4.2 Liderazgo Educacional: 4.2.1 Conceptos
4.4 Los valores y la educación
4.4.1 Los valores
4.4.1.2. Características de los valores
The product at the electrodes during electrolysis depend on a number of factors (a) Nature of the electrode
Certain electrodes take part in electrolytic reactions and must be used with caution. Platinum and graphite electrodes are inert electrodes and do not take part in electrolytic reactions. They are safer to use to avoid unwanted results.
(b) Concentration of the ions of the electrolyte
Increase in the concentration of any ion in the electrolyte increases the chance of its discharge especially when they are close in the electrochemical series. The electrochemical series gives the order of discharge of ions from solution. This factor is important for mixed electrolyte or for aqueous solutions of electrolytes.
Recall that water also ionizes to give H+ and OR. Water is a weak electrolyte. e.g for NaCc q) Cations are Na+ +H 3 0+
Anions—÷CY OR
Na+ is too far from H+ on the electrochemical series so Ft is discharged as F1 2(g) at the cathode and C12 at the anode.
(c) Position on the electrochemical series
K+ F-
Na+
ca2+
Mgr*
AIM Increasing ease zu2+ for discharge To Cathode 4--
Fe2+
Sn2+
Pb 2+
Increasing ease for Fit
discharge cu2+
Hr
Ag+
Au+
Pt+
Fig 9.2 Electrochemical/Reactivity Series 148
so42-
CI- Br-
OH-
To Anode
All other factors being constant, ions that are lower is the series will be discharge in preference to any other above in the series. The electrochemical series is the arrangement of cations and anions in the order they are discharged during electrolysis.
9.2.4 Examples of electrolysis of some salts
Let
us illustrate what we have discussed so far with examples: the electrolysis of copper (ii) tetraoxosulphate (iv) and sodium chloride solutions using platinum or carbon electrodes. The ions present in sodium chloride solutions are Nat, Hp+ which both migrate to the cathode and Ct and OW which migrate to the anode.
At
the cathode H, 0 ' is preferentially discharged and at the anode the CI is discharged +e
H + H Cl CI + Cl
► H 2
0 + H are cathode reactions
► H ag,
CI +e are anode reactions
l■ CI, g
H21g) is given off at the cathode and chlorine gas at the anode.
The cations, present in CuSo 4 solution are Cu ,I-1,0+. Copper is lower than hydroden in the electrochemical series. It is discharged in preference to the hydrogen.
At the anode the anions are OH - and SO 4 2- . The OH- is much lower in the series than SO 4 1- . It is discharged in preference to the SO 4 1- . The result is deposition of copper at the cathode and liberation of oxygen at the anode.
CCP •
ay)+ 2e Cu(s)
40H- ► °2(g)
+ 2H 2 Ofi,
Some more examples of electrolysis of salts are given in the table 9.1 caw
Table 9:1 Products at the electrodes during the electrolysis of some salt
Salt Type of electrode Product at the
• Molten sodium chloride
• Very dilute sodium chloride solution
• Concentrated sodium chloride solution
• Concentrated sodium hydroxide solution
• Dilute sodium hydroxide solution
• Dilute tetraoxosulphate (vi) acid
cathode anode cathode anode
Carbon Carbon Carbon Platinum or carbon Carbon Carbon
Carbon Carbon Carbon Platinum or Carbon Carbon Carbon
H, H
H
2(,
H
2(g) H 2
(g)02(g)
°2(g)
0 g)
°2
4.)Using the information provided in the table, explain why the concentrations of H2SO4fagI and Na0H(aq) increase during electrolysis of dilute H 2 SO4() aq and Na0Hoql respectively.
9.3 Redox Reactions 9.3.1 Oxidation and reduction
Oxidation is defined as:
(a) addition of oxygen to a substance.
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(b) removal of hydrogen from a substance.
(c) loss of electrons.
(d) increase in oxidation number.
Oxidation numbers are numbers given to elements in the free and combined states according to a set of rules.
Oxidation number allows us to keep track of electrons during chemical reactions. It is the charge that an atom of the elements would have if both electrons of each bond were assigned to the more electronegative element. Elements are assigned zero oxidation numbers in their elemental forms.
