Electrochemistry Redox PDF

ELECTROCHEMISTRY Electron Transfer Fun Electron Transfer ▪ The movement of electrons is the driving force behind redox reactions ▪ When an atom or molecule loses electrons it is oxidized ▪ It becomes more positive in charge ▪ Reducing agent ▪ Atoms or molecules gaining electrons become...

ELECTROCHEMISTRY Electron Transfer Fun Electron Transfer ▪ The movement of electrons is the driving force behind redox reactions ▪ When an atom or molecule loses electrons it is oxidized ▪ It becomes more positive in charge ▪ Reducing agent ▪ Atoms or molecules gaining electrons become reduced ▪ Oxidizing agent ▪ Redox reactions must happen in pairs ▪ In order for one to be oxidized another MUST be reduced  Valence vs Oxidation State Oxidation is the number of Valency is the number of electrons that a particular atom electrons present in the can lose, gain or share with outermost shell another atom Does NOT indicate electrical Indicates the electrical charge of charge an atom Determines the maximum Does not indicate the number of number of bonds that an atom bonds that a particular atom can can have have Valency of a pure element Oxidation state of a pure element depends on the number of is always zero electrons in the outer most shell Importance of Redox Reaction Almost all batteries (including the one in your phone) Combustion of carbon based molecules in oxygen – Including cellular respiration and photosynthesis – Hydrocarbon burning for fuel Single displacement reactions RULES FOR ASSIGNING OXIDATION NUMBERS Assigning Oxidation Numbers N in N2O3 S in H2SO4 Half Reaction Equations Used to balance redox reactions Represents one of the two parts of the redox reaction Balancing Using ½ Reaction Write unbalanced equation and assign oxidation numbers to all entities Write the ½ reactions Balance each half reaction Use electrons to balance the charge in each half reaction Use water to balance oxygens; then use H+ and OH- to balance hydrogen Equalize the electron transfer in the two half reactions Add the half reactions (Hess’s Law) Ex. potassium permanganate (KMnO4) and sodium sulfite (Na 2SO3) Unbalanced net ionic: – MnO4 − + SO3−2(aq) → MnO4−2(aq) + SO4−2(aq) Manganese changes from +7 to +6 (reduction) and sulfur changes from +4 to +6 (oxidation). This makes MnO4− the oxidizing agent and SO32- the reducing agent: Now balance each half-reaction by starting with all atoms except H and O. Then balance for O atoms and add H2O to the side of the reaction missing O. Then balance for H by adding H+ to the side missing H Oxidation:MnO4− (aq) → MnO42−(aq) Reduction:SO32−(aq) + H2O(l) → SO42−(aq) + 2H+(aq) Now we add electrons to the most positively charged side of the reaction to balance the charges on each side. The coefficients and the charge play a role. – In the oxidation step, an electrons is added on the left to match the -2 charge on the right. – In the reduction step, electrons are added on the right to match the -2 charge on the left. Oxidation: MnO4− (aq) + 1e- → MnO42−(aq) Reduction: SO32−(aq) + H2O(l) → SO42−(aq) + 2H+(aq) + 2e- Next multiply each half-reaction so that the number of electrons in each are equal. Oxidation x2 : 2MnO4− (aq) + 2e- → 2MnO42−(aq) Now we can add the two half reactions together. Any species that occur on both sides of the reaction are canceled If a reaction occurs in a basic or acidic solution. The last step is to add H+ or OH− to the reactants to account for those conditions. More on this later… Since most Redox Reaction take place in aqueous solution H2O molecules can be used to balance in acidic or basic situations In acidic solutions H+ is a participant in the reactants; in basic solutions OH- In order to balance H+; OH- and H2O is added to balance the equation