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1 Chapter 20 Electrochemistry. 2 Chapter Goals 1. Electrical Conduction 2. Electrodes Electrolytic Cells 3. The Electrolysis of Molten Potassium Chloride.

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Presentation on theme: "1 Chapter 20 Electrochemistry. 2 Chapter Goals 1. Electrical Conduction 2. Electrodes Electrolytic Cells 3. The Electrolysis of Molten Potassium Chloride."— Presentation transcript:

1 1 Chapter 20 Electrochemistry

2 2 Chapter Goals 1. Electrical Conduction 2. Electrodes Electrolytic Cells 3. The Electrolysis of Molten Potassium Chloride 4. The Electrolysis of Aqueous Potassium Chloride 5. The Electrolysis of Aqueous Potassium Sulfate 6. Faraday’s Law of Electrolysis 7. Commercial Applications of Electrolytic Cells

3 3 Chapter Goals Voltaic or Galvanic Cells 8. The Construction of Simple Voltaic Cells 9. The Zinc-Copper Cell 10. The Copper-Silver Cell Standard Electrode Potentials 11. The Standard Hydrogen Electrode 12. The Zinc-SHE Cell 13. The Copper-SHE Cell 14. Standard Electrode Potentials 15. Uses of Standard Electrode Potentials

4 4 Chapter Goals 16. Standard Electrode Potentials for Other Half- Reactions 17. Corrosion 18. Corrosion Protection Effect of Concentrations (or Partial Pressures) on Electrode Potentials 19. The Nernst Equation 20. Using Electrohemical Cells to Determine Concentrations 21. The Relationship of E o cell to  G o and K

5 5 Chapter Goals Primary Voltaic Cells 22. Dry Cells Secondary Voltaic Cells 23. The Lead Storage Battery 24. The Nickel-Cadmium (Nicad) Cell 25. The Hydrogen-Oxygen Fuel Cell

6 6 Electrochemistry Electrochemical reactions are oxidation-reduction reactions. The two parts of the reaction are physically separated.  The oxidation reaction occurs in one cell.  The reduction reaction occurs in the other cell. There 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.

7 7 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.  Negatively charged ions, anions, move toward the positive electrode.

8 8 Electrodes The following convention for electrodes is correct for either electrolytic or voltaic cells: The cathode is the electrode at which reduction occurs. The cathode is negative in electrolytic cells and positive in voltaic cells. The anode is the electrode at which oxidation occurs. The anode is positive in electrolytic cells and negative in voltaic cells.

9 9 Electrodes Inert electrodes do not react with the liquids or products of the electrochemical reaction. Two examples of common inert electrodes are graphite and platinum.

10 10 Electrolytic Cells Electrical energy is used to force nonspontaneous chemical reactions to occur. The process is called electrolysis. Two examples of commercial electrolytic reactions are: 1. The electroplating of jewelry and auto parts. 2. The electrolysis of chemical compounds.

11 11 Electrolytic Cells Electrolytic cells consist of: 1. A container for the reaction mixture. 2. Two electrodes immersed in the reaction mixture. 3. A source of direct current.

12 12 The Electrolysis of Molten Potassium Chloride Liquid potassium is produced at one electrode.  Indicates that the reaction K + ( ) + e -  K ( ) occurs at this electrode.  Is this electrode the anode or cathode? Gaseous chlorine is produced at the other electrode.  Indicates that the reaction 2 Cl -  Cl 2(g) + 2 e - occurs at this electrode.  Is this electrode the anode or cathode?

13 13 The Electrolysis of Molten Potassium Chloride Diagram of this electrolytic cell. Porous barrier e-e- K + + e -  K ( ) cathode reaction 2Cl -  Cl 2 (g) + 2e - anode reaction Generator-source of DC - electrode + electrode e-e- molten KCl

14 14 The Electrolysis of Molten Potassium Chloride The nonspontaneous redox reaction that occurs is:

15 15 The Electrolysis of Molten Potassium Chloride In all electrolytic cells, electrons are forced to flow from the positive electrode (anode) to the negative electrode (cathode).

