Electrochemistry Part 1 12/03/24 Electron Transfer Reactions: Electrochemistry Electrochemistry: The branch of chemistry that deals with electron transfer reactions • Electron transfer occurs in many chemical reactions but is not always obvious by looking at the chemical equation Example: πΆπ’ππ4 β 5π»2 π ππ + ππ π → ππππ4 β 5π»2 π + πΆπ’ π πΆπ’2+ (ππ) + 2π − → πΆπ’ (π ) Copper gains electrons ππ π → 2π − + ππ2+ (ππ) Zinc loses electrons Oxidation and Reduction From the previous example, Copper gained electrons so we would say that the Copper was reduced, while Zinc lost electrons so we would say that Zinc was oxidized. Oxidation: A process where a chemical species loses electrons Reduction: A process where a chemical species gains electrons LEO the Lion Says GER LEO: Lose Electrons Oxidation GER: Gain Electrons Reduction We often call electrochemistry reactions redox reactions for short! Electron Bookkeeping: Oxidation States Since chemical equations don’t always tell us that electron transfer is occurring, how do we know if a particular reaction is a redox reaction? Oxidation State Method: A scientific model used to identify electron transfer in a chemical reaction. • The method works by comparing the number of electrons a particular element has to how many valence electrons that element would have in a neutral state. • During this method, we must draw the structure of the chemical species in question, and assign the electrons as belonging to a particular atom • Lone pairs belong to the atom they are on, however; bonding electrons belong to the more electronegative atom Shortcut Method • Atoms in their natural state (H2, O2, He, Sn) have an oxidation state of 0 • Oxygen bonded to another atom has an oxidation state of -2 • Hydrogen bonded to another atom has an oxidation state of +1 • Ions have an oxidation state equal to their charge • The oxidation states of all the atoms in a compound must add up to the charge of that compound Electron Transfer Examples On Board For each reaction, determine if it is a redox reaction. If so, which species is being reduced and which species is being oxidized. • H2SO4 + 8HI ο H2S + 4I2 + 4H2O • Na2CO3 + 2HCl ο 2NaCl + H2O + CO2 • Pb(NO3)4 + 2Fe(NO3)2 ο 2Fe(NO3)3 + Pb(NO3)2 Balancing Redox Reactions • Often times redox reactions are given in an unbalanced form, and we need to make sure everything is balanced • The first step is to find the oxidation states to identify which species is reduced and which is oxidized Balancing Redox Reactions Example P1 πΉπ 2+ ππ + πππ4− ππ → πΉπ 3+ ππ + ππ2+ (ππ) Fe:2+ to 3+ losing 1 electron oxidized Mn:-1 to 2+ gaining 3 electrons reduced Next, we split the reaction into two half reactions: πΉπ 2+ aq → πΉπ 3+ aq πππ4− ππ → ππ2+ (aq) If elements need to be balanced, we can add coefficients to balance them. If there are any missing oxygens in the half reactions we add water molecules to balance. Balancing Redox Reactions Example P2 πΉπ 2+ aq → πΉπ 3+ aq πππ4− ππ → ππ2+ ππ + 4π»2 π If there are any missing Hydrogens, we add H+ to balance πππ4− ππ + 8π» + (ππ) → ππ2+ + 4π»2 π This should result in all of the atoms being balanced for both half reactions Next, we need to balance the half reactions by adding electrons πΉπ 2+ aq → πΉπ 3+ aq + 1π − πππ4− ππ + 8π» + ππ + 5 π − → ππ2+ + 4π»2 π The next step is to multiply by coefficients until both reactions have the same number of electrons 5πΉπ 2+ aq → 5πΉπ 3+ aq + 5e− πππ4− ππ + 5π − + 8π» + (ππ) → ππ2+ ππ + 4π»2 π Balancing Redox Reactions Example P3 Finally, we can add both reactions together cancelling any species on both sides of the overall equation, which should result in cancelling out the electrons. 