Review: Balancing REDOX reactions Balance the following redox reactions: 1. MnO4-(aq) + Cl-(aq) Mn2+(aq) + Cl2(aq) acidic solution
Electrochemistry Oxidation-Reduction reactions. LEORA (loose electron oxidation reducing agent) GEROA (gain electron reduction oxidizing agent)
2. H2O2(aq) + ClO2(aq) ClO2-(aq) + O2(g) basic solution 3. NO2-(aq)solution
Cu2+ + Zn+ Cr2O72-(aq) Cr3+(aq)
Zn2+ + Cu
NO3-(aq) acidic Reduction: Cu2+ + 2e- Cu Oxidation: Zn Zn2+ + 2e(Half-reactions of the redox process)
The two parts of the reaction can be physically separated. The oxidation reaction occurs in one cell . The reduction reaction occurs in the other cell. A cell is a compartment for the half-reaction The cell must contain all physical forms of the species involved Reduction half-reaction cell contains aqueous Cu2+ and solid Cu Oxidation half-reaction cell contains aqueous Zn2+ and solid Zn
ElectrochemistryThere are two kinds electrochemical cells. 1. Electrochemical cells containing in nonspontaneous chemical reactions are called electrolytic cells. 2. Electrochemical cells containing spontaneous chemical reactions are called voltaic or galvanic cells.
The combination of two cells (reduction and oxidation cell) is called an electrochemical cell
Electrical Conduction Metals conduct electric currents well in a process called metallic conduction. In metallic conduction there is electron flow with no atomic motion. In ionic or electrolytic conduction ionic motion transports the electrons. Positively charged ions, cations, move toward the negative electrode (cathode) Negatively charged ions, anions, move toward the positive electrode (anode)
Electrodes The surface in which oxidation or reduction halfreaction occurs is called the ELECTRODE. Electrode part of reaction: active electrode Electrode not a part of reaction: INERT electrode
Voltaic or Galvanic Cells Electrochemical cells in which a spontaneous chemical reaction produces electrical energy. Cell halves are physically separated so that electrons (from redox reaction) are forced to travel through wires and creating a potential difference. Examples of voltaic cells include:
Convention for electrodes (correct for either electrolytic or voltaic cells): Cathode : reduction Negative in electrolytic cells and positive in voltaic cells.
Anode : oxidation Positive in electrolytic cells and negative in voltaic cells.
Construction of Simple Voltaic Cells
The Zinc-Copper CellCell components for the Zn-Cu cell are:1. 2. 3. A metallic Cu strip immersed in 1.0 M copper (II) sulfate. A metallic Zn strip immersed in 1.0 M zinc (II) sulfate. A wire and a salt bridge to complete circuit
The Zinc-Copper Cell Short hand notation for voltaic cells. The Zn-Cu cell provides a good example.Single vertical bar: electrode Single vertical bar: electrode
Voltaic cells consist of two half-cells which contain the oxidized and reduced forms of a substances in contact with each other. A simple half-cell consists of: A piece of metal immersed in a solution of its ions. A wire to connect the two half-cells. And a salt bridge to complete the circuit, maintain neutrality, and prevent solution mixing.
The cells initial voltage is 1.10 volts
Zn | (1.0 M) Zn2+ || Cu2+ (1.0 M) | CuOxidation half-reaction Reduction half-reaction Double vertical bar : Salt bridge
In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).
The Copper - Silver Cell Cell components:1. 2. 3. A Cu strip immersed in 1.0 M copper (II) sulfate. A Ag strip immersed in 1.0 M silver (I) nitrate. A wire and a salt bridge to complete the circuit.
The Copper - Silver Cell
The Copper - Silver Cell These experimental facts demonstrate that Cu2+ is a stronger oxidizing agent than Zn2+.
The initial cell voltage is 0.46 volts.
In other words Cu2+ oxidizes metallic Zn to Zn2+. Similarly, Ag+ is is a stronger oxidizing agent than Cu2+.
Because Ag+ oxidizes metallic Cu to Cu 2+. Compare the Zn-Cu cell to the Cu-Ag cell The Cu electrode is the cathode in the Zn-Cu cell. The Cu electrode is the anode in the Cu-Ag cell. If we arrange these species in order of increasing strengths, we see that:
Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.
Standard Electrode Potential Potential means the tendency to transfer electrons To measure relative electrode potentials, we must establish an arbitrary standard. That standard is the Standard Hydrogen Electrode (SHE). The SHE is assigned an arbitrary voltage of 0.00000000 V To determine the ability of a reaction to transfer or accept electrons relative to the SHE
The Zinc-SHE Cell For this cell the components are:1. 2. 3. A Zn strip immersed in 1.0 M zinc (II) sulfate. The other electrode is the Standard Hydrogen Electrode. A wire and a salt bridge to complete the circuit.
