Download Chemistry B2A Chapter 18 Oxidation

Chemistry B2A
Chapter 18
Oxidation-Reduction Reactions & Electrochemistry
Oxidation is the loss of electrons and Reduction is the gain of electrons.
Zn(s) + Cu2+(aq) → Zn2+(aq) + Cu(s)
Zn(s) → Zn2+ + 2eCu2+ + 2e- → Cu(s)
Zn is oxidized (loses electrons) so it is a reducing agent.
Cu2+ is reduced (gains electrons) so it is an oxidizing agent.
In some reactions, it is not easy to see the electron loss and gain, so chemists developed
another definition of oxidation and reduction:
Oxidation is the gain of oxygen atoms and/or the loss of hydrogen atoms. Reduction is the
loss of oxygen atoms and/or the gain of hydrogen atoms.
CH4(g) + 2O2(g) → CO2(g) + 2H2O(g)
Therefore, CH4 is the reducing agent and O2 is the oxidizing agent.
Note: When the oxidizing or reducing agent is named, the whole compound is specified, not
just the element that undergoes the change.
Oxidation states (oxidation numbers): it lets us keep track of electrons in oxidationreduction reactions by assigning charges to the various atoms in a compound. However,
sometimes these charges are quite apparent.
Rules for assigning oxidation states:
1. The oxidation state of an atomic in an uncombined element is 0.
Ar, Na, K, O2, and N2
2. The oxidation state of a monatomic ion is the same as its charge.
3. Oxygen is assigned an oxidation state of -2 in most of its covalent compounds.
Important exception: peroxides (compounds containing the O22- group), in which each
oxygen is assigned an oxidation state of -1. For example in H2O2, oxidation state for
oxygen is -1.
4. In its covalent compounds with non-metals, hydrogen is assigned an oxidation state of
+1.
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
5. In binary compounds, the element with the greater electronegativity is assigned a
negative oxidation state equal to its charge as an anion in its ionic compounds. It
means, the most electronegative atom controls or possesses the charge electrons.
For example, in SO2 oxygen has a higher electronegativity. The oxidation state of
oxygen is -2. Therefore, 2(-2) = -4 and sulphur should be +4.
Note: The most electronegative element are: F, O, N, and Cl.
F: -1
O: -2
N: -3
Cl: -1
Note: If two of these elements are found in the same compound, we assign them in
order of electronegativity.
F ˃ O ˃ N ˃ Cl
For example, NO2 Oxygen has a higher electronegativity. Therefore,
O: 2(-2) = -4
So, nitrogen should be +4.
6. For an electrically neutral compound, the sum of the oxidation states must be zero.
7. For an ionic species, the sum of the oxidation states must equal the overall charge.
Note: in some redox reactions ions are produced:
Note: in some redox reactions ions are not produced (redox reactions between non-metals).
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
Oxidation: is an increase in oxidation state (a loss of electrons).
Reduction: is a decrease in oxidation state (a gain of electrons).
Oxidizing agent (electron acceptor): the reactant containing the element that is reduced
(gains electrons).
Reducing agent (electron donor): the reactant containing the element that is oxidized (loses
electrons).
Half-reaction method for balancing:
1. Identify and write the equations for the oxidation and reduction half–reactions.
2. For each half–reaction:
A. Balance all the elements except H and O.
B. Balance O using H O.
2
+
C. Balance H using H .
D. Balance the charge using electrons.
3. If necessary, multiply one or both balanced half-reactions by an integer to
equalize the number of electrons transferred in the two half-reactions.
4. Add the half-reactions, and cancel identical species.
5. Check that the elemnts and charges are balanced.
Example: how can we balance the following equation?
1. Separate into half-reactions.
2. Balance elements except H and O.
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
3. Balance O’s with H2O and H’s with H+.
4. How many electrons are involved in each half-reaction? Balance the charges.
5. Multiply the whole reactions by a whole number to make the number of electrons gained
equal to number of electrons lost.
6. Combine half-reactions cancelling out those reactants and products that are the same on the
both sides, especially the electrons.
Galvanic cell: this apparatus generates an electric current by using a redox reaction. We also
call this device a voltaic cell or, more commonly, a battery.
A single galvanic cell consists of an oxidizing agent and reducing agent which are physically
separated by one or more electrolytes. Each side of the galvanic cell is known as a half-cell.
Each half-cell consists of an electrode (the metal of the half-reaction) and a solution
containing the corresponding cation of the half-reaction (see Figure 1).
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
Figure 1:Galvanic cell
Standard Reduction Potentials (Eº): reduction potential is a measure of the tendency of a
chemical species to acquire electrons and thereby be reduced. Reduction potential is measured
in volts (V), or millivolts (mV). Scientists have chosen one half-cell as a reference and
compared all other half-cells to it. The chosen is standard hydrogen electrode, (SHE):
2H+ + 2e- → H2
Eº = 0.00 V
[H+] = 1.0 M, P = 1 atm., T = 25ºC
Table 1 shows some of the extreme values for standard reduction potentials:
Table 1: Standard Reduction Potentials
Zn2+ + 2e- → Zn(s)
Eº = -0.76 V
Cu2+ + 2e- → Cu(s)
Eº = +0.34 V
The greater the tendency of the ion to gain electrons and undergo reduction, the less negative
(or more positive) the reduction potential of the ion. In the zinc/copper cell, the copper has a
greater tendency to undergo reduction than the zinc. Therefore:
Oxidation takes place at the zinc electrode, liberating electrons to the external circuit:
Zn(s) → Zn2+ + 2eReduction takes place at the copper electrode, consuming electrons coming from the external
Circuit:
Cu2+ + 2e- → Cu(s)
The electrode at which oxidation occurs is called the anode (Zn).
The electrode at which reduction occurs is called the cathode (Cu).
Since oxidation releases electrons to the anode, it is designated the negative electrode in the
galvanic cell (Zn). Reduction removes the electrons from the cathode; it is the positive
electrode. The electrodes of the half-cells are connected by a wire through which the electrons
can flow.
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
Salt bridge: the electrons flow over the wire from anode (Zn) to cathode (Cu). We cannot
have positive charges in any place without an equivalent number of negative charges. Salt
bridge is necessary to carry ions and keep the charge neutral. A salt bridge is usually an
inverted glass U-tube that connects two beakers together. The salt bridge is filled with a
strong electrolyte such as NaCl, KCl, or NaNO3.
Lead storage battery: they contain a liquid electrolyte.
Dry cell battery (acid version): they do not contain a liquid electrolyte.
Dry cell battery (alkaline version): they do not contain a liquid electrolyte.
Dry cell battery (other version): they do not contain a liquid electrolyte.
Dr. Behrang Madani
Chemistry B2A
Bakersfield College
Silver cell – Zn anode, Ag2O cathode
Mercury cell – Zn anode, HgO cathode
Nickel-Cadmium: Rechargeable
Electrolysis: the process of electrolysis involves forcing a current through a cell to produce a
chemical change that would not otherwise occur. Electric energy is converted to the chemical
reactions in electrolysis.
Note: electrolysis is a nonspontaneous process.
Dr. Behrang Madani
Chemistry B2A
Bakersfield College