Chem 30 - Unit 8 - Electrochemistry

Oxidation

The loss of electrons in a chemical reaction

Operational Definition of Oxidation

Any chemical reaction in which an atom or compound reacted with molecular oxygen

Reduction

The gain of electrons in a chemical equation

Operational Definition of Reduction + Example

When ores are reduced from oxides back to smaller, pure metals. Ex. Magnetite, Fe3O4(S), to Fe(3+)

Oxidation-Reduction Reactions/Redox Reactions

Reactions in which electrons are gained by one atom or ion and lost by another

What reactions are always Redox reactions?

Combustion and single-replacement reactions are always redox reactions

Oxidizing Agent

A reactant that oxidizes another reactant by being reduced. Ex. Metal ions

Reducing Agent

A reactant that reduces another reactant by being oxidized. Ex. Metal atoms

Spontaneous Redox Reactions

Reactions which proceed with no addition of energy or any other stimulus. Redox reactions are spontaneous when the stronger reducing agent is losing electrons and stronger oxidizing agent is gaining electrons

Net ionic equation

Shows only the ions involved in a chemical reaction

Half-reaction

Reactions which describe the changes in only the compound that is oxidized or reduced

Oxidation half-reaction

Zn(s) -> Zn+2(aq) + 2e-

Reduction half-reaction

Cu2+(aq) + 2e- -> Cu(s)

Balancing half-reactions in acidic or basic solutions

1. Write unbalanced half-reactions that show the formulas of the given reactant and product. 2. Balance any atoms other than oxygen and hydrogen first. 3. Balance any oxygen atoms by adding water molecules. 4. Balance any hydrogen atoms by adding hydrogen ions. 5. Adjust for basic conditions by adding the same number of Hydroxide ions as the number of hydrogen present on both sides of the equation 6. Combine hydrogen ions and hydroxide ions into water 7. Cancel out water molecules Balance the charges by adding electrons

Disproportionation Reaction

A reaction where some atoms of an element are oxidized and others are reduced

Smelting

Heating iron ore with charcoal to extract metallic iron

Oxidation Number

A number assigned to an atom in a compound if the electrons were completely held by the atom with the greatest electronegativity. Ex. H2O, oxygen oxidation number is -2, while hydrogen is +1. Will always add up to the charge of the compound or polyatomic ion

Electric cell

A device that continuously converts chemical energy into electrical energy or vice versa

Voltaic/Galvanic Cell

A cell which converts chemical energy into electric using spontaneous redox reactions

Electrode

Solid in half cell

Anode

Electrode which is oxidized, the negative end, electrons flow away, gets smaller

Cathode

Electrode which is reduced, the positive end, electrons flow towards, gets bigger

Electrolyte

Solution in each half cell, contains ions of the electrodes element

Salt Bridge

A solution with a porous boundary at each end that helps keep the two electrolytes neutral without mixing, preventing resistance to the movement of electrons. Gets attached to an external circuit

Shorthand notation

Uses a / to separate phases, // to indicate the porous cup or salt bridge, and , to separate chemical species in the same phase

Inert electrode

An electrode which does not participate in the reaction but provides a current. Used if the reaction doesn't have two solids. Pt or graphite (C) are most common

Electric potential difference/energy (V)

The difference per unit charge, or amount of energy released. In standard conditions, Eonet = Cathode

What is Ep relative too?

Hydrogen ions = 0.00 volts.

What happens to Ep under nonstandard conditions?

It slows until it dies, or reaches equilibrium at 0

Charge (Q measured in C)

The total charge transferred by a cell or battery by the movement of charged particles, measured in coulombs. Q=It

Electric Current (I measured in A or C/s)

Rate of flow of charge past a point in a circuit, measured in amperes by an ammeter. The larger the electric cell, the greater the current that can be produced

Power (W or J/s)

The rate at which a cell or battery produces electrical energy, measured in watts. The product of the current and the voltage of the battery

Energy density/Specific energy (J/kg)

The quantity of energy stored per unit mass

Corrosion

Reaction with metals and oxygen in the air, returning it to an ore-like state. Most metals form an oxidized layer around them and prevent further oxidization

Rusting

The corrosion of iron or steel, forming rust. Iron/steel is the anode, oxygen is the cathode, and things like rain act as electrolytes. Forms Fe2O3 x H2O or Fe(OH)3

Rust prevention

○ Cover iron with paint, tar, grease, or chromium
○ Attach a Sacrificial anode

Cathodic protection

Fueling iron with electrons to force it to become the cathode, using an impressed current or a sacrificial anode

Galvanizing

Applying a protective zinc layer around the steel or iron

Battery

A group of two or more electric cells connected to each other in series. Voltage is the sum of the voltages of the individual cells

Dry cell battery

Battery which uses manganese dioxide and zinc as the redox reactions. Non-rechargeable due to the OH- produced. Clocks, remote controls, toys, etc.

