Electrochemical Reactions The first chemical process to produce electricity was described in 1800 by the Italian scientist Alessandro Volta, a former high school teacher. Acting on the hypothesis that two dissimilar metals could serve as a source of electricity, Volta set up a reaction between two metals (silver and zinc) with different electron affinities. Small plates of silver and zinc were separated by a pad of absorbent material soaked in saltwater. The saltwater would conduct electricity, and therefore allow charge to transfer from one metal to the other if the circuit was completed (by a direct connection between the two metal plates, such as a wire). This type of setup is called a galvanic cell. Galvanic cell: a chemical reaction involving electron transfer, arranged so that the electrons must transfer through an external wire in order to complete the circuit, thus generating electrical current. (The chemical reaction causes the electric current to flow.) To increase the flow of electricity, Volta stacked several of these reactions one on top of another, creating a pile of alternating silver and zinc plates (separated by pads of absorbent material soaked in saltwater). When Volta moistened his fingers and touched the top and bottom metal plates, he experienced small electric shocks. Volta’s creation—the “voltaic pile,” as it came to be called—was the first battery. Battery: a repeating “pile” of galvanic cells. Piling galvanic cells (arranging them in series) generates greater voltage, allowing batteries to act as a practical source of electricity. Within months, William Nicholson and Anthony Carlisle in England attempted to confirm the production of electric charges on the upper and lower plates in a voltaic pile using an electroscope. In order to connect the plates to the electroscope, Nicholson and Carlisle added some water to the uppermost metal plate and inserted a wire to the electroscope. To their surprise, Nicholson and Carlisle observed the formation of a gas, which they identified as hydrogen. Nicholson and Carlisle then filled a small tube with river water and inserted wires from the voltaic pile into each end of the tube. Two different gases were generated, one at each wire. Nicholson and Carlisle had discovered electrolysis—the electrical current caused water molecules to decompose into hydrogen and oxygen gases. Electrolytic cell: a chemical reaction involving electron transfer which only proceeds if an external source of electrical current (electricity) is applied. (The electric current causes the reaction to occur.)

Electrochemistry is the study of the relationship between electrical forces and chemical reactions. The term “redox” is often used interchangeably with electrochemistry. Redox is a shortening (and reversal) of “oxidation-reduction.” This term refers to the fact that electrochemical processes require the simultaneous oxidation of one element and reduction of another.  In an oxidation half-reaction, an element loses electrons. This causes that element’s oxidation number (charge) to increase. (Because electrons have a negative charge, removing some of those negatively charged electrons from an element makes the element more positive.)  In a reduction half-reaction, an element gains electrons. This causes the element’s oxidation number (charge) to decrease. (Because the electrons it gains have a negative charge.)  The electrons lost by the oxidized element are transferred to the reduced element. The reduced element gains electrons from the oxidized element. This is how flow of electrical current is generated.

Batteries in Real Life A regular Duracell or Energizer battery always has a label showing which end is positive and which is negative so that you can line it up correctly in the devices it powers. These two ends, called electrodes, would be the top piece of silver and the bottom piece of zinc in the voltaic pile or the two wires in the water in the Nicholson/Carlisle electrolysis. In chemistry, we specify these two electrodes as the anode (where oxidation occurs in galvanic cells) and the cathode (where reduction occurs in galvanic cells). For a battery (galvanic cell) to run, there must be an electrolyte (a conductive solution) connecting the two electrodes. In both Volta’s pile and the Nicholson/Carlisle electrolysis, salt water served this purpose. (Remember from last unit: ionic substances are best at dissolving in water and conducting electricity.) Modern batteries use a variety of chemicals to power their reactions. Common battery chemistries include: Primary Cells (non-rechargeable batteries): o

Zinc-carbon battery: The zinc-carbon chemistry is common in many inexpensive AAA, AA, C and D dry cell batteries. The anode is zinc, the cathode is manganese dioxide, and the electrolyte is ammonium chloride or zinc chloride.

o

Alkaline battery: This chemistry is also common in AA, C and D dry cell batteries. The cathode is composed of a manganese dioxide mixture, while the anode is a zinc powder. It gets its name from the potassium hydroxide electrolyte, which is an alkaline substance.

Secondary Cells (rechargeable batteries): o

Lithium-ion battery (rechargeable): Lithium Ion is the most common rechargeable battery type for high-performance devices, such as cell phones, digital cameras and even electric cars. A variety of substances are used in lithium batteries, but a common combination is a lithium cobalt

oxide cathode and a carbon anode. (Nickel-metal hydride (NiMH) and nickel-cadmium (NiCd) batteries were also once very prevalent, but have become less common as LiOn batteries have become easier to produce.) o

Lead-acid battery (rechargeable): This is the chemistry used in a typical car battery. The electrodes are usually made of lead dioxide and metallic lead, while the electrolyte is a sulfuric acid solution.

