Chapter 23 Electrochemistry

Slides:



Advertisements
Similar presentations
Chpater 4 Oxidation-Reduction
Advertisements

Chapter 20 Electrochemistry
Electrochemistry Chapter 20.
Unit 11- Redox and Electrochemistry
Created by C. Ippolito March 2007 Updated March 2007 Chapter 22 Electrochemistry Objectives: 1.describe how an electrolytic cell works 2.describe how galvanic.
Topic: Electrochemical Cells Do Now: 5 color pencils.
Oxidation Reduction Reactions
Electrochemical Cells
Cells and Voltage.
Cells and Voltage.
ELECTROCHEMISTRY Chapter 17. W HAT IS ELECTROCHEMISTRY Electrochemistry is the science that unites electricity and chemistry. It is the study of the transfer.
Oxidation Reduction Chemisty: Redox Chemistry
Lesson 2. Galvanic Cells In the reaction between Zn and CuSO 4, the zinc is oxidized by copper (II) ions. Zn 0 (s) + Cu 2+ (aq) + SO 4 2-  Cu 0 (s) +
Electrochemistry 18.1 Balancing Oxidation–Reduction Reactions
1 Electrochemical Cells: The Voltaic Cell Mr. ShieldsRegents Chemistry U14 L03.
Aim: What are electrochemical cells?
Redox: Oxidation and Reduction Definitions Oxidation: loss of e- in an atom increase in oxidation number (ex: -1  0 or +1  +2)  Reduction: gain of.
Electrochemistry Electrons in Chemical Reactions.
Chapter 26. An electrochemical cell A device that converts chemical energy into electrical energy. A Daniell cell is a device that could supply a useful.
Aim Redox 1 – Why is redox so important in your life?
Chapter 22 REDOX.
Oxidation-Reduction Reactions LEO SAYS GER. Oxidation and Reduction (Redox) Electrons are transferred Spontaneous redox rxns can transfer energy Electrons.
Electrochemistry. Electrochemical Cells  Electrons are transferred between the particles being oxidized and reduced  Two types –Spontaneous = Voltaic.
GALVANIC AND ELECTROLYTIC CELLS
ELECTROCHEMICAL CELLS
Electrochemistry.
 Deals with the relation of the flow of electric current to chemical changes and the conversion of chemical to electrical energy (Electrochemical Cell)
1 Chapter Eighteen Electrochemistry. 2 Electrochemical reactions are oxidation-reduction reactions. The two parts of the reaction are physically separated.
Chapter 20 Electrochemistry and Oxidation-Reduction.
Electrochemical Cells - producing an electric current with a redox reaction.
1 Electrochemistry. 2 Oxidation-Reduction Rxns Oxidation-reduction rxns, called redox rxns, are electron-transfer rxns. So the oxidation states of 1 or.
CONTENT OBJECTIVE make qualitative or quantitative predictions about galvanic (voltaic) cells based on half-cell reactions and potentials and analyze these.
Mr. Chapman Chemistry 30 ELECTROCHEMICAL CELLS AND REDOX REACTIONS.
ELECTROCHEMICAL CELLS In redox reactions, there is a chemical reaction and an exchange of electrons between the particles being oxidized and reduced. An.
Electrochemistry - Section 1 Voltaic Cells
Topic: Redox Aim: What are electrochemical cells? Do Now: Which of the following ions is most easily reduced? 1)Li+ 2) K+ 3) Ca 2+ 4) Na+ HW:
Galvanic Cells ELECTROCHEMISTRY/CHEMICAL REACTIONS SCH4C/SCH3U.
Unit 2: Electrochemistry Electrolysis
Chapter 19 Last Unit Electrochemistry: Voltaic Cells and Reduction Potentials.
Section 14.2 Voltaic Cells p Voltaic cells Voltaic cells convert chemical energy to electrical energy. In redox reactions, oxidizing agents.
9.2 Electrochemical cells. Two types of electrochemical cells Voltaic cell Spontaneous Chemical  Electrical Uses activity differences between two metals.
ELECTROCHEMICAL CELLS. ELECTROCHEMISTRY The reason Redox reactions are so important is because they involve an exchange of electrons If we can find a.
Electrochemical Cells. Electrochemical Electrochemical cells are a way of storing chemical potential energy. When batteries operate, electrons in high.
CE Chemistry Module 8. A. Involves electron changes (can tell by change in charge) Cl NaBr 2NaCl + Br 2 B. Oxidation 1. First used.
Chapter 19: Electrochemistry: Voltaic Cells Generate Electricity which can do electrical work. Voltaic or galvanic cells are devices in which electron.
1 REVERSIBLE ELECTROCHEMISTRY 1. Voltaic Or Galvanic Cells Voltaic or Galvanic cells are electrochemical cells in which spontaneous oxidation- reduction.
9.2 Electrochemical Cells
Topic 19 Oxidation and Reduction. 1)What is the oxidation number of P in PO 4 -3 ? 2)If Cu and Zn and connected, which is the anode? 3)What reaction (oxidation.
Assigning Oxidation Numbers RULESExamples 2Na + Cl 2  2NaCl Na = 0 or written Na 0 Cl 2 = 0 or written Cl 2 0 RULESExamples 1. Each Uncombined Element.
Electrochemistry Chapter 18. Electrochemistry –the branch of chemistry that studies the electricity- related application of oxidation-reduction reactions.
You will have to completely label a diagram to look like this
Chapter 20 Electrochemistry
Electrochemical Cells
Redox in Action: Voltaic cells
Voltaic Cells Aim: To identify the components and explain the functions of an electrochemical (voltaic) cell.
Electrochemical cells
Electrochemistry RedOx: Part Deux.
Chapter 10 Electrolytic Cells 10.7.
Electrochemistry.
10.2 Electrochemistry Objectives S2
Redox #’s 1-5 #1) The reaction absorbs energy, therefore it is electrolytic (A). #3) Electrolysis requires an external power source (A). #4) Reduction.
Electrochemistry.
Chapter 10 ELECTROLYTIC CELLS 10.7.
Chemistry/Physical Setting
You will have to completely label a diagram to look like this
Redox & Electrochemistry.
Electrochemistry Lesson 3
AP Chem Get HW checked Work on oxidation # review
Chapter 21: Electrochemistry
What is a redox reaction?
Presentation transcript:

