Ib Hl Chemistry Assessment Statements Topics 9 And 19

Andrew Voyles IB Chemistry Assessment Statements: Topic 9: Redox Topic 19: Redox 9.1.1: OILRIG – Oxidation Is Loss of electrons, Reduction Is Gain of electrons. 9.1.2: The oxidation number of an element is zero. The oxidation number of an ion is equal to the charge of the ion. In compounds, Hydrogen has an oxidation number of +1. In compounds, Oxygen (usually, but not in peroxides) has an oxidation number of -2. Typically, in a molecule, an atoms’ oxidation number is what its charge would be if the molecule were a polyatomic ion. 9.1.3: Iron oxide, comes in two forms. FeO and Fe2O3, Iron (II) oxide and Iron (III) oxide, respectively. 9.1.4: If an element is oxidized, its oxidation number will go up. If an element is reduced, its oxidation number will go down. 9.2.1: Oxidation half equation... Mg Mg2+ + 2e-. Reduction half equation... O + 2e O2-. 9.2.2: MnO4- + I- → I2 + Mn2+ 2I- → I2 MnO4- → Mn2+ 2I- → I2 MnO4- + 8 H+ → Mn2+ + 4 H2O 2 I- → I2 + 2e5 e- + 8 H+ + MnO4- → Mn2+ + 4 H2O 5(2I- → I2 +2e-) 2(5e- + 8H+ + MnO4- → Mn2+ + 4H2O) 10 I- + 10 e- + 16 H+ + 2 MnO4- → 5 I2 + 2 Mn2+ + 10 e- + 8 H2O 10 I- + 16 H+ + 2 MnO4- → 5 I2 + 2 Mn2+ + 8 H2O 9.2.3: An oxidizing agent is an element which causes oxidation (and is reduced as a result) by removing electrons from another species.

A reducing agent is an element which causes reduction (and is oxidized as a result) by giving electrons to another species. 9.2.4: Mg

Mg2+ + 2e-. Reducing agent O + 2eO2- Oxidizing agent

9.3.1: A more reactive metal will displace a less reactive one from a compound and a more reactive halogen will displace a weaker one from a compound. This can be generalized to say a stronger reducing agent will displace a weaker one from a compound, and a stronger oxidizing agent will displace a weaker one from a compound. Thus, if a metal displaces another, we know it must be more reactive and the same for halogens (which are the given examples). 9.4.1: Add the two half equations in question together (the oxidation half-reaction will have to be reversed, invert the sign of the standard reduction potential value also) If the total standard reduction potential value is positive, the reaction is possible. If not, it isn’t. 9.4.1: In the redox reaction that powers the voltaic cell, reduction (addition of electrons) occurs to cations at the cathode, while oxidation (removal of electrons) occurs to anions at the anode. The electrodes do not touch each other but are electrically connected by the electrolyte, which can be either solid or liquid.

9.4.2: Oxidation occurs at the negative electrode (the anode), and reduction occurs at the positive electrode (the cathode), in a voltaic cell.

9.5.1: This is, fundamentally, the reverse of an electrochemical cell. In this case a greater electromotive force is applied from the external circuit using a battery or power source and this forces the species within the cell to perform the reverse reaction to that which they would normally tend to do. It is important not to confuse electrochemical cells which generate electricity by means of a redox reaction and electrolytic cells which use electricity to perform chemical reactions.

9.5.2: Oxidation occurs at the positive electrode (the anode), and reduction occurs at the negative electrode (the cathode), in a voltaic cell. 9.5.3: Due to the fact that the electrical potential for electrolysis is negative, this reaction will not proceed spontaneously, so an external power source (i.e., a battery) must be applied to force the reaction to occur. 9.5.4: The cation is reduced to form solid metal, and the anion is oxidized to form a gas, typically. The Pb2+ ions go to the cathode where they pick up electrons and become Lead atoms Pb2+ + 2e Pb The Br- ions go to the anode where they lose electrons and become Bromine molecules 2BrBr2

19.1.1: The standard hydrogen electrode consists of a solution of H3O ions at 1 mol dm-3 in a beaker. Placed into this is a platinum electrode surrounded by a gas tube submerged in the solution, with hydrogen gas bubbling over it at 1 atm inside. The circuit to the other half cell is then attached to the platinum electrode, and a salt bridge saturated in potassium chloride. The entire process should take place at 298K and 1 atm pressure.

19.1.2: The potential difference between a given half cell (at 1 mol dm3 conc) and the standard hydrogen electrode. 19.1.3: The potential difference between two half cells (if one half equation is reversed and the two equations are added, the cell potential will be given. It should be positive if you reversed the right one, if it's negative the reaction occurs in the opposite direction to the one you're writing. 19.1.4: Most reactions with positive E-zero values will occur, however it is possible that under non-standard conditions reactions may not occur, or that some reactions may have very high activation energy, meaning the will not occur at any significant rate. Due to the fact that the electrical potential for electrolysis is negative, this reaction will not proceed spontaneously, so an external power source (i.e., a battery) must be applied to force the reaction to occur. 19.2.1: Electrolysis is where the above cells are forced to run in reverse by attaching an electricity source to overcome the potential difference. In aqueous solutions water is also present, and will sometimes be oxidized/reduced in preference to the dissolved salts (or whatever). It is possible to use the standard electrode potentials to predict this, in that species above water (when it is on the left) will not be oxidized, and species below water (on the right) will not be reduced in an aqueous solution. If necessary, this can be checked by working out the cell potential for all possible combinations (involving the, presumably, two elements and water). The reaction with the smallest negative potential difference will be the one which occurs. Highly concentrated solutions may overcome this to some degree however (i.e. it is possible for Cl2 to be oxidized in a concentrated solution). 19.2.2: Faraday's law states that the mass of product produced will be proportional to the charge passed. (Note: The equation charge = current x time , or q=It may be necessary/useful). Faraday's law may also be restated as "The number of faradays required to discharge 1 mol of an ion at an electrode equals the number of charges on that ion". 19.2.3: In the electrolysis of an aqueous solution, as the cations are reduced to from solid metals, these are deposited at the cathode. By connecting a metal object to the cathode, it is possible to plate the deposited metal on the object.