Unit 11- Redox and Electrochemistry

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

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

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

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)???

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

Try these: • HNO3 • CO2 • K2PtCl6 • PCl5 • H2SO4

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)

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

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

Example: • 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. Cu + AgNO3 Cu(NO3)2 + Ag

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

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

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.

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.

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.

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

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

Reduction potentials for many electrodes has already been measured

Ecell  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.  = Ered (cathode) −Ered(anode) 

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

= (anode) (cathode) − Ered Ered   Ecell  Cell Potentials = +0.34 V − (−0.76 V) = +1.10 V

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.

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.

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.

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)

Corrosion con’t…

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.

…Corrosion Prevention

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

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

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

Unit 11- Redox and Electrochemistry