Chapter 23 Electrochemistry

Sections 23.1-23.2 Electrochemical Cells OBJECTIVES: Describe how RedOx rxns produce useful electricity Explain the structure and function of Voltaic (Galvanic) Cells [i.e. batteries]

The Nature of Voltaic (Galvanic) Cells You have already seen that when a strip of metal is placed in a solution containing a less active metal, a single replacement rxn will occur This is a classic example of a RedOx rxn. Whether the process is spontaneous or not can easily be predicted by using Table J in your Reference Tables for Physical Setting / CHEMISTRY

Voltaic/Galvanic Cells If the half-rxns which define a RedOx rxn are allowed to occur in separate beakers, the electrons can be made to flow through an external wire and used to perform work The problem is that as the RedOx rxn tries to proceed there is an imbalance of ions in the beakers. Nature will not allow this, so we must supply ions from an external source to keep each solution electrically neutral. The external source of ions is called a “Salt Bridge”. It is composed of a salt- saturated gel contained by a glass U-tube.

The Daniell Cell The Daniell Cell was the first “wet cell” battery. It is composed of a copper electrode in a copper (II) sulfate solution, a zinc electrode in a zinc sulfate solution, a salt bridge (usually containing sodium chloride), an external wire, and a voltmeter. The electrons spontaneously flow from the Zn electrode to the Cu electrode. How the cell works is described in the next slide.

Daniell Cell Half-Rxns Oxidation Half-Rxn: Zn(s) → Zn+2(aq) + 2e-1 ANODE (-) Reduction Half-Rxn: Cu+2(aq) + 2e-1 → Cu(s) CATHODE (+) An Ox, Red Cat

So What Else Happens in the Daniell Cell? The electrons flow from the Zn electrode to the Cu electrode The two RedOx rxns happen simultaneously The cation in the salt bridge moves toward the cathode (Cu Electrode) The anion in the salt bridge moves toward the anode (Zn Electrode) The Zn electrode gets lighter in mass The Cu electrode gets heavier in mass The rxn stops (the battery is dead) when the salt in the salt bridge runs out, or the Zn electrode is used up, or the Cu+2 ions run out

Cell Voltages Standard cell conditions are defined as: 1 Molar Solute, 25◦ C, 1 Atm A “Standard Hydrogen Electrode” under standard conditions has a back voltage applied so the observed cell voltage appears to be zero (see diagrams in the next slides) If a stated voltage (E) has a zero superscript (E◦), the experiment was done under standard conditions

(a) The Standard Hydrogen Cell and (b) Close up of the Hydrogen Standard Electrode.