The Zinc-SHE Cell
The initial cell voltage is 0.763 volts.
The cathode is the Standard Hydrogen Electrode. In other words Zn reduces H+ to H2.
The anode is Zn metal. Zn metal is oxidized to Zn2+ ions.18
The Copper-SHE Cell The cell components are:1. 2. 3. A Cu strip immersed in 1.0 M copper (II) sulfate. The other electrode is a Standard Hydrogen Electrode. A wire and a salt bridge to complete the circuit.
The Copper-SHE Cell
Uses of Standard Electrode Potentials Electrodes that force the SHE to act as an anode are assigned positive standard reduction potentials. Electrodes that force the SHE to act as the cathode are assigned negative standard reduction potentials. Standard electrode (reduction) potentials tell us the tendencies of half-reactions to occur as written.Potassium (K0 ) has greater tendency to give electrons (to be OXidized relative to SHE)
The initial cell voltage is 0.337 volts.
In this cell the SHE is the anode The Cu2+ ions oxidize H2 to H+.
The Cu is the cathode. The Cu2+ ions are reduced to Cu metal.20
Fluorine (F20 ) has greater tendency to accept electrons (to be REDuced relative to SHE)
Uses of Standard Electrode Potentials Standard electrode potentials are used to predict whether an electrochemical reaction (at standard state conditions) will occur spontaneously.1. Strongest Oxidizing Agent
Uses of Standard Electrode PotentialsChoose the appropriate half-reactions from a table of standard reduction potentials. 2. Write the half-reaction equation with the more positive E0 value first, along with its E0 value. 3. Write the other half-reaction equation as an oxidation (reverse the tabulated reduction half-reaction) and change the sign of the tabulated E0. 4. Balance the electron transfer. Do not multiply the E0 values by the coefficient! (E0 values are INTENSIVE properties) 5. Strongest Reducing Agent Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E0cell is positive, which indicates that the forward reaction is spontaneous (galvanic). If the system has a negative E0cell , the system must be nonspontaneous (electrolytic).
Standard electrode potentials are used to predict which redox couple (if there are many) will occur at the specified standard conditions
Will silver ions (Ag+) oxidize metallic zinc (to Zn2+) or will zinc ions (Zn2+) oxidize metallic Ag (to Ag+)? (Which is more spontaneous?)
Gibbs Free Energy and Electrical Potential at Standard Conditions For a redox reaction at standard states G0 = -nFE0cell96 485 J G 0 = - 2 mole e - + 1.5662 V V mole e G 0 = - 302230 J/mol rxn or - 302.23 kJ/mol
Will permanganate ions, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution? (Which is more spontaneous?)
n F E0cell G0
# of electrons involved in the balanced reaction 96 485 J/Vmol e- (Faradays Constant) cell potential at standard conditions standard Gibbs Free Energy
Relationship of E0cell and Equilibrium constant Keq G0 = -nFE0cell = -RTlnKeq 0 E cell = RTlnKeq / nF or lnKeq = nFE0cell / RT
96 485 J G 0 = - 4 mole e- +0.74 V V mole e G 0 = - 285595 J/mol rxn or - 285.60 kJ/mol
Consider the following reduction half-reactions involving water+ O 2(g) + 4H (aq) + 4e - 2H 2 O (l)
Electrolytic Cells Nonspontaneous electrochemical cells E0cell is negative G0 is positive [G0 = -nF(-E0cell value)= + value]Eletrolysis means forcing a nonspontaneous redox reaction to occur (at the electrodes) by applying potential+ O 2(g) + 4H(aq) + 4e- 2H2 O(l)
2(H 2(g) + 2OH
2H 2 O (l) + 2e )
0.83 V 2.06 V
O 2(g) + 2H 2 (g) 2H 2 O (i)
More spontaneous, but cathode requires H+ (will come from H2 gas)
O 2(g) + 4H + + 4e - 2H 2 O (l) (aq) O 2(g) + 2H 2 O (l) + 4 e - 4OH -(aq) 2H 2 O (l) + 2 e - H 2(g) + 2 OH -(aq)O 2(g) + 2H 2 O (l) + 4e - 4OH -(aq) 2(H 2(g) + 2OH -(aq) 2H 2 O (l) + 2e - ) O 2(g) + 2H 2 (g) 2H 2 O (i) 0.40 V 0.83 V 1.23 V
1.23 V 0.40 V - 0.83 VLess spontaneous, but cathode requires H2O
1.23 V 0.83 V 2.06 V
2(H2(g) + 2OH-(aq) 2H 2 O(l) + 2e - ) O 2(g) + 2H2 (g) 2H2 O(l)
Spontaneous, releases 212.3 kJ/mole of energy Nonspontaneous, requires 212.3 kJ/mole of energy
Electrolytic Cells Electrolytic cells requires external potential source In electrolytic cells, re