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37 Cards in this Set
- Front
- Back
Redox reaction |
oxidation-reduction reaction
electrons are transferred form one atom to another atom that loses electrons is oxidized atom that gains electrons is reduced |
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Oxidized
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atom loses electrons in redox reaction
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reduced
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atom gains electrons in redox reaction
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oxidation states
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possible charge values that an atom may hold within a molecule
necessary for redox reactions must add up to charge on molecule or ion |
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oxidation state = 0
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atoms in their elemental form
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oxidation state of fluorine (F)
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equals -1
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oxidation state of hydrogen (H)
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equals +1
except when bonded to a metal; then -1 |
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oxidation state of oxygen (O)
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equals -2
except when it is in a peroxide like H2O2 |
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Oxidation state = +1
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group 1 elements
alkali metals |
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oxidation state = +2
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group 2 elements
alkaline earth metals |
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oxidation state = -3
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group 15 elements
nitrogen family |
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oxidation state = -2
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group 16 elements
oxygen family |
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oxidation state = -1
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group 17 elements
halogens |
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LEO the lion says GER
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LEO: Lose Electrons Oxidation
GER: Gain Electrons Reduction |
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Reducing agent
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reductant
giving electrons to an atom losing electrons, is oxidized |
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Oxidizing agent
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compound containing atom being reduced
gains electrons, is reduced oxidizes other atom |
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Electric potential (E)
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associated with any redox reaction
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Half-reaction
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Each component of a redox reaction
oxidation half-reaction potential is opposite reduction half-reaction potential usually listed as reduction potentials (sign is reversed for oxidation potential) |
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Half-reaction of standard hydrogen electrode
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2H+ + 2e- --> H2
Half-reaction potential = 0.00V |
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Balancing redox reactions in acidic solutions
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1. divide reaction into its corresponding half reactions
2. balance the elements other than H and O 3. Add H20 to one side until O atoms are balanced 4. Add H+ to one side until H atoms are balanced 5. Add e- to one side until charge is balanced 6. multiply each half reaction by an integer so that an equal number of electrons are transferred in each reaction 7. add the 2 half reactions and simplify |
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Balancing redox reactions in basic solutions
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same steps as acidic solutions
neutralize H+ ions by adding same number of OH- ions to both sides of reaction |
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Galvanic cell (voltaic cell)
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uses electric potential between phases to generate a current of electrons from one phase to another in a conversion of chemical energy to electrical energy
turns chemical energy into electrical energy |
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Salt bridge
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ionic conducting phase
electrolyte solution phase impermeable to electrons type of liquid junction that minimizes potential difference between different solutions carries current in form of ions |
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Terminals
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electronic conductors such as metal wires (T)
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Electrodes
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electronic conductors (E)
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Ionic conductor
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salt bridge (I)
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Simple galvanic cell
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T-E-I-E'-T'
has 2 electrodes: anode (-) and cathode (+) oxidation half reaction takes place at anode reduction half reaction takes place at cathode 2 terminals of cell is made from same material |
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cell potential (E)
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electromotive force (emf)
potential difference between the terminals when they are not connected connecting the terminals reduces the potential difference due to internal resistance within the cell drop in emf increases, as current increases current flows in direction opposite electron flow electrons flow from anode to cathode always positive, always has chemical energy than can be converted to work |
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RED CAT
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REDuction CAThode
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AN OX
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ANode OXidation
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Positive cell potential
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equals a negative Gibbs free energy (deltaG), which equals a spontaneous reaction (work is being done by system and not on system)
deltaG = -nFEmax deltaG: Gibbs free energy n: number of moles of electrons that are transferred in balanced redox reaction F: Faraday's constant E: voltage Free energy (deltaG) represents the product of total charge (nF) times voltage (E) |
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Reactions that do not occur at standard state
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deltaG = deltaG(not) + RT[ln(Q)]
deltaG: Gibbs free energy deltaG(not): Gibbs free energy (standard conditions) T: temperature Q: reaction quotient |
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reactions that are at equilibrium conditions
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at equilibrium, there is no available free energy with which to do work; deltaG = 0
deltaG(not) = -RT[ln(K)] |
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relationship between K and deltaG(not)
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if K = 1, then deltaG(not) = 0
if K > 1, then deltaG(not) < 0 if K < 1, then deltaG(not) > 0 |
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Concentration cell
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limited form of a galvanic cell with a reduction half reaction taking place in 1 half cell and the exact reverse of that half reaction taking place in the other half cell
type of galvanic cell it is never at standard conditions, so Nerst equation is required to solve for cell potential if concentrations were equal on both sides, the concentration cell potential would be zero |
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galvanic cell
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positive cell potential
spontaneous |
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electrolytic cell
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negative cell potential
forced by outside power source to run backwards cathode is negative anode is positive RED CAT & AN OX still the same |