Chapter 14 Gen Chem II
Chapter 14 Gen Chem II AS.030.102
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This 5 page Bundle was uploaded by Juliet Villegas on Sunday February 21, 2016. The Bundle belongs to AS.030.102 at Johns Hopkins University taught by Sunita Thyagarajan in Spring 2016. Since its upload, it has received 55 views. For similar materials see Introductory Chemistry II in Chemistry at Johns Hopkins University.
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Date Created: 02/21/16
Chemistry Exam I: Oxidation: Loss of electrons by a species (increase oxidation #) Reduction: Gain of electrons by a species (decrease oxidation #) Oxidizing Agent: electron acceptor, species is reduced Reducing agent: electron donor, species is oxidized Balancing Redox Reactions: 1) Write out all oxidation numbers to see what’s being oxidized, reduced 2) Write out half reactions 3) Balance Oxygens by adding H2O and balance Hydrogens by adding H+ 4) For acidic solutions: Balance net charges by adding electrons 5) Multiply equations so electrons cancel out 6) Add two half reactions 7) Add OH- to make water and cancel out H+ 8) Cancel out what you can (Especially water molecules) Types of Redox reactions: o Direct: No external circuit (no useful electric current) o Indirect: External circuit (useful current, external wire used) Electrochemical cells: Apparatus that allows a redox reaction to occur by transferring electrons through an external connector (circuit). o Galvanic/voltaic cell: Produces electrical energy. Chemical change produces electric current. DeltaGrxn<0. Product favored reaction. Ex. Batteries o Electrolytic cell: Requires electrical energy. Electrical current used to cause chemical change in a non- spontaneous reaction. DeltaGrxn>0. Reactant favored reaction. Ex. Gold/silver plate jewelry Galvanic cell parts: o Anode on left: Oxidation occurs at this electrode (Zn(s) -> Zn2+(aq)+2e-). Positive charge builds up. o Cathode on right: Reduction occurs at this electrode (Cu2+ (aq)+2e- -> Cu(s). Negative charge builds up. o Electrons flow in the wire from anode to cathode o Ammeter: measures current (between wires) Electrolytic cell: Increase the electrostatic potential energy of electrons in the cathode to make it flow towards the anode, so anode becomes cathode and vice versa. Energetics: o Faraday expressed quantitative relationship between amounts reacted and totally electrical charge passing through the cell. o Faraday’s Laws (of electrolysis): 1) The mass of a given produced or consumed at an electrode is proportional to quantity of electrical charge passed through cell. 2) Equivalent masses of different substances are produced or consumed at an electrode by passing of a given quantity of electric charge. o Charge on single electron (e)= 1.60217646 X 10-19 C (Multiply by Avogadro’s # for Faraday’s) o Faraday constant: F= 96,485.34 C mol-1 o Electric current I (units: amperes)=amount of charge Q (coulombs) flowing through circuit per unit time (I = Q/t) o Moles of electrons: It/96,485 C mol-1 o Change in energy (Delta energy) = Q*V Q=charge V=voltage Q= 1C = 1.602x10-19 C V= 1 volt DeltaEnergy = 1 J = 1 eV 1 eV = 1.602 x 10-19 J The Gibb’s Free Energy and Cell Voltage: o Work: welec = -QDeltaE = -ItDeltaE DeltaE positive for galvanic cell o At constant T&P: -welec,max = |DeltaG| o DeltaG = welec = -QDeltaE = -nFDeltaE (reversible)=q+welec-TdeltaS o Qrev=TdeltaS o Electrical work produced only if DeltaG<0 (DeltaE>0) Standard State: o DeltaG^o=-nFDeltaE^o DeltaE^o is cell voltage (potential difference) of galvanic cell in which reactants and products are in their standard states Cell Potential: o Electrons move from anode to cathode by electromotive force (emf) which is change in potential o Anode (negative electrode) is the supplier of electrons o Cathode (positive electrode) is the acceptor of electrons o Anode, standard is flipped and becomes negative for cell potential o Cell potential is cathode-anode o Best oxidizing agents have higher reduction potential o Best reducing agents have low reduction potential o Lower reduction potentials can reduce those with higher ones, but not vice versa o MAKE SURE THINGS ARE BALANCED!!! o Computing cell voltage from free energy change: o o To get half-cell voltage using half potentials, ((mol)(cathode V)-(mol)(anode V))/(mol product) Mol is electrons transferred Reduction potential diagrams: single substance gets oxidized and reduced. Occurs if reduction on right is larger than on left Reduction potential diagram: Concentration and Nernst Equation Electrolyte PH?? Equilibrium Constants: Ion Selective Probes: respond to presence of specific ion o Most common is pH probe sensitive to H+ ions How to get pH: Galvanic cells arranged in series form a battery of cells o Primary cells discarded when electrical energy discharged o Secondary cells can be recharged o Dry cell: Disadvantage Concentrations change w/ time and voltage falls with use Replace NH4Cl with KOH o Alkaline dry cell: Dissolved species doesn’t appear in overall reaction, so steadier voltage o Rechargeable batteries: Can undergo many cycles of charge-discharge, and energy/kg low for lead-acid battery o Corrosion factors: oxygen, water, salt (electrolyte), and acidity o “Crabpot” prevents corrosion with sacrificial anode like Zn, Al, and Mg (they get consumed instead because oxidize more easily) o Paint inhibits but can chip o Passivation is a protective oxide layer, stainless steel o Electrometallurgy: Most metals occur naturally as oxide. Can react oxides with coke to reduce metal. Electrolytic reduction works when smelting doesn’t
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