Download Gen Chem Redox Reactions Summary

Gen Chem
Unit: Redox Reactions
Name: ________________________
Redox Summary
Corresponding Book Chapters: p. 644-674 in AW text and Sections 19.1-19.2 in the Glencoe packet.
Learning Objectives:
3.1: Identify redox reactions beyond the classroom
3.2: Define oxidation and reduction in terms of electron loss or gain
3.3: Manipulate oxidation number rules to determine unknown oxidation numbers
3.4: Identify oxidation number changes in chemical reactions to predict oxidation and reduction
processes
3.5: Balance simple half reactions in redox reactions
Question: What is a redox reaction?
Answer:
Oxidation is when electrons are __________________ and reduction is when electrons are _________________.
When writing redox equations, remember that oxidation looks like this:
Na  Na+ + e…and reduction looks like this:
Cl2 + 2e-  2ClWith oxidation, lost electrons are written on the PRODUCT side of the equation. With reduction, gained
electrons are written on the REACTANT side of the equation.
Review and/or copy Table 19.1 in the Glencoe packet (p. 683). It is a good summary of the above.
Question: What is a spectator ion?
Question: What do the following mean and why are they necessary to include in redox equations?
A. (aq)
B. (l)
C. (s)
D. (g)
E. 
As this relates to electronegativity (EN): Elements with low ENs (groups 1 & 2) are strong reducing
agents, and those with high ENs (groups 16& 17) are strong oxidizing agents. Check p. 684, Figure 19.4.
*Know the rules for assigning oxidation numbers. You may or may not be given them on the exam.
Example: For the following equation, identify the element that is oxidized, the element that is reduced,
and the spectator ion.
2KBr + Cl2  2KCl + Br2
OX:
RED:
SPECTATOR:
Balancing Redox Equations
Question: When atoms LOSE electrons, what happens to the oxidation number? _____________________
When atoms GAIN electrons, what happens to the oxidation number? _____________________
**You must make sure that the total # of electrons lost equals the total # of electrons gained.
Work the practice problems on p. 690 in the Glencoe packet, #15-17.
Half Reactions
Recall that in the oxidation portion of a redox equation, electrons appear as PRODUCTS because they are
LOST.
Sn+2  Sn+4 + 2e-
In the reduction portion of a redox equation, electrons appear as REACTANTS because they are GAINED.
Fe+3 + e-  Fe+2
Rules for balancing redox equations in an acid solution:
--For the HALF REACTION
1.
2.
3.
4.
Balance all atoms other than H and O
Balance O by adding water (H2O) to the other side of the equation
Balance H by adding the hydrogen ion (H+) to the other side of the equation
Balance the net charge by adding electrons to the other side of the reaction
--THEN
5. Make sure the number of electrons gained equals the number of electrons lost. In other words,
the net charge should be equal to zero. You may have to use coefficients to do this.
6. Add the half reactions and then cancel anything that is the exact same on both sides of the
equation. The electrons SHOULD cancel out if step 5 is done correctly. Simplify the coefficients if
you need to.
Practice using the steps on the following problem:
Fe(s) + CuSO4(aq)  Cu(s) + Fe2(SO4)3(aq)
Electrochemistry
Electrochemistry translates the chemical energy of a reduction–oxidation reaction into electrical energy.
Points to remember:
1. When the substances involved in oxidation and reduction half–reactions are physically
separated, it is called an electrochemical cell.
2. Each half reaction occurs on the surface of an electrically conductive solid called an electrode.
3. Each electrode is immersed in a solution containing ions needed for the half–reaction.
4. The electrodes are connected by a wire so that electrons can move from the oxidation half–
reaction to the reduction half–reaction.
5. The solutions are connected by a salt bridge so that ions can move between solutions.
6. In an electrochemical cell the chemical potential energy can be harnessed as the substances
undergoing oxidation push electrons through the wire to the substances undergoing reduction.
Diagram of an electrochemical cell:
Source: http://staff.prairiesouth.ca/~chemistry/chem30/6_redox/redox2_2.htm
There are two main parts to an electrochemical (or voltaic) cell.
Anode:
The anode is the electrode in the electrochemical cell where oxidation occurs. Anions are attracted to the
anode because the anode is positively charged since it (the anode) is losing electrons in the half reaction.
Cathode:
The cathode is the electrode in the electrochemical cell where reduction occurs. Cations are attracted to
the cathode because the cathode is negatively charged since it (the cathode) is gaining electrons in the
half reaction.
Salt Bridge: (the upside-down U in the middle)
The salt bridge moves ions into the solutions to maintain electric neutrality without mixing them.
Electrochemical Cells:
A characteristic of electrochemical cells is that the redox reaction may occur spontaneously (voltaic cell),
or non-spontaneous reactions can be forced to occur (electrolytic cell & electroplating).
1.
Voltaic cells are typically used to produce electrical energy. In fact, batteries are a voltaic cell.
2. Electrolytic cells use electricity to bring about a redox reaction that would normally be nonspontaneous. In other words, low-energy reactants become high-energy products.
3.
Electroplating is a process that coats a material with a layer of metal based on metallic activity.
4. Rechargeable batteries: Spontaneous redox reactions eventually deplete the electrons available at
the anode (e.g. causing a battery to become “dead”). The redox reactions can be “reversed” by a nonspontaneous reaction.
a. Voltaic cells are “batteries” that supply potential (voltage or electrical energy) to run motors, light
bulbs, etc.
b. Eventually, the potential in the voltaic cells will deplete because the electrons from the anode
(oxidation) are depleted and built up at the cathode (reduction).
c. A batteries or outside voltage supply can be connected to the voltaic cell so that the anode is
connected to the negative terminal (supplying electrons that were lost) and the cathode is connected to
the positive terminal (removing the electrons that were gained).
Voltaic cells
a. In a voltaic cell, electrons are spontaneously emerging at the cathode so that reduction can occur.
i. The electrons are “pumped” from the anode towards the cathode, supplying the resistance (bulb,
motor, cell, etc.).
ii. Since the electrons leave the anode, it is designated with a positive charge. The electrons are built
up at the cathode; therefore, the cathode is often denoted with a positive sign.
iii. Strictly speaking, the term “cathode” was derived because “cations” (positively charged
ions/particles) were attracted to it. The “anode” attracts the “anions” (negatively charged ions/particles).
iv. The salt bridge completes the circuit so that current (electrons) can flow. If the circuit is
“broken”, the half reactions will not occur.
b. The electrons are the source of potential in the voltaic cell.