Presentation is loading. Please wait.

Presentation is loading. Please wait.

Unit 11- Redox and Electrochemistry

Similar presentations

Presentation on theme: "Unit 11- Redox and Electrochemistry"— Presentation transcript:

1 Unit 11- Redox and Electrochemistry
Anode Cathode Electrochemical cell Electrode Electrolysis Electrolyte Electrolytic cell Half-reaction Oxidation Oxidation number Redox Reduction Salt bridge Voltaic cell

2 What’s the point ? REDOX reactions are important in …
30/09/99 What’s the point ? REDOX reactions are important in … C3H8O + CrO3 + H2SO4  Cr2(SO4)3 + C3H6O + H2O Purifying metals (e.g. Al, Na, Li) Producing gases (e.g. Cl2, O2, H2) Electroplating metals Electrical production (batteries, fuel cells) Protecting metals from corrosion Balancing complex chemical equations Sensors and machines (e.g. pH meter)

3 What is redox? Oxidation- loss of electrons by an atom or ion
Reduction- gain of electrons by an atom or ion **since one can’t occur without the other Combine terms to Redox Mnemonic: LEO the lion says GER Lose Electrons Oxidation Gain Electrons Reduction

4 Oxidation numbers On periodic table
Determines what is oxidized and reduced in a reaction If they change it’s a redox reaction What type of  reaction is this (besides redox)???

5 Assigning Oxidation numbers
Identify the formula If element is free (uncombined) its ox # is 0 Monotomic ions- ox # is same as ion charge Metals in Groups 1,2 and 3 have ox #’s of +1, +2 and +3 respectively Fluorine is always -1 in a compound Hydrogen is always +1 unless it’s combined with a metal then it’s -1 Oxygen is usually -2, except when combined with a more electronegative element then it’s +2 *sum of oxidation #’s in a compound must be 0 *sum of oxidation #’s in a polyatomic ion must equal its charge

6 Try these: HNO3 CO2 K2PtCl6 PCl5 H2SO4

7 Redox reactions Once you determine oxidation numbers you can see what element was oxidized and what was reduced Oxidizing agent- substance that was reduced (gained electrons) Reducing agent substance that was oxidized (lost electrons) Oxidizing agent (element reduced) because it gains electrons from oxidized element so it allowed for oxidation to take place because it was a place for the electrons to go

8 Half-reactions Show oxidation or reduction of redox rx Ex:
Shows conservation of mass and charge Charge does not have to be 0

9 Balancing redox rx’s Assign oxidation numbers to determine what is oxidized and what is reduced. Write the oxidation and reduction half-reactions. Balance each half-reaction. Balance charge by adding electrons. Multiply the half-reactions by integers so that the electrons gained and lost are the same

10 Example: Cu + AgNO3  Cu(NO3)2 + Ag
Add the half-reactions, subtracting things that appear on both sides. Make sure the equation is balanced according to mass. Make sure the equation is balanced according to charge. Example: Cu + AgNO3  Cu(NO3)2 + Ag

11 Practical use for redox reactions
Electrochemical cells Involves a chemical reaction and flow of electrons 2 types: Voltaic- spontaneous Electrolytic- requires electric current (nonspontaneous) Each have 2 electrodes- site of oxidation and reduction Oxidation occurs at the anode Reduction occurs at the cathode An Ox Red Cat Anode- oxidation, reduction-cathode

12 Voltaic cells Once even one electron flows from the anode to the cathode, the charges in each beaker would not be balanced and the flow of electrons would stop

13 Voltaic cells Therefore, we use a salt bridge, usually a U-shaped tube that contains a salt solution, to keep the charges balanced. (completes the circuit) Cations move toward the cathode. Anions move toward the anode.

14 Voltaic Cells In the cell, then, electrons leave the anode and flow through the wire to the cathode. As the electrons leave the anode, the cations formed dissolve into the solution in the anode compartment.

15 Voltaic Cells As the electrons reach the cathode, cations in the cathode solution are attracted to the now negative cathode. The electrons are taken by the cation, and the neutral metal is deposited on the cathode.

16 Activity series helps identify anode and cathode
Metal higher on chart- oxized (anode) Metal lower on chart- site of reduction (cathode)

17 Determining electric potential
Voltmeter is used Voltage is compared to the reduction of H which is 0 volts The more “+” the reading; reduction is more likely E is positive- happens naturally E is negaitve- not spontaneous Standard electrode potentials-Eo= ability for an electrode to gain electrons

18 Reduction potentials for many electrodes has already been measured

19 Cell potentials At standard conditions can be determined using this equation: The strongest oxidizers have the most positive reduction potentials. The strongest reducers have the most negative reduction potentials. Ecell = Ered (cathode) − Ered (anode)

20 Cell Potentials Ered = −0.76 V  Ered = +0.34 V 
For the oxidation in this cell, For the reduction, Ered = −0.76 V Ered = V

21 Cell Potentials Ecell  = Ered (cathode) − (anode)
= V − (−0.76 V) = V

22 Examples of Voltaic Cells: Dry Cells
Dry cells use two electrodes and a “paste” as an electrolyte. Some pastes are acidic and others are alkaline. Carbon is generally used as the cathode and zinc as the anode.

23 Lead-Acid Batteries Lead-Acid batteries usually contain six cells.(2 V each) The battery contains lead plates, lead oxide plates, dividers, and a sulfuric acid electrolyte. The lead plate is the anode and the lead oxide plate is the cathode. Each cell is connected to form one cathode and one anode on the top or side of the battery.

24 Fuel Cells Fuel cells bring in the oxidizing and reducing agents as gases Graphite is typically the anode and cathode for the reaction which produced electricity. Fuel cells are clean and efficient.

25 Corrosion Corrosion is defined as the disintegration of metals.
Corrosion is typically caused by oxygen (O2). A familiar example of corrosion is iron rusting. Corrosion is a result of a redox reaction involving a metal. Iron oxide (rust)

26 Corrosion con’t…

27 Corrosion Prevention Typical corrosion protection involves plating the iron with another metal. The production of steel (iron and carbon) reduces the rate of corrosion of the iron. Aluminum, zinc, titanium are some metals which corrode slowly, or have different properties used to protect iron.

28 …Corrosion Prevention

29 Electrolytic Cells Electricity is used to force a chemical reaction
Electrolysis Used to obtain metals from molten salts Starting/keeping a car running Plating metals When starting a car- spont rx occurs in battery to provide electricity to start the car To keep the car running-alternator (non-spont rx) recharges the battery Also making aluminum Electrolysis of water into h and o gas

30 Electroplating Item to be plated is cathode
Metal that will plate is anode Put in solution containing ions- electrolyte

31 Electroplating con’t Benefits Drawbacks Resists corrosion
Improves appearance Cheaper Drawbacks Plating isn’t always even Can wear off Solutions are toxic

Download ppt "Unit 11- Redox and Electrochemistry"

Similar presentations

Ads by Google