Assigning Oxidation States Rules: 1. Element will always have an oxidation state of 0 2. Monatomic ions have oxidation states equal to their charge 3. Sum of oxidation states has to equal overall charge 4. Metals. . . .
⁃ Group I metal ion. . . +1
⁃ Group II metal ion. . . +2
5. Nonmetals. . .
⁃ Fluoride: -1
⁃ Hydrogen: +1
⁃ Oxygen: -2
⁃ Group VII: -1
⁃ Group VI: -2
⁃ Group V: -3
Balance Redox Reactions :
1. Identify the atoms involved in redox
2. Set up half-reactions
3. Balance the atoms involved
4. Determine the electrons in each half reaction 5. Balance electrons between half reactions 6. Add up half reactionsWe also discuss several other topics like What are the positive points of obe?
Balance Redox Reactions (cont.): 7. Balance any oxygen with water (H2O)
8. Balance any hydrogen with hydrogen ions (H+) 9. Balance any other spectators
10. Add as many OH- to both sides as H+ 11. Combine OH- and H+ to produce water 12. Add up or cancel H2O if needed
• We write the oxidation half reaction on the left and the reduction in the right. A double vertical line, indicating the salt bridge, separates the two half reactions
• Substances in different phases are separated by a single vertical line which represents the boundary between the phases Don't forget about the age old question of How many joules are needed to change the temperature of 1 gram of water by 1?
• For some Redox Reactions the reactants and products of one or both of the half reactions may be in the same phase. In these cases we separate the reactants and products from each other with a comma in the line diagram. Such cells use an inert electrode such as platinum (Pt) or graphite as the anode or cathode or both
Don't forget about the age old question of Divide the name into the name of the branches and the name of the stem.
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Energy Spontaneous Voltage
Summarizing Standard Electrode Potentials:
• The electrode potential of the standard hydrogen electrode (SHE) is exactly 0 • The electrode in any half-cell with a greater tendency to undergo reduction is positively charged relative to the SHE and therefore has a positive E standard
• The electrode in any half-cell with a lesser tendency to undergo reduction (or greater tendency to undergo oxidation) is negatively charged relative to the SHE and therefore has a negative E standard
• The cell potential of any electrochemical cell is the difference between the electrode potentials of the cathode and the anode
• The standard E cell is positive for spontaneous reactions and negative for nonspontaneous reactions
• The half reaction with the more positive electrode potential attracts electrons more strongly and will undergo reduction
• The half reactions with the more negative electrode potential repels electrons ,ore strongly and will undergo oxidation
• Any reduction reaction in Table 18.1 is spontaneous when paired with the reverse of any of the reactions listed below it on the table If you want to learn more check out What are the characteristics of the phylum porifera?
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Summarizing Characteristics of Electrochemical Cell Types: • In all electrochemical cells:
⁃ Oxidation occurs at the anode
⁃ Reduction occurs at the cathode
• In voltaic cells:
⁃ The anode is the source of electrons and has a negative charge (anode -) ⁃ The cathode draws electrons and has a positive charge (cathode +)
• In electrolytic cells:
⁃ Electrons are drawn away from the anode which must be connected to the positive terminal of the external power source ( anode +)
⁃ Electrons are forced to the cathode which must be connected to the negative terminal of the power source (cathode -)