5πΉπ 2+ aq → 5πΉπ 3+ aq + 5e− πππ4− ππ + 5π − + 8π» + (ππ) → ππ2+ ππ + 4π»2 π 5πΉπ 2+ aq + πππ4− ππ + 5π − + 8π» + (ππ) → 5πΉπ 3+ aq + 5e− + ππ2+ (ππ) + 4π»2 π 5πΉπ 2+ aq + πππ4− ππ + 8π» + (ππ) → 5πΉπ 3+ aq + ππ2+ (ππ) + 4π»2 π If everything was done correctly, you should end up with a fully balanced equation in terms of all atoms and charges with no electrons. Basic Conditions Under basic reaction conditions, we add one more step to the process because any protons will be neutralized by the OH- present in solution, so we add hydroxide to both sides to cancel out the H+ 5πΉπ 2+ aq + πππ4− ππ + 8π» + ππ + 8ππ» − ππ → 5πΉπ 3+ aq + ππ2+ ππ + 4π»2 π ππ + 8ππ» − (ππ) The hydroxide and protons on the same side of the reaction combine to form water: 5πΉπ 2+ aq + πππ4− ππ + 8π»2 π ππ → 5πΉπ 3+ aq + ππ2+ ππ + 4π»2 π ππ + 8ππ» − (ππ) Once the water forms, we can cancel it from one side: 5πΉπ 2+ aq + πππ4− ππ + 4π»2 π ππ → 5πΉπ 3+ aq + ππ2+ + 8ππ» − (ππ) Why Redox Reactions Happen • Some atoms and molecules hold on to electrons more tightly than others, and therefore have a stronger or weaker electric potential energy • This potential energy difference between species causes electrons to flow from one species to another in order to minimize the overall energy of the system When Redox Reactions Become Useful P1 πΆπ’ππ4 ππ + ππ π → ππππ4 ππ + πΆπ’ (π ) When Redox Reactions Become Useful P2 • What if instead of simply mixing the reactants together and allowing the electrons to flow in random directions, we only allow the electrons to flow in one direction – Create electrical work (current) • Using some clever engineering, we can force the electrons to only move one direction through a wire, and utilize the current as an electrical energy source (battery) When Redox Reactions Become Useful P3 πΆπ’ππ4 ππ + ππ π → ππππ4 ππ + πΆπ’ (π ) Galvanic Cells P1 Electron Flow • Use Redox Reactions to do Electrical Work • Consist of Five Main Pieces – Electrodes Anode/Cathode – Species to be Reduced – Species to be Oxidized – Current Collectors (wire) – Salt Bridge Current Collector Salt Bridge • Electrons Flow – Always Anode to Cathode Cathode Cu Metal Species to be Reduced Cu2+ Anode Mg Metal Species to be Oxidized Galvanic Cells P2 There are several new terms that are being used and must be defined Half-Cell: One of the two half reactions that is attached to the galvanic cell Electrode: A conductive surface where a redox reaction occurs Anode: The electrode where oxidation takes place Cathode: The electrode where reduction takes place An Ox and a Red Cat Salt Bridge: A device that balances the charges in a galvanic cell by allowing ions to flow into both reaction vessels Standard Electric Potential: The voltage difference between the specified half-reaction and a ° standard hydrogen electrode (πΈπΆπππ ) -Can be looked up in a half reaction table -Usually all values are written as reduction reactions ° -If you need to flip an equation around and make it oxidation, you multiply πΈπΆπππ by negative 1 Galvanic Cell Example • Draw a galvanic cell for the following reaction, make sure to label the anode, cathode, and direction of electron flow. πΉπ 3+ ππ + πΆπ π → πΉπ π + πΆπ 3+ (ππ) Galvanic Cells P3 Calculating total cell potential: The total cell potential is found by adding the reduction half-reaction potential and oxidation half-reaction potential ° ° ° πΈπΆπππ = πΈππππ’πππ + πΈππ₯ππππ§ππ • This differs from your book slightly, but my method is easier, because you don’t need to remember which one to subtract • A positive voltage means the reaction is spontaneous as written Cell Potential Example Calculate the Standard Cell Potential of a Galvanic Cell made using the following reaction: ππ π + πΆπ’2+ ππ → ππ2+ ππ + πΆπ’ (π ) You find the reduction half reactions in a table: ° ππ2+ ππ + 2π − → ππ π πΈπΆπππ = −0.76 π ° 2+ − πΆπ’ ππ + 2π → πΆπ’ π πΈπΆπππ = 0.34 π Note: We do not multiply by the coefficients when calculating cell potential, it is an intrinsic property Galvanic Cells P4 Cell Line Notation: ππ π ππ2+ ππ πΆπ’2+ ππ πΆπ’ π • The oxidation half-reaction is on the left of the double line, and the reduction half-reaction is on the right of the double line • Different Phases are separated by a single line • If there are multiple species in a particular phase, they are separated by a comma • The solid electrodes are placed at the far left (anode) and far right(cathode) • If a reaction does not have a solid phase, then a Platinum electrode is typically used Line Notation Example Write the line notation for the following balanced redox reaction, draw this galvanic cell. 5πΉπ π + 2πππ4− ππ + 16π» + ππ → 5πΉπ 2+ ππ + 2ππ2+ ππ + 8π»2 π π πΉπ π πΉπ 2+ πππ4− ππ , π» + ππ , ππ2+ ππ ππ‘(π )
