Solutions of Electrolytes

Electrochemistry

Electric Current Energy is released when electrons

are transferred This can be heat energy, or electric

energy Zn + CuSO4 example – heat energy When separated by a barrier –

electrical energy is produced

Electric Energy By using a porous barrier to

separate the reactants, electrons being exchanged between ions are forced to travel through a connecting wire, producing electricity This type of set-up is called an

electrochemical cell

Voltaic Cells A voltaic cell is an electrochemical

cell in which a spontaneous redox reacton produces a flow of electrons through an external circuit

A circuit is a closed loop path for current to flow In the wire the negative electrons flow

from anode to cathode

Parts of the Voltaic Cell The voltaic cell is actually composed of

two halves, each called a half-cell Anode – the source of the electrons, it is

the electrode at which oxidation occurs as electrons are lost by a substance

Cathode – accepts the electrons, it is the electrode at which reduction occurs as electrons are gained by a substance

Inner workings of the voltaic cell Any element that loses electrons has

undergone oxidation. This takes place at the anode

Any element that gains electrons has undergone reduction, this takes place at the cathode

Oxidation and reduction must occur together

Any reaction in which electrons are gained or lost in equal numbers is a redox reaction Zn + Cu2+ Zn2+ + Cu

Redox in the voltaic cell In the cell, redox reactions result in

a continuous flow of electrons from anode to cathode

Assignment 1-5 page 580

Lab 1 Purpose:1. To create a chemical cell (voltaic

cell) 2. To understand how a chemical cell

works

Materials:

1.(2) 250 mL beakers

2.A paper towel

3.(1) copper electrode

4.(1) zinc electrode

5.Voltmeter (reading between 0 v and 1.5 v)

6.1M CuSO4

7.1M NaCl

8.1M ZnSO4

Procedure: 1.Attach the Zinc electrode to the negative (common) side of the voltmeter and the Copper electrode to the positive side of the voltmeter. 2.Pour the 1M CuSO4 solution into one 250 mL beaker.

3.Pour the 1M ZnSO4 solution into one 250 mL beaker.

4.Place the Zinc electrode in the ZnSO4 solution.

5.Place the Copper electrode in the CuSO4 solution.

6.Soak the paper towel in the NaCl solution and place one end each beaker. 7.Record the voltage that is created by the chemical cell.

Current The activity series allows you to

predict which ions are best at donating and accepting ions

Half-Reaction – the reaction occurring at each electrode

The amount of electric energy that can be generated by each half-reaction is determined by its reduction potential

Current Standard Reduction Potential- The E0

of a half-cell connected to the standard hydrogen electrode when ion concentrations in the half-cells are 1 M, gases are at a pressure of 1 atm, and the temperature is 25 degrees C

E0 represents the voltages generated by half-reactions

SHE The standard hydrogen electrode (SHE)

has the half-reaction 2 H+ + 2e- H2

The concentration of hydrogen ion is 1 M, the temp is 25 degrees C, and the pressure of the hydrogen gas is 1 atm

Is considered the anode in each cell If E0 is (+) it indicates that hydrogen is more

willing to give up electrons than the metal If E0 is (-) it indicates that the metal is more

willing to give up electrons than hydrogen

The activity series Ordering the potentials of the

metals results in a table of reduction potentials Pg 585 of your text

Calculating E0

E0 cell = E0 cathode - E0 anode Example: calculate E0 for the

zinc/copper cell Zn2+ + 2e- Zn E0 = -0.76 V Cu2+ + 2e- Cu E0 = +0.34

V Reduction of Zn has a lower E0 so Zn

is the anode E0 for the cell is the difference

between the two, +1.10 V

Assignment 6-10 pg 586 Due tomorrow

Oxidation States An Oxidation State (or oxidation

number) is a numerical representation of an atoms share of the bonding electrons In ionic compounds it is equal to the

ionic charge In covalent compounds it is the

average charge assigned to an atom according to electronegativities

Rules for assigning ox states

1. A free element is 02. For an ion it is the ionic charge3. The more electronegative element in a

binary compound is assigned the number equal to the charge it would have if it were an ion

4. H is +1, unless it is combined with a metal, in which case it is –1

5. F is always –1, because it is the most electronegative element

Rules cont.6. O is –2, unless combined with F

(+2), or in a peroxide (-1)7. Al has +1,2,or 3 in compounds,

group 1 and group 2 respectively8. The sum of all ox #’s in a

compound is 09. The sum of all ox #’s in a

polyatomic ion must = the charge of the ion

Redox and ox #’s Changes in oxidation #’s indicate a

redox reaction Any time a pure element appears

as a reactant or product, you have a redox reaction

Oxidation numbers are used to balance redox reactions

Balancing Redox equations OIL RIG

Oxidation Is Loss of electron Reduction Is Gain of electron

The half-reaction method See page 593 in text

Assignment Problems 13-15 page 594

How Batteries Work Batteries use redox reactions to

convert chemical energy into electrical

A battery is a single voltaic cell or group of voltaic cells that are connected together

When designing a battery:

1. Redox rxn must be spontaneous2. Half-reactions must produce the

desired voltage3. If you want a rechargeable

battery, the redox reaction must be easy to reverse

4. Environmental concerns must be addressed

Zinc-Carbon batteries Most common battery Also known as the dry cell battery Uses reaction involving zinc and

manganese dioxide Zinc container is the anode Carbon rod is the cathode

Alkaline batteries Alkaline batteries are zinc-carbon

batteries that use KOH instead of NH4Cl

Has no carbon post Perform better than normal zinc-

carbon batteries

Mercury Batteries Similar to the alkaline battery HgO is reduced at the cathode

The production of liquid mercury is a major disadvantage for mercury batteries

Used mercury batteries must be recycled so that the elemental mercury can be recovered

Rechargeable Batteries The process in which electric

energy is used to drive a redox reaction is electrolysis

The container in which electric energy drives a nonspontaneous redox reaction is an electrolytic cell

Car Batteries Car batteries are lead-acid batteries Contains six cells Produces 12V Lead serves as the anode The electrolytic solution is sulfuric acid

(H2SO4) The voltage produced from the cars

alternator reverses the half reactions and regenerates Pb, recharging the battery

Ni-Cad batteries Are also rechargable Cadmium is the anode NiO(OH) is reduced at the cathode Must be recycled because of the

toxicity of cadmium

Assignment 16-21 page 599

Electroplating Electroplating is the deposition of a

metallic coating onto an object by putting a negative charge onto the object and immersing it into a solution which contains a salt of the metal to be deposited. The metallic ions of the salt carry a positive charge and are attracted to the part. When they reach it, the negatively charged part provides the electrons to reduce the positively charged ions to metallic form.

Electroplating example

Electroplating activity Objective:Electroplate silver on to

copper

Research Paper Write a research paper on the

development and use of rechargable batteries in LEV’s. Format:

Must be typed, single spaced Body must be 2-3 pages Style: MLA

Worth a possible 25 points Due Friday, Dec. 7

Electrochemical cells at nonstandard conditions The Nernst Equation

E = E0 – 0.0257 V ln Q at 25˚ C

n

Q = reaction quotient n = number of moles of electrons

transferred