Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109 - General Chemistry - Spring 2011

Lecture Notes 11: 11 February

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Chemical Reactions, cont.

Ionic reactions - dissolving and precipitates, cont.

Ba2+(aq) + 2 Cl-(aq) + 2 K+(aq) + SO42-  right arrow BaSO4(s) + 2 K+(aq) + 2 Cl-

Again, we want to write a net ionic equation showing only the ions which reacted:

Ba2+ + SO42-  right arrow BaSO4(s)

Notice that net ionic equations are very general expressions. Essentially they are saying that any time we have these species present they will react, regardless of what else happens to be there! (Sometimes folks are confused when they add ions which should react and they don't. This is usually a case where something else reacted first, so the ions of interest really weren't there!).

Since all ions in aqueous reactions are considered to be hydrated, we do not generally include (aq) as part of the formula. But remember, it is assumed!

Solubility Rules

It is useful to remember some simple "rules" (really more like guide lines) to help in predicting reactions. For common compounds such as we see in general chemistry we can use the following rules:

  1. Nitrates (NO3-) are all soluble.
  2. Alkali metal (Li+, Na+, K+, Cs+, and Rb+) and ammonium (NH4+) salts are all soluble, with the exception of a few Lithium salts.
  3. Chloride, bromide, and iodide (Cl-, Br-, and I-) salts are generally soluble, except for the salts of silver, lead(II) and mercury(I) (Ag+, Pb2+ and Hg22+).
  4. Sulfates are soluble, except for the salts of barium {BaSO4}, lead(II) {PbSO4}, mercury(II) {HgSO4}, and calcium {CaSO4}.
  5. Most hydroxides are only slightly soluble (but see rule 2).
  6. Sulfides (S2-), carbonates (CO32-), phosphates (PO43-), and chromates (CrO42-) are only slightly soluble (but see rule 2).

(Additional examples of solving net-ionic equations can be found in the Discussion Module.)

Oxidation/Reduction (Redox) Reactions

In these reactions we see a transfer of electrons from one atom or molecule to another. First let's look at some terms.

CH4 + 2 O2 right arrow CO2 + 2 H2O

Notice that the methane is oxidized by the oxygen. We say that the carbon and hydrogen are both oxidized to give the new covalent products, water and carbon dioxide, because the electrons are not evenly shared, they are pulled toward the oxygens in each bond.

Examples:

Ag+ + NO3- + Cu0 right arrow Ag0 + NO3- + Cu2+

Balancing: 2Ag+ + Cu0 right arrow 2Ag0 + Cu2+ 

H+ + Cl- + Ca0 right arrow H2 (g) + Ca2+ + Cl-

Balancing: 2 H+ + Ca0 right arrow H2 (g) + Ca2+

Cu2+ + SO42- + Fe0 right arrow Cu0 + Fe3+ + SO42-

Balancing: 3 Cu2+ + 2 Fe0 right arrow 3 Cu0 + 2 Fe3+

(Additional examples of solving net-ionic equations can be found in the Discussion Module.)

Balancing Redox Equations

There are two common methods for balancing redox reactions: the oxidation number method and the half-reaction method. The half-reaction method works very well for ionic reactions, it is relatively easy to give partial credit, and it is the only method I will use in this class. If you know how to do the other method you are welcome to do so, but be careful to make sure you show your work or I won't be able to give partial credit!

The Half-Reaction Method

In the half-reaction method what we do is first break an equation into two parts and then balance the parts individually. Presented stepwise:

Acid Solution:

Separate the reaction into two half-reactions.

Balance each half-reaction separately:

  1. Balance atoms other than O & H by inspection.
  2. Balance O by adding H2O to the opposite side.
  3. Balance H by addding H+ as appropriate.
  4. Balance the charge by adding electrons (e-) - add to same side as excess of positive charge, or opposite side if excess negative charge.
  5. Balance the charges of the two half-reactions by multiplying appropriately.

Add two equations together

Cancel items appearing on both sides.

Example. Balance the following equation as it occurs in acid solution:

MnO4- + Cl- right arrow Mn2+ + Cl2

First break the equation into two half reactions, one for Mn and one for Cl

MnO4- right arrow Mn2+

  1. MnO4- right arrow Mn2+
  2. MnO4- right arrow Mn2+ + 4 H2O
  3. 8 H+ + MnO4- right arrow Mn2+ + 4 H2O
  4. 5 e- + 8 H+ + MnO4- right arrow Mn2+ + 4 H2O
  5. 10 e- + 16 H+ + 2 MnO4- right arrow 2 Mn2+ + 8 H2O

Cl- right arrow Cl2

  1. 2 Cl- right arrow Cl2
  2. ...
  3. ...
  4. 2 Cl- right arrow Cl2 + 2 e-
  5. 10 Cl- right arrow 5 Cl2 + 10 e-
10 e- + 16 H+ + 2 MnO8-+ 10 Cl- right arrow 2 Mn2+ + 8 H2O + 5 Cl2 + 10 e-

16 H+ + 2 MnO4- + 10 Cl- right arrow 2 Mn2+ + 8 H2O + 5 Cl2

(Additional examples for balancing Redox equations in acidic solution can be found in the Discussion Module.

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© R A Paselk

Last modified 11 February 2011