16 16 The Electrolysis of Aqueous Potassium Chloride In this electrolytic cell, hydrogen gas is produced at one electrode.  The aqueous solution becomes basic near this electrode. You do it!  What reaction is occurring at this electrode? You do it! Gaseous chlorine is produced at the other electrode. You do it!  What reaction is occurring at this electrode? You do it! These experimental facts lead us to the following nonspontaneous electrode reactions:

17 17 The Electrolysis of Aqueous Potassium Chloride 2 H 2 O + 2e -  H 2 (g) + 2 OH - cathode reaction 2Cl -  Cl 2 (g) + 2e - anode reaction Cell diagram Battery, a source of direct current e - flow - electrode+ electrode e - flow aqueous KCl Cl 2 gasH 2 gas + pole of battery- pole of battery

18 18 The Electrolysis of Aqueous Potassium Sulfate In this electrolysis, hydrogen gas is produced at one electrode.  The solution becomes basic near this electrode.  What reaction is occurring at this electrode? You do it! Gaseous oxygen is produced at the other electrode  The solution becomes acidic near this electrode.  What reaction is occurring at this electrode? You do it! These experimental facts lead us to the following electrode reactions:

19 19 The Electrolysis of Aqueous Potassium Sulfate

20 20 The Electrolysis of Aqueous Potassium Sulfate 2 H 2 O + 2e -  H 2 (g) + 2 OH - cathode reaction 2H 2 O  O 2 (g) + 4H + + 4e - anode reaction Cell diagram Battery, a source of direct current e - flow - electrode+ electrode e - flow aqueous K 2 SO 4 O 2 gas H 2 gas + pole of battery- pole of battery

21 21 Electrolytic Cells In all electrolytic cells the most easily reduced species is reduced and the most easily oxidized species is oxidized.

22 22 Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Faraday’s Law - The amount of substance undergoing chemical reaction at each electrode during electrolysis is directly proportional to the amount of electricity that passes through the electrolytic cell. A faraday is the amount of electricity that reduces one equivalent of a species at the cathode and oxidizes one equivalent of a species at the anode.

23 23 Counting Electrons: Coulometry and Faraday’s Law of Electrolysis A coulomb is the amount of charge that passes a given point when a current of one ampere (A) flows for one second. 1 amp = 1 coulomb/second

24 24 Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Faraday’s Law states that during electrolysis, one faraday of electricity (96,487 coulombs) reduces and oxidizes, respectively, one equivalent of the oxidizing agent and the reducing agent.  This corresponds to the passage of one mole of electrons through the electrolytic cell.

25 25 Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Example 21-1: Calculate the mass of palladium produced by the reduction of palladium (II) ions during the passage of 3.20 amperes of current through a solution of palladium (II) sulfate for 30.0 minutes.

26 26 Counting Electrons: Coulometry and Faraday’s Law of Electrolysis Example 21-2: Calculate the volume of oxygen (measured at STP) produced by the oxidation of water in example 21-1.

27 27 Commercial Applications of Electrolytic Cells Electrolytic Refining and Electroplating of Metals Impure metallic copper can be purified electrolytically to  100% pure Cu.  The impurities commonly include some active metals plus less active metals such as: Ag, Au, and Pt. The cathode is a thin sheet of copper metal connected to the negative terminal of a direct current source. The anode is large impure bars of copper.

28 28 Commercial Applications of Electrolytic Cells The electrolytic solution is CuSO 4 and H 2 SO 4 The impure Cu dissolves to form Cu 2+. The Cu 2+ ions are reduced to Cu at the cathode.

29 29 Commercial Applications of Electrolytic Cells Any active metal impurities are oxidized to cations that are more difficult to reduce than Cu 2+.  This effectively removes them from the Cu metal.

30 30 Commercial Applications of Electrolytic Cells The less active metals are not oxidized and precipitate to the bottom of the cell. These metal impurities can be isolated and separated after the cell is disconnected. Some common metals that precipitate include:

31 31 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:

32 32 The Construction of Simple Voltaic Cells Voltaic cells consist of two half-cells which contain the oxidized and reduced forms of an element (or other chemical species) 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.