Nickel-Cadmium Battery

Rechargeable battery (secondary cells) which uses nickel oxide and basic cadmium as the redox reaction

Lead storage battery

Car battery which uses lead, lead coated in lead dioxide, and sulfuric acid. Six cells in series

Fuel cell

Cells where reactants are continuously supplied, and the energy is then used to run machines

Hydrogen-Oxygen fuel cells

Uses hydrogen gas at anode and oxygen gas at cathode. Need constant supply of hydrogen, from reformers which convert CH4 into H2 and CO2. Not very efficient, but more so than gas power. Produces greenhouse gases

Electrolytic cell

Cells which turn electrical energy into chemical. Non-spontaneous, negative. Uses Electrolysis. Cathode is negative and anode is positive

Electrolysis

Supplying electrical energy to cause a non-spontaneous redox reaction to occur, to separate elements from their ores

What are electrolytic cells used for?

To electroplate, recharge batteries, and split compounds into useful gases

Silver tarnish

Silver sulfide, formed from hydrogen sulfide in the air reacting with oxygen and the silver. 4Ag_(s) + 2H₂S_(aq) + O_2(g) -> 2Ag₂S_(s) + 2H₂O

What can be used to remove silver tarnish? What’s the disadvantage of this?

Silver polish, but it also removes some of the silver along with it, which can lead to degrading over time.

What are the similarities between electrolytic and voltaic cells?

Cathode is reduced, anode is oxidizes, electrons move from the anode to the cathode, anions move to anode, cations move to cathode

Secondary Cell

A rechargeable cell. When it discharges, it is acting voltaic, and when it recharges, it is acting electrolytic. Ex. Nickel-Cadmium

The Chloride Anomaly

When chloride and water are the only reducing agents, chlorine gas will be produced instead of oxygen due to a variation in activation energies (not needed to understand this super well, just memorize the rule).

Chlor-Alkali process

The process of producing chlorine, sodium hydroxide, and hydrogen gas from the electrolysis of aqueous sodium chloride. Utilizes the chloride anomaly.

Electrorefining

Using an electrolytic cell to obtain high-grade metals at the cathode from impure metal at the anode. Example photo:

Electrowinning

Reducing metal cations from a molten or aqueous electrolyte at the cathode of an electrolytic cell. Some metals cannot be reduced in water-containing solutions as water is a stronger OA, and so must use molten metals as electrolytes

How is aluminum produced? What did the discovery of this method do?

Al₂O_3(s) must be dissolved in inert solvent cryolite to allow for the electrowinning of aluminum. This discovery changed aluminum's price from $45 000 per kg to only 90 cents.

Electroplating

Plating a metal at the cathode of an electrolytic cell, usually to keep costs cheap while preventing corrosion and improving appearance. Science still cannot successfully predict what may work for electroplating and so much of it is trial and error

What is hydrogen gas used for?

To produce hydrogen peroxide, ammonia, margarine, and petroleum

What is chlorine gas used for?

Sanitizing water, and producing bleach, plastics, pesticides and solvents

What is sodium hydroxide used for?

To make cellophane, pulp and paper, aluminum, and detergents

Coulomb

The charge transferred by a current of one ampere during one second

Faraday’s Law

Mass is directly proportional to the time a cell operates given a constant current

Faraday’s Constant

9.65×10^4 C/mole-, the coulombs of charge transferred for every mole of electrons flowing in a cell

Moles of electrons

ne- = Q/F or It/F

Shorthand for electrolytic cells

Anode | Ions , ions | Cathode

Plating of iron with zinc (cathode, anode, half reactions)

Uses iron at the cathode, zinc at the anode, and the half reactions of both the oxidation and reduction of zinc.