Electrochemical Cells Using The Nernst Equation

A fuel cell is, in principle, a very simple electrochemical device. The chemical reaction that powers hydrogen fuel cells is the same as that which occurs when hydrogen burns. The chemical equation for this reaction is: 2H2 + O2 ( 2H2O + energy. "Normally hydrogen burns, reacting with oxygen from the air, producing water, heat and light. ... In the fuel cell the chemical reaction is exactly the same, but instead of producing light and heat energy, electrical energy is produced."2 All fuel cells consist of an electrolyte (a substance that allows only the passage of ions) sandwiched between two electrodes. When a fuel containing hydrogen is passed over the negative electrode, otherwise known as an anode, it is ionized. Ionization of the fuel, often accomplished with the assistance of a catalyst, removes electrons from the hydrogen creating positively charged hydrogen ions and negatively charged free electrons. Since only the ions can pass through the electrolyte situated between the electrodes, the electrons must find another route to the positive electrode or cathode, where they will be reunited with the hydrogen ions and combined with oxygen atoms to form water. The electrons passing around the electrolyte constitute an electric current, and thus can be used to provide power during their journey from anode to cathode.3

In Part 1 of the lab, a solar cell was created and tested for its capability to conduct electricity. After researching the processes that contribute to the conductive property, it was found that the oxidized substance is the dye, as it donates an excited electron to the titanium oxide. Consequently, titanium oxide is reduced before it donates an electron to the cathode. The electrolyte solution was found to replenish the dye with electrons so it could continue to act as a reducing agent.

Redox reactions are an important class of reactions in organic chemistry that involve the transfer of electrons from

82. Oxidation occurs when there is a removal of electrons and/or hydrogen atoms from a

The purpose of this experiment was to investigate the chemical reactions involved with oxovanadium ions VO3- and VO2+ to help determine the standard electrode potential for the reaction: VO3− + 4H+ + e− →VO2+ +2H2O. This procedure helped determine the theoretical Eocell.

Oxidation-reduction reactions can be used to stereochemically control and produce many different organic molecules. The oxidation step in this process increases the number of carbon oxygen bonds by losing a hydrogen and breaking

#1) What is an electrochemical gradient, and why is it so important when discussing the movement of ions across cell membranes?

The chemical reaction which occurs between the electrolyte and the positive electrode internally in the battery then produces a buildup of positive ions, which are atoms that are missing electrons. The positive ions buildup at the cathode of the battery which

Reducing Agent - Electron donor in redox reactions. Oxidizing Agent - Electron recipient in redox reactions. NAD+

In the 18th century Italian Physicist Alessando Volta developed the electrolytic cell. Volta experimented with electric current, becoming the first to stimulate the auditory system electrically. In which he attached two metal rods to a battery of approximately 50V and inserted one rod in each ear. When the circuit was completed, he felt a sensation

Oxidation can be defined as a gain in oxygen or loss of hydrogen, loss in electrons, and therefore loss of potential energy. Reduction is gain of hydrogen, or gain in electrons, and gain of potential energy. The organic food molecules are OXIDIZED to CO2. This is the process of glycolysis and making energy. The carbon molecules gain oxygen and are oxidized, and release all that energy to generate ATP. NADH is the REDUCED from of NAD+, because of the gain in hydrogen, or gain in electrons.

Chemical reactions involving the transfer of electrons from one reactant to another are called oxidation-reduction reactions. In a redox reaction, two half-reactions occur; one reactant gives up electrons (undergoes oxidation) and another reactant gains electrons (undergoes reduction). In the case for galvanic cells, redox reactions occur spontaneously. A measure of the tendency for a reduction to occur is its reduction potential (E), measured in units of volts. At standard conditions, 298 K and concentrations of 1.0 M, the measured voltage of the reduction half- reaction is defined as the standard reduction potential (E°).

Step #4: Balance the charges. We can’t cancel out the number of electrons that we’ve added because it is not the same with the oxidation half- reaction and reduction half- reaction. Multiply the 1e to the common factor so that it will be equal to 2e.

When several batteries were built using a Cu/Cu2+ half-cell with a variety of metals, the half-cell containing Mg/Mg2+ would generate the most potential (2.71V) according to the standard reduction potential table. The experiment proved this correct. The Cu/Cu2+ and Mg/Mg2+ cell generated 1.74 V. This occurred because the magnesium had the least reduction potential and reacted as the anode so its reduction number changed signs from -2.37 to +2.37. The reactions may not have reached the predicted potentials due to the temperature of the lab. The reactions of the standard reduction potential table are set at 25° C, but the room is kept at a lower temperature.

Background Students had been taught ‘Electrochemistry’ as outlined in the IB Chemistry syllabus. Investigation Design an experiment that allows you to investigate a variable affecting the rate of electroplating. Your research question must be focussed and specific and must enable you to carry out your experiment safely and within the allocated time. Safety Show your research question to your teacher. Complete a safety hazard assessment before writing a full plan (a + b). Ensure your teacher approves this. Experiment If your plan is safe you will be allowed