33 33 The Construction of Simple Voltaic Cells

34 34 The Zinc-Copper Cell Cell components for the Zn-Cu cell are: 1. A metallic Cu strip immersed in 1.0 M copper (II) sulfate. 2. A metallic Zn strip immersed in 1.0 M zinc (II) sulfate. 3. A wire and a salt bridge to complete circuit The cell’s initial voltage is 1.10 volts

35 35 The Zinc-Copper Cell In all voltaic cells, electrons flow spontaneously from the negative electrode (anode) to the positive electrode (cathode).

36 36 The Zinc-Copper Cell There is a commonly used short hand notation for voltaic cells.  The Zn-Cu cell provides a good example.

37 37 The Copper - Silver Cell Cell components: 1. A Cu strip immersed in 1.0 M copper (II) sulfate. 2. A Ag strip immersed in 1.0 M silver (I) nitrate. 3. A wire and a salt bridge to complete the circuit. The initial cell voltage is 0.46 volts.

38 38 The Copper - Silver Cell 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. Whether a particular electrode behaves as an anode or as a cathode depends on what the other electrode of the cell is.

39 39 The Copper - Silver Cell These experimental facts demonstrate that Cu 2+ is a stronger oxidizing agent than Zn 2+.  In other words Cu 2+ oxidizes metallic Zn to Zn 2+. Similarly, Ag + is is a stronger oxidizing agent than Cu 2+.  Because Ag + oxidizes metallic Cu to Cu 2+. If we arrange these species in order of increasing strengths, we see that:

40 40 Standard Electrode Potential 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 … V

41 41 The Zinc-SHE Cell For this cell the components are: 1. A Zn strip immersed in 1.0 M zinc (II) sulfate. 2. The other electrode is the Standard Hydrogen Electrode. 3. A wire and a salt bridge to complete the circuit. The initial cell voltage is volts.

42 42 The Zinc-SHE Cell The cathode is the Standard Hydrogen Electrode.  In other words Zn reduces H + to H 2. The anode is Zn metal.  Zn metal is oxidized to Zn 2+ ions.

43 43 The Copper-SHE Cell The cell components are: 1. A Cu strip immersed in 1.0 M copper (II) sulfate. 2. The other electrode is a Standard Hydrogen Electrode. 3. A wire and a salt bridge to complete the circuit. The initial cell voltage is volts.

44 44 The Copper-SHE Cell In this cell the SHE is the anode  The Cu 2+ ions oxidize H 2 to H +. The Cu is the cathode.  The Cu 2+ ions are reduced to Cu metal.

45 45 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. For example, the half-reaction for the standard potassium electrode is: The large negative value tells us that this reaction will occur only under extreme conditions.

46 46 Uses of Standard Electrode Potentials Compare the potassium half-reaction to fluorine’s half- reaction: The large positive value denotes that this reaction occurs readily as written. Positive E 0 values denote that the reaction tends to occur to the right.  The larger the value, the greater the tendency to occur to the right. It is the opposite for negative values of E o.

47 47 Uses of Standard Electrode Potentials Use standard electrode potentials to predict whether an electrochemical reaction at standard state conditions will occur spontaneously. Example 21-3: Will silver ions, Ag +, oxidize metallic zinc to Zn 2+ ions, or will Zn 2+ ions oxidize metallic Ag to Ag + ions? Steps for obtaining the equation for the spontaneous reaction.

48 48 Uses of Standard Electrode Potentials 1. Choose the appropriate half-reactions from a table of standard reduction potentials. 2. Write the equation for the half-reaction with the more positive E 0 value first, along with its E 0 value. 3. Write the equation for the other half-reaction as an oxidation with its oxidation potential, i.e. reverse the tabulated reduction half-reaction and change the sign of the tabulated E 0. 4 Balance the electron transfer. 5 Add the reduction and oxidation half-reactions and their potentials. This produces the equation for the reaction for which E 0 cell is positive, which indicates that the forward reaction is spontaneous.