Cell Voltages You text book has a table of Reduction Potentials (voltages) on pg. 688 Find the two half reactions for your cell; in this case we have: Cu+2(aq) + 2e-1 → Cu(s) E◦ = +0.34 V Zn+2 (aq) + 2e-1 → Zn (s) E◦ = -0.76 V Since you must have both a Red and an Ox rxn, turn the half rxn with the smaller voltage around and change the sign of the voltage (next slide)

Cell Voltages (continued) In the present case we have: Cu+2(aq) + 2e-1 → Cu(s) E◦ = +0.34 V Zn(s) → Zn+2 (aq) + 2e-1 E◦ = +0.76 V If the number of electrons on each side is the same simply add the half rxns together and simplify; the voltages are also added together in a similar fashion: Cu+2(aq) + Zn(s) → Zn+2 (aq) + Cu(s) Overall Rxn +0.34 + (+0.76) = +1.10 Volts E◦cell

The Daniell Cell under Standard Conditions (notice the cell voltage!)

Another Example (pg. 1 of 5) PROBLEM: Suppose someone gave you Al(s), Zn(s), Al(NO3)3(aq), Zn(NO3)2(aq), and NaC2H3O2 [as a paste]… construct a Voltaic cell and label or explain all components SOLUTION: First, identify the electrodes (usually solids), then immerse them in their appropriate electrolytes, and let the paste be the salt bridge… so in this case we have – Al(s) in Al(NO3)3(aq) Zn(s), in Zn(NO3)2(aq) NaC2H3O2 [as a paste] is in the salt bridge

Another Example (pg. 2 of 5) Now write the Reduction Half Rxns using pg. 688: Al+3(aq) + 3e-1 → Al(s) E◦ = -1.66 V Zn+2(aq) + 2e-1 → Zn(s) E◦ = -0.76 V You must have the same number of electrons on both sides of the arrow, so multiple the first rxn by 2 and the second rxn by 3. The voltages, however, are not changed: 2Al+3(aq) + 6e-1 → 2Al(s) E◦ = -1.66 V 3Zn+2(aq) + 6e-1 → 3Zn(s) E◦ = -0.76 V

Another Example (pg. 3 of 5) Turn the smaller voltage rxn around, change sign, and add: 2Al(s) → 2Al+3(aq) + 6e-1 (Ox) E◦ = +1.66 V 3Zn+2(aq) + 6e-1 → 3Zn(s) (Red) E◦ = -0.76 V 2Al(s) + 3Zn+2(aq) → 3Zn(s) + 2Al+3(aq) Over All Net Ionic Rxn (balanced by charge & mass!) E◦cell = +0.90 V = (+1.66 V) + (-0.76 V) Note that all cell voltages must be positive!

Another Example (pg. 4 of 5) Thus, we know: Electrons flow from Al(s) → Zn(s) The cell voltage is +0.90 V Reduction occurs at the Zn electrode, and it is the cathode (+) Oxidation occurs at the Al electrode, and it is the anode (-) Na+1(aq) from the salt bridge flows to the cathode (Zn) C2H3O2(aq)-1 from the salt bridge flows to the anode (Al) The Zn electrode gets heavier while the Al electrode gets lighter

Another Example (pg. 5 of 5)

Other Batteries : Lead Storage Cell

Other Batteries : Standard Dry Cell

Other Batteries : Alkaline Battery

Section 23.3 Electrolytic Cells & More OBJECTIVES: Describe how RedOx rxns can be used to electroplate Discuss some other practical applications of electrochemistry

Electrolytic Cell Basics Electrolytic Cells require an external D.C. power supply The power supply forces a RedOx rxn to take place backwards (rxn normally has a positive DG◦ or a negative E◦) Usually, there is only one beaker containing both the electrolyte and electrode AnOx & RedCat still works .. but the cathode is now (-) and the anode is now (+). This is backwards compared to a Voltaic Cell!

Electroplating of Silver Metal The oxidation of sliver metal has an E◦ value of -0.80 V, so it is not a spontaneous reaction. The external power source supplies this voltage to drive the rxn as seen in the adjacent diagram. The silver electrode is oxidized (loses weight) and the spoon is plated by the reduction of silver ions.

The Formation of Rust Water acts as the medium for the RedOx rxn between solid Fe (oxidized) and molecular oxygen (reduced) + water. The rxn causes a pit to form that will eventually go all the way though the metal. Some metals, such as Al or Ag, form a protective oxidized layer that prevents pitting.

How to Prevent Rust Coat the metal with paint or lacquer to seal out oxygen Use a “sacrifice metal” This implies a metal that oxidizes (rusts) easier than the metal you which to protect, is allowed to rust such that the electrons it produces are fed into the protected metal (see next slide) Apply an external D.C. voltage This is essentially the same as the previous method, but instead of a sacrifice metal a battery is used

One Type of Cathodic Protection

Tarnishing Some metals, such as aluminum and silver, form a thin layer of oxide (rust), which protects the metal from corroding or pitting all the way through Other metals, such as iron, have no such protection and rusts all the way through to the other side Sorry… this is the last time we will use the “slide notes” this year!!!