Read Sections 8.3, and 8.4 before viewing the slide show.

Unit 29 Electrochemistry (Chapter 8) Description of an Electrochemical Cell (8.3) Electrochemistry Terminology (8.3) Electrochemistry as Applied to Batteries (8.3) Corrosion (8.4)

Electrochemistry (8.3) Electricity is due to the motion of electrons. Since oxidation-reduction reactions involve an exchange of electrons, such reactions can be used to generate electricity through applications such as batteries. Image from

Electrochemistry Cont. (8.3) In the reaction from the previous page, the copper goes from being in the elemental form to dissolving in solution as Cu 2+ ions which cause the blue color in solution. The Ag + ions, previously dissolved in solution, become elemental silver and form the “hangings” seen in the second beaker. In equation form: Cu (s) → Cu 2+ (aq) + 2 e - Ag + (aq) + e - → Ag (s) An important aspect in understanding electrochemistry is to understand how these two equations may be combined to form the overall reaction.

Electrochemistry Cont. (8.3) Each of the reactions below is called a half-reaction. One represents an oxidation and the other a reduction. Cu (s) → Cu 2+ (aq) + 2 e - Ag + (aq) + e - → Ag (s) Since the electrons donated by the copper are the ones accepted by the silver, the number of electrons being accepted and donated must match. In order for the electrons to balance, each half-reaction is multiplied by an integer as necessary to ensure that the number of electrons donated matches those accepted. In this example, the Cu equation involves two electrons and the Ag equation involves only one so multiplying the Ag equation by 2 will give two electrons accepted to go along with the two electrons donated by copper (continued on next page).

Electrochemistry Cont. (8.3) Multiplying the first equation by “1” and the second by “2” gives: 1 × (Cu (s) → Cu 2+ (aq) + 2 e - ) 2 × (Ag + (aq) + e - → Ag (s) ) These simplify to: Cu (s) → Cu 2+ (aq) + 2 e - 2 Ag + (aq) + 2 e - → 2 Ag (s) Adding these two equations gives: Cu (s) + 2 Ag + (aq) → Cu 2+ (aq) + 2 Ag (s)

Electrochemical Cells – Terminology (8.3) Semipermeable Membrane – only allows solvent and nitrate ions to pass through. Anode Cathode Copper Silver Cu 2+ SO 4 2- Ag + NO 3 - Electrons Rather than carrying out the previous reaction in one container, the two halves of the reaction may be separated to allow the electrons to transfer externally to the cell – see the figure to the right. The semipermeable membrane allows the nitrate ions to transfer through to the left while the electrons transfer to the right through the wire at the top. The copper metal connected to the wire is called an electrode – specifically an anode since that is where oxidation occurs. The silver metal is another electrode called the cathode – the electrode at which reduction occurs.

Image from Electrochemical Cells – the Implementation (8.3) The figure below illustrates the construction of a dry cell typically used in flashlights and other portable devices. A simplified version of the reaction that occurs is: Zn + 2 MnO 2 + H 2 O → Zn 2+ + Mn 2 O OH - Alkaline cells replace use KOH in the paste – these are typically more expensive but last longer. Can you tell which substance is oxidized – is it zinc or manganese dioxide?

Image from The Lead Storage Battery (8.3) The lead storage battery is commonly found in cars and boats. It is a rechargeable battery though it is quite heavy and involves corrosive materials. The typical 12-volt lead storage battery is made of six cells of two volts each. During the discharge of a lead storage battery (starting your car) the net reaction is: Pb + PbO H 2 SO 4 → 2 PbSO H 2 O During the recharging, while the car is running, the reverse reaction occurs through the action of the car’s alternator.

Image from Corrosion (8.4) Estimates are the corrosion in the US alone costs about $276 billion per year. Approximately 20% of iron and steel production annually in the US is used to replace corroded items. In the corrosion process, iron metal is initially oxidized to Fe 2+ while oxygen in the air is reduced to the hydroxide ion. This ultimately leads to iron (III) hydroxide, which is the material commonly identified as rust. Electrons transferred in this process through the metal itself, but an electrolyte is required to complete the circuit. Thus, corrosion is more prevalent in northern climates in which salt is used on the roads and in areas near salt water. Often another metal that is more easily oxidized is used as a “sacrifical” anode. Such a material is destroyed preferentially to the structural metal and is easily replaced. See next slide.

Image from Corrosion Example

Image from Another Corrosion Example

Image from Sacrificial Anodes The small metal ingots (some highlighted in the image below) are called “sacrificial” anodes. In a salt water environment, the “sacrificial” anodes will be destroyed prior to the hull of the ship. Occasional replacement of the anodes is a relatively simple and inexpensive task that does not affect the integrity of the hull. Sacrificial Anodes