49 49 Uses of Standard Electrode Potentials

50 50 Electrode Potentials for Other Half-Reactions Example 21-4: Will permanganate ions, MnO 4 -, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution? Follow the steps outlined in the previous slides. Note that E 0 values are not multiplied by any stoichiometric relationships in this procedure.

51 51 Electrode Potentials for Other Half-Reactions Example 21-4: Will permanganate ions, MnO 4 -, oxidize iron (II) ions to iron (III) ions, or will iron (III) ions oxidize manganese(II) ions to permanganate ions in acidic solution? Thus permanganate ions will oxidize iron (II) ions to iron (III) and are reduced to manganese (II) ions in acidic solution.

52 52 Electrode Potentials for Other Half-Reactions Example 21-5: Will nitric acid, HNO 3, oxidize arsenous acid, H 3 AsO 3, in acidic solution? The reduction product of HNO 3 is NO in this reaction. You do it!

53 53 Corrosion Metallic corrosion is the oxidation-reduction reactions of a metal with atmospheric components such as CO 2, O 2, and H 2 O.

54 54 Corrosion Protection Some examples of corrosion protection. 1 Plate a metal with a thin layer of a less active (less easily oxidized) metal.

55 55 Corrosion Protection 2. Connect the metal to a sacrificial anode, a piece of a more active metal.

56 56 Corrosion Protection 3. Allow a protective film to form naturally.

57 57 Corrosion Protection 4 Galvanizing, the coating of steel with zinc, provides a more active metal on the exterior. Zinc Steel

58 58 Corrosion Protection 5. Paint or coat with a polymeric material such as plastic or ceramic.

59 59 Effect of Concentrations (or Partial Pressures) on Electrode Potentials The Nernst Equation Standard electrode potentials, those compiled in appendices, are determined at thermodynamic standard conditions. Reminder of standard conditions M solution concentrations 1.00 atm of pressure for gases All liquids and solids in their standard thermodynamic states. Temperature of 25 0 C.

60 60 The Nernst Equation The value of the cell potentials change if conditions are nonstandard. The Nernst equation describes the electrode potentials at nonstandard conditions. The Nernst equation is:

61 61 The Nernst Equation

62 62 The Nernst Equation Substitution of the values of the constants into the Nernst equation at 25 o C gives:

63 63 The Nernst Equation For this half-reaction: The corresponding Nernst equation is:

64 64 The Nernst Equation Substituting E 0 into the above expression gives: If [Cu 2+ ] and [Cu + ] are both 1.0 M, i.e. at standard conditions, then E = E 0 because the concentration term equals zero.

65 65 The Nernst Equation

66 66 The Nernst Equation Example 21-6: Calculate the potential for the Cu 2+ / Cu + electrode at 25 0 C when the concentration of Cu + ions is three times that of Cu 2+ ions.

67 67 The Nernst Equation Example 21-6: Calculate the potential for the Cu 2+ / Cu + electrode at 25 0 C when the concentration of Cu + ions is three times that of Cu 2+ ions.

68 68 The Nernst Equation Example 21-6: Calculate the potential for the Cu 2+ / Cu + electrode at 25 0 C when the concentration of Cu + ions is three times that of Cu 2+ ions.

69 69 The Nernst Equation Example 21-7: Calculate the potential for the Cu 2+ /Cu + electrode at 25 0 C when the Cu + ion concentration is 1/3 of the Cu 2+ ion concentration. You do it!

70 70 The Nernst Equation Example 21-7: Calculate the potential for the Cu 2+ /Cu + electrode at 25 0 C when the concentration of Cu + ions is 1/3 that of Cu 2+ ions.

71 71 The Nernst Equation

72 72 The Nernst Equation Example 21-8: Calculate the electrode potential for a hydrogen electrode in which the [H + ] is 1.0 x M and the H 2 pressure is 0.50 atmosphere.

