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Question	Answer
Historical Definitions	Reduction: The loss of mass when a metal is refined from its ore Oxidation: reacting a substance with oxygen (often a metal)
Modern Definitions	Reduction: The gain or partial gain of electrons Oxidation: The loss or partial loss of electrons
Half-Reactions	Show how stable substances gain or lose electrons to form their respective charged ions. Can be a reduction half-reaction or an oxidation half-reaction
Net Ionic Equations	Two half-reactions combined to form a singular equation. Will always be a reduction half-reaction mixed with an oxidation half-reaction
Electronegativity	Can play an important role in the ways that electrons are transferred or shared
Redox Reactions	Must have a reducing agent and an oxidizing agent One substance is reduced and one substance is oxidized The net reduction must equal the net oxidation
Oxidizing Agents	Substances that cause oxidation The substance that is being reduced is the OA because it causes the other to be oxidized
Reducing Agents	Substances that cause reduction The substance that is being oxidized is the RA because it causes the other to be reduced
Redox Table	Shows the half-reactions of substances and combinations of substances Shows the relative strengths of OAs and RAs
Spontaneous Reaction	When a reaction occurs naturally and products are formed. The strongest oxidizing agent is on the left side of the equation
Non-Spontaneous Reaction	When a reaction does not occurs naturally and products are not formed. The strongest oxidizing agent is on the right side of the equation
Oxidation Numbers	Numbers assigned to all types of substances as if they were an ionic charge. The rules are: Elements are 0 Oxygen is (-2) unless it is peroxide or w/ F Hydrogen is (+1) unless it is a hydride The oxidation # for an ion is the same as its charge The oxidation number for a compound is either 0 or its charge
Types of Redox Questions	Constructing a Redox Table Predicting a Redox Reaction Balancing Redox Equations (Oxidation numbers or Half-Reactions) Titration Questions/Redox Stoichiometry
Electrochemical Cell	Contains solid electrodes dipping into an electrolyte solution, is composed of two half-reactions (one at each electrode), and allows for a load to use energy (using the flow of electrons)
Electrochemical Cells Must Include	Oxidation Half-Reaction Reduction Half-Reaction Net Ionic Equation Voltage of the Cell. Line Notation Labelled Anode Labelled Cathode Labelled Electrolyte Direction of Electron Flow Ion Movement
Anode	A solid electrode that serves as the location of oxidation. Anions migrate toward the anode. Electrons are produced.
Cathode	A solid electrode that serves as the location of reduction. Cations migrate toward the cathode. Electrons are consumed.
Types of Electrochemical Cells	Voltaic and Electrolytic
Voltaic Cells	Are always a spontaneous reaction, even if it is normally non-spontaneous. Always goes from the anode to the cathode. Takes place in two separate containers that are connected with a salt-bridge or porous boundary. Converts chemical energy to electrical energy. Has a positive cell potential.
Electrolytic Cells	Always a non-spontaneous reaction, even if it is normally spontaneous. Always goes from the anode to the cathode. Takes place in one container. Converts electrical energy to chemical energy. Has a negative cell potential.
Dry Cells	A carbon cell with a paste electrolyte and carbon rod. A normal battery would be a dry cell.
Alkaline Cells	Batteries that improve on long term efficiency compared to a dry cell
Button Cells	Small circular cells that have the cathode on the bottom that is either mercury or silver oxide.
Lead Acid Battery	Multiple cells set up in a series with a lead plate separated with porous material
Nicads	Nickel cadmium cells that are not efficient because they develop a memory as the reverse reaction becomes harder to do.
Fuel Cells	Highly efficient fossil fuel cells that make a continuous energy input and output
Primary Cells	Non-reversible reactions which are used and the discarded.
Secondary Cells	Reversible reactions that store energy and are able to recharge after energy is depleted.
Net Voltage	The sum of reduction and oxidation potential Listed in the data book by the half-reactions The sign is reversed when the reaction is reversed
Cell Potential	The energy needed for electrons to move to the cathode Listed in the data book by the half-reactions Do not change the signs from original The cathode — the anode
S.H.E.	Standard Hydrogen Electrode is the reference point for the net voltage of other half-reactions in the redox table. If the standard is changed, the difference in voltage potential will also stay the same.
Line Notation	A way of abbreviating the components of a cell. Single line indicates change is state Reactants come before the products Double line indicates a bridge or separation Anode is written before cathode (oxidation before reduction)
Electrolysis of Water	Shows how the reduction half -reaction for water and oxidation half-reaction for water join together to create 2H2O=2H2+O2 Remember that Hydrogen ions and hydroxide ions add together to form water
Chloride Anomoly	When the strongest reducing agent is between water and chloride, the chloride half-reaction is used.
Uses for Electrolysis	To protect precious metals from corrosion To produce metals from salts To remove precious metals from other substances
Down's Cell	When there is electrolysis of a molten salt and a molten metal is gathered as the product. Has a different set-up and usually has a gas outlet for the other product.
Corrosion Prevention	1. Prevent Interaction 2. Paint or Coat metal 3. Galvanize metal
Sacrificial Anodes	Is also called cathodic protection and uses a stronger reducing agent so a different substance gets oxidized instead.
Cell Stoichiometry	Uses the formula: Ne = It/F Faraday's Constant is 96500 C/mol of e-

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	Created by Natasha Gidluck over 6 years ago