Electrochemistry

The movement of electrons between chemical species is the heart of chemical reactions. A. Redox Reactions Reduction- chemical process where electrons are GAINED Oxidation - chemical process where electrons are LOST. - LEO goes GER (Lose Electrons is Oxidation / Gain Electrons is Reduction) Redox (reduction/oxidation) reactions - Are coupled chemical processes where electrons are transferred from one substance to another -oxidizing agent (oxidant) - the chemical species which oxidizes another (is reduced itself) -reducing agent (reductant)- the chemical species which reduces another (is oxidized itself) B. Oxidation states (numbers) The oxidation state of an atom (or bonded atoms) reflects the degree of oxidation of the chemical substance, usually written as an integer value with positive, negative or no charge. Atoms that gain or lose electrons become ions. The charge of the ion is a true representation of its oxidation state. The formation of covalent bonding may not produce ions but can change the distribution of electrons around the atoms which can still be represented by an oxidation state. C. Assigning oxidation numbers D. Writing & balancing redox reactions Half-reactions show the individual reduction and oxidation reactions as separate equations. These are written as ionic equations (showing charges) Voltaic cells (Galvanic cells) [|Voltaic cells] are electrochemical cells that harvest energy from redox reactions. The spontaneous redox processes move electrons through an external path instead of passing directly between redox partners. Components/characteristics: cells- aqueous solutions of different salt composition or concentration external circuit- provides a path for electron to move which always move from the anode to the cathode electrodes- solid metals that are used as redox partners which are connected to an external circuit -cathode-- reduction, higher electronegativity, increases in mass, positive electrode (e- flow towards) -anode-- oxidation, lower electronegativity, decreases in mass, negative electrode (e- flow away from) salt bridge- a porous connective apparatus which allows for the movement of ions between the cells. This provides a way to keep neutral solutions. -anions migrate toward the anode & cations migrate toward the cathode Video :[| Constructing a Voltaic Cell] Resource: [|Activity Series of metals] FLASH - [|Galvanic Cell] Khan Academy : [|Galvanic Cell] Electromotive force or cell potential (Emf or Ecell) measures the potential energy difference of electrons between the two redox partners (electrodes). -electrons tend to spontaneously move to a lower energy state --> lower electronegative metal to a higher electronegative metal (e.g. zinc to copper) Each electrode has a specific reduction potential based on the metal cations (or nonmetal in solution) ability to gain electrons. The greater the ability an atom/anion has for gaining eletrons the greater its reduction potential. Emf is measured in volts. A volt is the potential energy difference associated with 1 coulomb of charge imparting 1 joule of energy. Resource: [|Electrical Potential] Resource: [|Standard Reduction Potentials] All reduction potentials are standardized to hydrogen ions gaining electrons, this is called the [|standard hydrogen electrode] potential Standard conditions for relative comparison are defined as 1M concentration for all solutions, ambient temperature (298.15 K) & 1 atm pressure. The standard reduction potential for any material is identified as Eo and the standard cell potential is Eocell Important points 1. All half-reaction potentials are written as reduction reactions 2. Calculating cell potentials are not stoichiometrically significant. We don't consider coefficients in calculations 3. Don't change signs of the standard reduction potentials. The equation is compensating for the oxidation reactions (minus instead of plus--remember Hess's Law) 4. The more positive the value of Eored, the greater the driving force for reduction. Cathodes are defined by the large Eored.
 * The oxidation number for an atom in its elemental form is zero
 * Fe: The oxidation number of Fe = 0
 * Cl2 : The oxidation number of Cl = 0
 * The oxidation number of a monoatomic ion = charge of the monatomic ion
 * Oxidation number of S-2 is -2.
 * Oxidation number of Al+3 is +3
 * The oxidation number of all Group 1A metals = +1 (unless elemental).
 * The oxidation number of all Group 2A metals = +2 (unless elemental).
 * Hydrogen (H) has two possible oxidation numbers:
 * +1 when bonded to a nonmetal
 * -1 when bonded to a metal
 * Oxygen (O) has two possible oxidation numbers:
 * -1 in peroxides (O2-2)
 * -2 in all other compounds
 * The oxidation number of fluorine (F) is always -1.
 *  Negative oxidation numbers are commonly limited to nonmetals.
 * <span style="font-family: 'Times New Roman'; font-family: 'Times New Roman',Times,serif; font-size: medium;"> Elements commonly exhibit positive oxidation numbers only when combined with more electronegative elements.
 * <span style="font-family: 'Times New Roman'; font-family: 'Times New Roman',Times,serif; font-size: medium;">The sum of the oxidation numbers of all atoms (or ions) in a neutral compound = 0.
 * <span style="font-family: 'Times New Roman'; font-family: 'Times New Roman',Times,serif; font-size: medium;">H2SO3 -- H = +1, S = +4, O = -2
 * <span style="font-family: 'Times New Roman'; font-family: 'Times New Roman',Times,serif; font-size: medium;">H2SO4 -- H = +1, S = +6, O = -2
 * <span style="font-family: 'Times New Roman'; font-family: 'Times New Roman',Times,serif; font-size: medium;">The sum of the oxidation numbers of all atoms in a polyatomic ion = charge on the polyatomic ion
 * =  ||= 1 coulomb (C) ||= 6.85 x 1018 electrons ||
 * =  ||= 1 electron (e-) ||= 1.60 x 10-19 C ||
 * =  ||= 1 volt (V) ||= 1 Joule / 1 Coulomb ( J/V) ||
 * =  ||= 1 Joule (J) ||= 1 kg m2 / s2 ||
 * =  ||= 1 Coulomb ||= 1 ampere x 1 second (A·s) ||