can be reduced at the cathode
Reduction: 2H2O (l) + 2e– → H2 (g) + 2OH– (aq)
99
The process of selective discharge occurs when
one species is preferentially discharged over another at either the anode or cathode of an electrolytic cell.
100
When there is more than one possible species that could be discharged at the electrodes, the choice of species that is actually discharged depends on three factors:
The standard electrode potential (E⦵) value of the half-equations for the reaction.
The relative concentrations of the solutions (such as for the halogens in aqueous solution).
The material that the electrodes are composed of (such as the use of graphite or copper electrodes in the electrolysis of copper(II) sulfate).
101
at the anode, with a dilute ocnc of sodium chloride, what is oxidised
H2O (l) is oxidised to form oxygen gas:
2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
102
at the anode, with a saturated conc of sodium chloride, what is oxidised
chloride ions are oxidised to form chlorine gas:
2Cl– (aq) → Cl2 (g) + 2e–
103
at the cathode, with sodium chloride solution, waht are the two possible reactions
The two possible reactions at the cathode are the reduction of sodium ions or the reduction of water:
Na+ (aq) + e– → Na (s) E⦵ = –2.71 V
2H2O (l) + 2e– → H2 (g) + 2OH– (aq) E⦵ = –0.83 V
H2O is a stronger oxidising agent than Na+ ions (H2O has a more positive E⦵ value), meaning that it is more easily reduced. Therefore, at the cathode, water undergoes reduction to form hydrogen gas.
104
The electrolysis of brine (NaCl) produces
hydrogen and chlorine gases, but the product of greatest economic importance is the alkali, sodium hydroxide.
105
where is sodium hydroxide produced in sodium chloride electrolysis
Note that sodium hydroxide is not produced at either the cathode or the anode of the cell, instead, it exists as the solution formed in the reaction.
106
why is pure water a poor conductor of electricity
it lacks any of the mobile ions in solution
107
what do you add to pure water to allow for its electrolyisis
To increase its conductivity and allow the solution to conduct electricity, a strong electrolyte such as sulfuric acid is added, effectively creating a dilute sulfuric acid solution. When very dilute sulfuric acid is electrolysed, one volume of oxygen gas is collected at the anode and two volumes of hydrogen gas are collected at the cathode
108
electrolyis of dilute sulfuric acid at the anode:
At the anode, water is oxidised to produce oxygen gas
Anode: 2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
109
electrolyiss of dilute sulfuric acid at the cathode
at the cathode, water is reduced to produce hydrogen gas:
Cathode: 4H2O (l) + 4e– → 2H2 (g) + 4OH– (aq)
110
The overall equation for the electrolysis of water is:
2H2O (l) → 2H2 (g) + O2 (g)
111
describe the changes in pH for the electrolyis of waater
The changes in pH that occur at the electrodes are as follows: at the anode, the pH decreases as H+ ions are produced; at the cathode, the pH increases as OH– ions are produced.
112
The electrolysis of copper(II) sulfate produces
pure copper at the cathode.
113
At the anode of the electrolysis of copper(II) sulfate
water is oxidised to form oxygen gas:
2H2O (l) → O2 (g) + 4H+ (aq) + 4e–
114
electrolysis of copper sulfate, If the graphite electrodes are replaced by copper electrodes, the reaction taking place at the cathode remains the same – copper(II) ions are reduced – but at the anode the copper electrode itself undergoes oxidation:
(give the equation)
Cu (s) → Cu2+ (aq) + 2e−
115
why is the electrolysis of copper sulfate used
in the refinement of impure copper
116
overall equation for the electrolyiss of copper sulfate
Cu (s) + Cu2+ (aq) → Cu2+ (aq) + Cu (s)
117
what does electrplating invovle
the coating of a piece of metal with a thin layer of another metal. Metal objects are electroplated to prevent corrosion or to improve their appearance. Copper, chromium, silver and tin are the most commonly used metals for electroplating. Examples of electroplated objects include jewellery and cutlery
118
silver electroplating at the cathode
silver ions undergo reduction to form silver atoms which coat the spoon:
Ag+ (aq) + e− → Ag (s)
119
silver electroplating at the anode
the silver atoms that make up the piece of silver undergo oxidation:
Ag (s) → Ag+ (aq) + e−
120
The object to be electroplated is
the cathode.
121
The anode is composed of the metal to be
plated onto the obect
122
electroplating: The electrolyte solution is a salt containing the ions of
the metal to be plated onto the object.
123
Q =
It
where: Q = charge (in coulombs)
I = current (in amps)
t = time (in seconds)
124
One mole of electrons has a charge of
96 500 coulombs.
125
From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:
if the current is doubled,
the charge is doubled
126
From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:
if the time is doubled
the charge is doubled
127
From the equation Q = It, we can see that charge (Q) is directly proportional to the current (I) and the time (t) for which the current is applied. For example:
and if both the curretn and time are doubled,
the charge will be four times greater
128
