Balanced Overall Chemical Equation

Balanced Overall Chemical Equation

Activity Series Lab


Although oxidation and reduction always occur simultaneously in redox reactions, the overall chemical equation can be separated into the oxidation part and the reduction part. Each part is called a half-reaction. Furthermore, overall redox reactions are often expressed as net ionic equations, omitting spectator ions. The following examples illustrate the treatment of chemical reactions in electrochemistry.

Balanced Overall Chemical Equation

2 AgNO3 + Cu(s)  Cu(NO3)2 + 2 Ag(s)

Balanced Net Ionic equation

2 Ag+ + Cu(s)  Cu2+ + 2 Ag(s)

Balanced Half Reactions

Oxidation Half-ReactionCu(s)  Cu2+ + 2 e-

Reduction Half-Reaction2 Ag+ + 2e-  2 Ag(s)

To see how half-reactions relate to the balanced overall net ionic equation, the electrons can be cancelled from both sides of the equation and the two half-reactions added.

Cu(s)  Cu2+ + 2 e-

+ 2 Ag+ + 2e-  2 Ag(s)

2 Ag+ + Cu(s)  Cu2+ + 2 Ag(s)

Given that there are many possible combinations of chemical reactions that can occur between metals and their solutions of dissolved ions, chemists have come up with a list that ranks metals from the most reactive to the least reactive, called the activity series of the metals. The ability to generate a part of this list, and interpret its meaning is fundamental to a further understanding of redox chemistry.


Develop an activity series given 4 metals and their solutions of dissolved ions.


  • Solid zinc, mossy
  • Solid tin, mossy
  • Solid lead
  • Solid copper

  • 0.1M Zn(NO3)2
  • 0.1M Sn(NO3)2
  • 0.1M Pb(NO3)2
  • 0.1M Cu(NO3)2


  • Test tube rack
  • 13 x 100 Test tubes (6)
  • 50 mL beaker (3)
  • Dropper pipet (3)


  1. Obtain all equipment.
  2. Pour 10mL of copper (II) nitrate solution into a clean, dry 50mL beaker.
  3. Small volumes (approx. 5 mL) of other solutions of ions will be required later, after you have determined which reactions are necessary to investigate.
  4. Obtain the smallest possible piece of each metal. More pieces will be required later, after you have determined which reactions are necessary to perform.
  5. Add even volumes of copper (II) nitrate solution to 3 test tubes with a dropper pipet.
  6. Add a piece of metal to each solution. Observe.
  7. Record your results in the table below.
  8. Perform as many reactions as necessary to complete the table below.

Qualitative Observations

Note that reactions will only occur in one direction. For example, if solid lead reacts with a solution of copper (II) ions, the opposite reaction will NOT occur. Otherwise stated, if solid copper is combined with a solution of zinc ions, no reaction will occur. Therefore, ideally, a minimum of 6 combinations should yield a complete table of results.

Results of Single Replacement Reactions

Red. Agents\Ox. Agents

/ Cu2+ / Zn2+ / Sn2+ / Pb2+
Cu(s) / Rx.No Rx. / Rx.No Rx. / Rx.No Rx.
Zn(s) / Rx.No Rx. / Rx.No Rx. / Rx.No Rx.
Sn(s) / Rx.No Rx. / Rx.No Rx. / Rx.No Rx.
Pb(s) / Rx.No Rx. / Rx.No Rx. / Rx.No Rx.



Student 1 Name:


Student 2 Name:

According to the results of the experiment, the reactivity series of the four metals, from most reactive metal to least reactive metal, is as follows.

Most Reactive / Cu2+Zn2+Sn2+Pb2+Cu(s)Zn(s)Sn(s)Pb(s) / Cu2+Zn2+Sn2+Pb2+Cu(s)Zn(s)Sn(s)Pb(s) / Cu2+Zn2+Sn2+Pb2+Cu(s)Zn(s)Sn(s)Pb(s) / Cu2+Zn2+Sn2+Pb2+Cu(s)Zn(s)Sn(s)Pb(s) / Least Reactive

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