Difference between revisions of "Chemistry Lab/Electrochemistry"

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(a bit more info here. i have to upload pics at some point (unless someone beats me to it))
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==Balancing Oxidation/Reduction Reactions==
 
==Balancing Oxidation/Reduction Reactions==
 
Redox reactions follow a simple set of steps to solve.
 
Redox reactions follow a simple set of steps to solve.
# Split it into 2 half-reactions, one being oxidation and the other being reduction. The reactant being oxidized loses electrons, so they start out without a charge and end up with a positive charge and an electron. The reactant being reduced gains electrons, so they start out charged and with an electron and end up uncharged. Keep in mind that the charge must be neutral on both sides; i.e., should the reducing agent lose 2 electrons, it must end up with a charge of +2 and with 2 extra electrons. Example: Cu (s) ----> Cu2+ + 2 e- An example of a reduction reaction would be 2 Ag+ (aq) + 2 e- ------> 2 Ag (s)
+
# '''Split it into 2 half-reactions'''. Usually, the basis of a redox reaction will be given with two compounds reacting to form two more compounds. Aside from oxygen and hydrogen, each reactant will have a corresponding element with a product. These two compounds will form one half-reaction; the other two form the other. Balance the coefficients of key compounds if necessary, i.e. Balancing Cr with Cr2O7. '''''Don't worry about balancing extra hydrogens or oxygens yet.'''''
# If it's in an acidic medium...
+
#'''Balance all non-hydrogen or oxygen elements'''. You may have balanced the key components of the half-reaction, but sometimes you'll have a pesky oxygen or hydrogen messing things up. To get this to balance, you must add ions to the other side to balance the reaction out. This process differs for acidic vs. basic solutions. Do '''NOT''' try to balance one equation with the other, this step '''ONLY''' comes at the end.
## Balance all non-hydrogen or oxygen elements
+
#If it's in an '''acidic medium''':
## Balance oxygen by adding H2O to the appropriate side
+
##'''Balance oxygen''' by adding H2O to the appropriate side. Add a coefficient to this H2O if necessary.
## Balance hydrogen by adding H+ to the appropriate side
+
##'''Balance hydrogen''' by adding H+ to the appropriate side. This can be on either side, depending on how many H2O's you added. Add a coefficient to this if necessary.
## Balance the charge by adding e-
+
##'''Balance the charge''' by adding electrons. Now you've balanced the equation in terms of elements, but the charge may not be balanced yet. To balance this, add electrons to the side with a higher charge until the total charge of each half of the half-reaction is the same.
# Multiply each reaction by an integer so that there are the same number of electrons on each side (i.e. they cancel out)
+
#If it's a '''basic medium'''
# Combine the half-reactions and cancel
+
##'''Balance oxygen''' as above.
# If it's in a basic medium, add OH- to each side until all H+ is gone; then, cancel again
+
##'''Balance hydrogen'''. Instead of balancing this with H+, you need to balance it with OH-. This means that you may get extra oxygens. If this happens, add another H2O on the other side and continue adding OH- until it balances. There is also '''another method''' that is detailed below.
 +
##'''Balance the charge''' as above.
 +
#Now, we can add the reactions together to come up with our final reaction. '''Multiply each reaction by an integer''' so that there are the same number of electrons on each side (i.e. they cancel out). This means that the electrons of one half-reaction should be on the OPPOSITE side of the electrons in the other half-reaction. If this is not the case, go back and check your work. More than likely, there's a mistake in there somewhere.
 +
#Combine the half-reactions and cancel. The elctrons should cancel out completely, and H2O's and H+'s may cancel somewhat. If you have time, it's usually a good idea to make sure that the equation is balanced by elements and by charge.
 +
 
 +
*'''Second method for balancing redox reactions in basic solution''': If it's in a basic medium, add OH- to each side of the final equation until all H+ is gone; then, cancel again. Remember that OH- + H+ ---> H2O in this step.
  