73 73 The Nernst Equation Example 21-8: Calculate the electrode potential for a hydrogen electrode in which the [H + ] is 1.0 x M and the H 2 pressure is 0.50 atmosphere.

74 74 The Nernst Equation The Nernst equation can also be used to calculate the potential for a cell that consists of two nonstandard electrodes. Example 21-9: Calculate the initial potential of a cell that consists of an Fe 3+ /Fe 2+ electrode in which [Fe 3+ ]=1.0 x M and [Fe 2+ ]=0.1 M connected to a Sn 4+ /Sn 2+ electrode in which [Sn 4+ ]=1.0 M and [Sn 2+ ]=0.10 M. A wire and salt bridge complete the circuit.

75 75 The Nernst Equation Calculate the E 0 cell by the usual procedure.

76 76 The Nernst Equation Substitute the ion concentrations into Q to calculate E cell.

77 77 The Nernst Equation

78 78 Relationship of E 0 cell to  G 0 and K From previous chapters we know the relationship of  G 0 and K for a reaction.

79 79 Relationship of E 0 cell to  G 0 and K The relationship between  G 0 and E 0 cell is also a simple one.

80 80 Relationship of E 0 cell to  G 0 and K Combine these two relationships into a single relationship to relate E 0 cell to K.

81 81 Relationship of E 0 cell to  G 0 and K Example 21-10: Calculate the standard Gibbs free energy change,  G 0, at 25 0 C for the following reaction.

82 82 Relationship of E 0 cell to  G 0 and K 1. Calculate E 0 cell using the appropriate half- reactions.

83 83 Relationship of E 0 cell to  G 0 and K 2. Now that we know E 0 cell, we can calculate  G 0.

84 84 Relationship of E 0 cell to  G 0 and K Example 21-11: Calculate the thermodynamic equilibrium constant for the reaction in example at 25 0 C.

85 85 Relationship of E 0 cell to  G 0 and K Example 21-12: Calculate the Gibbs Free Energy change,  G and the equilibrium constant at 25 0 C for the following reaction with the indicated concentrations.

86 86 Relationship of E 0 cell to  G 0 and K 1. Calculate the standard cell potential E 0 cell.

87 87 Relationship of E 0 cell to  G 0 and K 2. Use the Nernst equation to calculate E cell for the given concentrations.

88 88 Relationship of E 0 cell to  G 0 and K

89 89 Relationship of E 0 cell to  G 0 and K

90 90 Relationship of E 0 cell to  G 0 and K E cell = V, compared to E 0 cell = V. We can use this information to calculate  G. The negative  G tells us that the reaction is spontaneous.

91 91 Relationship of E 0 cell to  G 0 and K Equilibrium constants do not change with reactant concentration. We can use the value of E 0 cell at 25 0 C to get K.

92 92 Primary Voltaic Cells As a voltaic cell discharges, its chemicals are consumed. Once the chemicals are consumed, further chemical action is impossible. The electrodes and electrolytes cannot be regenerated by reversing current flow through cell.  These cells are not rechargable.

93 93 One example of a dry cell is flashlight, and radio, batteries. The cell’s container is made of zinc which acts as an electrode. A graphite rod is in the center of the cell which acts as the other electrode. The space between the electrodes is filled with a mixture of: 1. ammonium chloride, NH 4 Cl 2. manganese (IV) oxide, MnO 2 3. zinc chloride, ZnCl 2 4. and a porous inactive solid. The Dry Cell

94 94 The Dry Cell As electric current is produced, Zn dissolves and goes into solution as Zn 2+ ions. The Zn electrode is negative and acts as the anode.

95 95 The Dry Cell The anode reaction is: The graphite rod is the positive electrode (cathode). Ammonium ions from the NH 4 Cl are reduced at the cathode.

96 96 The Dry Cell The cell reaction is:

97 97 The Dry Cell The other components in the cell are included to remove the byproducts of the reaction. MnO 2 prevents H 2 from collecting on graphite rod. At the anode, NH 3 combines with Zn 2+ to form a soluble complex and removing the Zn 2+ ions from the reaction.

