Download Chapter 18 Review 18.1 Oxidation-Reduction Reactions Oxidation

Chapter 18 Review
18.1 Oxidation-Reduction Reactions
Oxidation-Reduction Reaction - where one or more electrons are transferred
Oxidation- loss of electrons
Reduction- gain of electrons
- If an atom gains an electron, the charge becomes negative
- If an atom loses an electron, the charge becomes positive
18.2 Oxidation States
Oxidation States- concept that allows to keep track of the elections in a redox reaction
Electronegativity- how much an atom wants electrons
- each element has one or more oxidation states
- uncombined elements have the charge of 0
- a compound has the charge of 0
- the two elements in a compound must cancel each other out to have a zero charge
- the most electronegative elements are F, O, N, and Cl
18.3 Oxidation-Reduction Reactions Between Nonmetals
Oxidation- increase in oxidation state
Reduction- decrease in oxidation state
Oxidizing Agent- causes the other compound to oxidize
Reducing Agent- causes the other compound to reduce
- transfers of electrons sometimes causes ions in the products
(Na → Na+, Cl→ Cl-)
- Combustion is an example of a Redox reaction
- the number of electrons need to match the charge and number of each atom in the reaction
- When an agent is names, the whole compound is specified, no an single element
18.4 Balancing Oxidation-Reduction Reactions be the Half-Reaction Method
Half-Reactions- equations that have electrons as reactants or products
- many redox reactions are balanced by trial and error
- make sure that there is the same number of each atom on both sides of the reaction
- to balance reactions that occur in aqueous solution, use the half- reactions
- One half-reaction represents the oxidation process, while the other represents the reduction
process
- In a reduction half-reaction, the electrons are on the reactant side
- In a oxidation half-reaction, the electrons are on the product side
- the number of electrons present should match the charge of the element
- balance the half-reaction to make sure there isn’t any lost of compounds
- the number of electrons lost must equal the number of electrons gained
- to balance the half-reactions by multiplying it by a coefficient to get the same number of
electrons in both reactions
- Add the half reactions together and cancel the electrons and any other catalysts
- most redox reactions occur in distinctly acidic or basic solutions
18.5 Electrochemistry: An Introduction
Electrochemistry- a study of the interchange of chemical and electrical energy
Galvanic Cell (aka electrochemical battery)- a device powered by an redox reaction
where the agents are separated, the electrons flow through a wire, and there is a salt bridge
connecting the two solutions
Anode- the electrode where oxidation occurs
Cathode- the electrode where reduction occurs
Electrolysis- electrical energy is used to produce a chemical change
- batteries uses energy from an redox reaction to produce an electric current
- Electrochemistry involves two types or processes:
1. the production of an electric current from a redox reaction
2. the use of an electric current to produce chemical change
- to use energy from the reaction, the oxidizing agent needs to be separated from the reducing
agent
- this causes the electrons to transfer through a wire
- the solutions also need to be connected to have electron flow (so the ions can flow)
- this keeps the net charge in each compartment zero
- the solutions are connected through a salt bridge
- the current produced in the wire by electron flow can be directed by a electric meter to do work
- this type of method is used to power many different machinery
- by forcing a current through water, oxygen and hydrogen gas can be produced
18.6 Batteries
Potential- energy due to position or composition
Lead Storage Battery
- lead storage battery has been a major factor in making automobile practical for transportation
- can function for several years under -30 F to 100 F
- the reducing agent is Pb (Lead)
- oxidizing agent is PbO (lead (II) oxide)
- the anode reaction is Pb + HSO → PbSO + 2 H+ + 2 e- the cathode reaction is PbO + HSO + 2 e- + 2 H+ → PbSO + 2 HO
- the overall reaction that happens is Pb + PbO + 2 HSO → 2 PbSO + HO
- the anode and cathode are separated and bathed in sulfuric acid
- the battery consists of metal grids with these solution on them
- depending on the ability of the reducing agent to accept electrons and the ability of the
oxidizing agent to give electrons, the amount of energy produced will differ
Dry Cell Battery
- don’t contain any liquid electrolytes
- In the acid version
-the anode reaction is Zn → Zn2+ + 2 e- the cathode reactions is 2 NH4+ + 2 MnO2 + 2 e- → Mn2O3 + 2 NH3 + H2O
- produces the potential of 1.5 volts
- in the alkaline version
- the anode reaction is Zn + 2 OH- → ZnO + H2O + 2 e- the cathode reaction is 2 MnO2 + H2O + 2 e- → Mn2O3 + 2 OH- this cell lasts longer
- other types are the silver cell and mercury cell
- there is the nickel-cadmium cell
- the anode reaction is Cd + 2 OH- → Cd(OH)2 + 2 e- the cathode reaction is NiO2 + 2 H2O + 2 e- → Ni(OH)2 + 2 OH- can be recharged indefinite amount of times
18.7 Corrosion
Corrosion- process of returning metals to their natural state
Cathodic Protection- the connection of an active metal to another to prevent corrosion
- corrosion involves the oxidation of metals
- this process creates great economic lose
- most metals produce a thin oxide coating, which protect their inner atoms from oxidizing
- a way to protect metals is to apply a coating of paint or metal plating
- Alloying is also used to prevent corrosion
- Cathode protection is used to protect steel in buried fuel tanks and pipelines
- furnishes electrons easily than iron
- the Mg used would need to be replaced periodically because it will dissolve after being
oxidized over a period of time
18.8 Electrolysis
Electrolysis- forcing a current through a cell to produce a chemical change that would not
otherwise occur
- a battery will normally run out of energy, but it can be recharged by using the method of
electrolysis
- this reverses the redox reaction happening in the battery
- electrolysis is also used in the production of metals from their ores
- when the redox reaction is reversed, it’s called an electrolytic cell
- in an electrolytic cell, the anode becomes the reduction side and the cathode becomes the
oxidation side
- the electron flow is reversed, putting everything in its original place