Reduction is also defined as (a) removal of oxygen (b) addition of hydrogen (c) gain of electrons
(d) decrease in oxidation number
Oxidation and reduction reactions always occur together. Recall that during electrolysis oxidation occurs at the anode and reduction at the cathode. The substance that accepts electrons is the oxidizing agent while the one that donates the electrons is the reducing agent.
93.2 Assignment of oxidation number
Oxidation is defined in the last section. Now we want to recall the rule for assigning oxidation numbers and use same to assign oxidation numbers to some elements in compounds.
The rules are:
(i) all atoms of elements in uncombined state have zero oxidation number. e.g. Na in Na ys), H in H2(
(ii) in simple ions, e.g Na*, Z11 2 ` AP' etc, the oxidation number is the charge on the ions.
(iii) the sum of the oxidation numbers must add up to zero for compounds and to the net charge on the ions if it is a complex ion
(iv) hydrogen and all alkali metals have +1 oxidation number in their compounds. Hydrogen has a —1 oxidation number in hydrides
(v) the oxidation number of oxygen is —2 in its compound except in peroxide where it is —I.
We now want to apply the rules to simple examples. H 2( (H), NaCI(C1), H 2 SO4(S), HC104(CI) and S0 32- (5). Assign oxidation numbers to the elements in bracket in the given compounds or ions. For H 2 , the oxidation number of H is zero (rule i)
NaCI — the rule assigns +1 to Na, CI must therefore have an oxidation number of —1 (rule ii) H2 SO4 (S). Let the oxidation number of S be x.
+2 + x — 8 = 0 (rules (iii) iv vi) x = + 6
Oxidation number of sulphur in H 2 SO4 is +6. This is why the IUPAC name is tetraoxosulphate (vi) acid.
Now complete by assigning oxidation numbers for Cl in HCIO 4 and S in S0 3 2- 9.3.3 Redox reactions and the electrochemical cell
In the previous section you learnt that oxidation-reduction reactions occur together and they involve loss and gain of electrons respectively. It is possible to have an arrangement in which the electron formed by the oxidation process are made to flow through a conductor (copper wire) to the reduction site. This flow of electrons will constitute an electric current a demonstration of the conversion of chemical energy to electrical
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energy. This type of arrangement is called an electrochemical cell or galvanic cell. An example of a galvanic cell is the Daniel cell which makes use of the following redox reaction.
CuSO4() Zn(s) --r-ZnS0 4(aq) + Cu( )
•■-- Voltmeter +Ve zinc rod
anode copper rod
cathode ZnSO4(aq)
- -
Cu SO4(aq) -ve
Fig 9.3 Electron transfer in a redox reaction
For a 1.0 mol dm-3 solutions of the electrolytes solutions at 298k the voltmeter will record about 1.10 volts.
The reactions that occur are Zn Zn + 2e - at the anode Cue' + 2e --F Cu(s) cathode.
Note here again that oxidation is at the anode and reduction at the cathode. 1.10 volts is called the electromotive force (emf) of the cell. The emf is the electric potential difference between the electrodes.
9.3.4 Comparison of the electrolytic and electrochemical cells
This comparison is given in the table 9.2, the differences and similarities are highlighted.
Table 9.2: Comparison of the electrolytic and electrochemical cells
Electrolytic Cell Electrochemical Cell
• Chemical decomposition occurs by passage Electricity is generated by chemical reaction of electricity through an electrolyte solution within the cell.
• Electrical energy is converted to chemical Chemical energy is converted to electrical
energy energy
• The anode is the positive electrode The anode is the negative electrode
• Oxidation occurs at the anode Oxidation occurs at the anode
• The cathode is the negative electrode The cathode is the positive electrode
• Reduction occurs at the cathode Reduction occurs at the cathode
• Eelctrons flow from anode to cathode Electrons flow from anode to cathode 9.4 Conclusion
The transformation of energy from electrical to chemical and vice versa has important consequences that are beneficial to humanity. The application of electrolysis are mainly in industries and natural phenomena such
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as corrosion. You can now understand why torch batteries supply electrical energy. You will also appreciate the importance of electrolysis when we look at the applications in the next unit.