98 98 The Dry Cell

99 99 The Dry Cell Alkaline dry cells are similar to ordinary dry cells except that KOH, an alkaline substance, is added to the mixture. Half reactions for an alkaline cell are:

100 100 The Dry Cell Alkaline dry cells are similar to ordinary dry cells except that KOH, an alkaline substance, is added to the mixture. Half reactions for an alkaline cell are:

101 101 Secondary Voltaic Cells Secondary cells are reversible, rechargeable. The electrodes in a secondary cell can be regenerated by the addition of electricity.  These cells can be switched from voltaic to electrolytic cells. One example of a secondary voltaic cell is the lead storage or car battery.

102 102 The Lead Storage Battery In the lead storage battery the electrodes are two sets of lead alloy grids (plates). Holes in one of the grids are filled with lead (IV) oxide, PbO 2. The other holes are filled with spongy lead. The electrolyte is dilute sulfuric acid.

103 103 The Lead Storage Battery Diagram of the lead storage battery.

104 104 The Lead Storage Battery As the battery discharges, spongy lead is oxidized to lead ions and the plate becomes negatively charged. The Pb 2+ ions that are formed combine with SO 4 2- from sulfuric acid to form solid lead sulfate on the Pb electrode.

105 105 The Lead Storage Battery The net reaction at the anode during discharge is: Electrons are produced at the Pb electrode. These electrons flow through an external circuit (the wire and starter) to the PbO 2 electrode. PbO 2 is reduced to Pb 2+ ions, in the acidic solution. The Pb 2+ ions combine with SO 4 2- to form PbSO 4 and coat the PbO 2 electrode. PbO 2 electrode is the positive electrode (cathode).

106 106 The Lead Storage Battery The cell reaction for a discharging lead storage battery is: As the cell discharges, the cathode reaction is:

107 107 The Lead Storage Battery The cell reaction for a discharging lead storage battery is: As the cell discharges, the cathode reaction is:

108 108 The Lead Storage Battery What happens at each electrode during recharging? At the lead (IV) oxide, PbO 2, electrode, lead ions are oxidized to lead (IV) oxide. The concentration of the H 2 SO 4 decreases as the cell discharges. Recharging the cell regenerates the H 2 SO 4.

109 109 The Lead Storage Battery What happens at each electrode during recharging? At the lead (IV) oxide, PbO 2, electrode, lead ions are oxidized to lead (IV) oxide. The concentration of the H 2 SO 4 decreases as the cell discharges. Recharging the cell regenerates the H 2 SO 4.

110 110 The Nickel-Cadmium (Nicad) Cell Nicad batteries are the rechargeable cells used in calculators, cameras, watches, etc. As the battery discharges, the half-reactions are:

111 111 The Hydrogen-Oxygen Fuel Cell Fuel cells are batteries that must have their reactants continuously supplied in the presence of appropriate catalysts. A hydrogen-oxygen fuel cell is used in the space shuttle  The fuel cell is what exploded in Apollo 13. Hydrogen is oxidized at the anode. Oxygen is reduced at the cathode.

112 112 The Hydrogen-Oxygen Fuel Cell Notice that the overall reaction is the combination of hydrogen and oxygen to form water.  The cell provides a drinking water supply for the astronauts as well as the electricity for the lights, computers, etc. on board. Fuel cells are very efficient.  Energy conversion rates of 60-70% are common!

113 113 Synthesis Question What are the explosive chemicals in the fuel cell that exploded aboard Apollo 13?

114 114 Synthesis Question The Apollo 13 fuel cells contained hydrogen and oxygen. Both are explosive, especially when mixed. The oxygen tank aboard Apollo 13 exploded.

115 115 Group Question Some of the deadliest snakes in the world, for example the cobra, have venoms that are neurotoxins. Neurotoxins have an electrochemical basis. How do neurotoxins disrupt normal chemistry and eventually kill their prey?

116 116 End of Chapter 21 Electrochemistry is an important part of the electronics industry.


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