 
==Activity Series==
 
==Activity Series==

Revision as of 21:43, 13 April 2012

This page refers to the 2011 and 2012 foci of Chem Lab.

Basic Information

A redox reaction, or an oxidation/reduction reaction, occurs when one reactant is oxidized, or loses electrons, and one reactant is reduced, or gains electrons. A simple way to tell the difference is OIL RIG (Oxidation Is Losing; Reducing Is Gaining) or LEO says GER (Lose Electrons - Oxidize; Gain Electrons - Reduce). The oxidizing agent is reduced, and the reducing agent is oxidized. A half-reaction is exactly what it sounds like - half a reaction. It focuses exclusively on one portion of the reaction, either oxidation or reduction.

Balancing Oxidation/Reduction Reactions

Redox reactions follow a simple set of steps to solve.

  1. Split it into 2 half-reactions. Usually, the basis of a redox reaction will be given with two compounds reacting to form two more compounds. Aside from oxygen and hydrogen, each reactant will have a corresponding element with a product. These two compounds will form one half-reaction; the other two form the other. Balance the coefficients of key compounds if necessary, i.e. Balancing Cr with Cr2O7. Don't worry about balancing extra hydrogens or oxygens yet.
  2. Balance all non-hydrogen or oxygen elements. You may have balanced the key components of the half-reaction, but sometimes you'll have a pesky oxygen or hydrogen messing things up. To get this to balance, you must add ions to the other side to balance the reaction out. This process differs for acidic vs. basic solutions. Do NOT try to balance one equation with the other, this step ONLY comes at the end.
  3. If it's in an acidic medium:
    1. Balance oxygen by adding H2O to the appropriate side. Add a coefficient to this H2O if necessary.
    2. Balance hydrogen by adding H+ to the appropriate side. This can be on either side, depending on how many H2O's you added. Add a coefficient to this if necessary.
    3. Balance the charge by adding electrons. Now you've balanced the equation in terms of elements, but the charge may not be balanced yet. To balance this, add electrons to the side with a higher charge until the total charge of each half of the half-reaction is the same.
  4. If it's a basic medium
    1. Balance oxygen as above.
    2. Balance hydrogen. Instead of balancing this with H+, you need to balance it with OH-. This means that you may get extra oxygens. If this happens, add another H2O on the other side and continue adding OH- until it balances. There is also another method that is detailed below.
    3. Balance the charge as above.
  5. Now, we can add the reactions together to come up with our final reaction. Multiply each reaction by an integer so that there are the same number of electrons on each side (i.e. they cancel out). This means that the electrons of one half-reaction should be on the OPPOSITE side of the electrons in the other half-reaction. If this is not the case, go back and check your work. More than likely, there's a mistake in there somewhere.
  6. Combine the half-reactions and cancel. The elctrons should cancel out completely, and H2O's and H+'s may cancel somewhat. If you have time, it's usually a good idea to make sure that the equation is balanced by elements and by charge.
  • Second method for balancing redox reactions in basic solution: If it's in a basic medium, add OH- to each side of the final equation until all H+ is gone; then, cancel again. Remember that OH- + H+ ---> H2O in this step.

Activity Series

One task participants may be asked to complete in this event is to construct an activity series based on what ions react with others.

Electrochemical Cells

Eelectrochemical cells results in an exchange of electrons in a redox reaction. There are two main types of electrochemical cells.

Voltaic Cells

A voltaic, or galvanic, cell is composed of two metals connected by a salt bridge. It uses the electron exchange to generate the current. It consists of two half-cells, each of which contains a metal solution with that metal submerged in it.

The cell in which oxidation occurs is called the anode, and the cell in which reduction occurs is called the cathode. You can remember this by knowing that reduction has a "c" in it, and cathode starts with a "c". Electrons flow from the anode to the cathode.

Electrolytic Cells

Electrolytic cells use a current to decompose chemical compounds.

Electron Potential

Links