Non-rechargeable batteries, or primary cells, and rechargeable batteries, or secondary cells, produce current exactly the same way: through an electrochemical reaction involving an anode, cathode and electrolyte. In a rechargeable battery, however, the reaction is reversible. When electrical energy from an outside source is applied to a secondary cell, the negative-to-positive electron flow that occurs during discharge is reversed, and the cell's charge is restored. (While attached to the charger, the reaction is electrolytic instead of galvanic.) When it comes to rechargeable batteries, not all batteries are created equal. NiCd batteries were among the first widely available secondary cells, but they suffered from an inconvenient problem known as the memory effect. Basically, if these batteries weren't fully discharged every time they were used, they would quickly lose capacity. NiCd batteries were largely phased out in favor of NiMH batteries. These secondary cells boast a higher capacity and are only minimally affected by the memory effect, but they don't have a very good shelf life. Like NiMH batteries, LiOn batteries have a long life, but they hold a charge better, operate at higher voltages, and come in a much smaller and lighter package. Essentially all high-quality portable technology manufactured these days takes advantage of this technology. However, LiOn batteries are not currently available in standard sizes such as AAA, AA, C or D, and they're considerably more expensive than their older counterparts. With NiCd and NiMH batteries, charging can be tricky. You must be careful not to overcharge them, as this could lead to decreased capacity. To prevent this from happening, some chargers switch to a trickle charge or simply shut off when charging is complete. NiCd and NiMH batteries also must be reconditioned, meaning you should completely discharge and recharge them again every once in a while to minimize any loss in capacity due to the memory effect. LiOn batteries, on the other hand, have sophisticated chargers that prevent overcharging and never need to be reconditioned because they do not suffer from the memory effect. Even rechargeable batteries will eventually die, though it may take hundreds of charges before that happens. When they finally do give out, be sure to dispose of them at a recycling facility, as they contain substances that can be dangerous if put in a landfill with regular garbage.

Based on this reading and the information with the questions on the reverse, write definitions for: Electrolytic Cell, Galvanic Cell, Battery, Electrolyte, Oxidation, and Reduction in your notebook for future reference.

Name: Period:

Intro to Batteries and Electrochemistry 1. Name the scientist(s) who used a chemical reaction to generate an electrical current.

2. Name the scientist(s) who used an electrical current to generate a chemical reaction.

3. What two elements were used in Volta’s reaction (the voltaic pile)?

4. What is “piled” in a voltaic pile? (Use the scientific term.)

5. Why do you think Volta did not feel a shock from his reaction until he created a pile?

6. What two elements are produced in the Nicholson/Carlisle reaction? 7.

Write a balanced chemical reaction for the electrolysis reaction performed by Nicholson & Carlisle. (Hint: the only reactant is H2O; there are only two products—both are elements in their diatomic states because they are produced as gases.)

8. In a water molecule, what is the charge of a hydrogen atom?

9. In a water molecule, what is the charge of an oxygen atom?

10. In diatomic hydrogen molecules (H2), the hydrogen atoms have a charge of zero. (Elements always have a charge of zero in their elemental state.) In the reaction below, determine the number of electrons that must be produced to account for the charge difference between the two types of hydrogen. __2__ H1+ + _____ e-  _ 1 _ H20 11. Electrochemical reactions always involve both an oxidation component—in which an element’s charge increases—and a reduction component—in which a (different) element’s charge decreases. Is the reaction in the previous question an oxidation or a reduction?

12. Which type of electrochemical reaction requires an applied current or voltage to proceed?

13. Which type of electrochemical reaction produces a current or voltage when it occurs?

14. Use #10 to determine the answer to circle in each statement: a. Reduction causes charge to increase/decrease. Electrons are left/right of the arrow. b. Oxidation causes charge to: increase/decrease. Electrons are left/right of the arrow. 15. Is a battery a galvanic cell or an electrolytic cell?

16. Given that a battery is a type of chemical reaction, why do batteries eventually “die”?

17. What is the difference between a primary cell and a secondary cell?

18. What is necessary (in terms of the chemical reactions involved) in order for a battery to be rechargeable?

19. What type of battery is used in automobiles?

20. In an alkaline battery: a. What substance is the cathode? b. What substance is the anode? c. What is the electrolyte? 21. What substance served as the electrolyte in Volta’s voltaic pile?

22. What is the “memory effect”?

23. Which types of battery is most affected by the memory effect?

24. Which rechargeable battery is least affected by the memory effect?