Humboldt State University ® Department of Chemistry

Richard A. Paselk

Chem 109	General Chemistry	Summer 2002
Lecture Notes:: 13 June	© R. Paselk 2002
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Aqueous Solutions, cont.

• How much 1.000 M MgSO₄ is needed to make 500.0 mL of a 0.25 M solution.

First convert volumes to Liters: 500.0 mL = 0.5000 L

Can solve as a ratio: (0.5000 L) (0.25 M) = (1.0000 M) (x)

x = (0.25M)(0.5000 L)/(1.0000 M) = 0.125 L

Chemical Reactions

Ionic reactions - dissolving and precipitates: Much of the chemistry around us involves the dissolution of ionic solids in water to give aqueous solutions and the precipitation of ions from aqueous solution to give precipitates (solids). So what I would like to do first is to look a little at the process of dissolving and the nature of aqueous solutions.

We looked at water in some detail earlier (12 June notes). The thing we need to keep in mind is that the ions in water are not independent - they dissolve because they substitute interactions with water molecules for interactions with counter ions. And they stay in solution because they are insulated from each other by the water "shells" around each ion. A couple of corollaries

• First, only so many ions will dissolve until we run out of water molecules to make the shells, then they will interact with each other and precipitate out.
• Second, the amount of a given ion which will dissolve in water depends on the competition between water and its counter ion for interaction with it. If the ion-ion interaction is very strong, then they will precipitate out at very low concentrations etc.

Let's consider some chemical processes:

• Dissolve NaCl in water (overhead) NaCl Æ Na⁺_(aq) + Cl^-_(aq)
• Add NaCl solution to a AgNO₃ solution. First recall that each of the ions in solution is in an (aq)ueous complex, so we can write:

• Mix barium chloride and potassium sulfate:

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

1. Nitrates (NO₃^-) are generally soluble.
2. Alkali metal (Li⁺, Na⁺, K⁺, Cs⁺, and Rb⁺) and ammonium (NH₄⁺) salts are nearly all soluble.
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 (other than barium sulfate {BaSO₄}, lead(II) sulfate {PbSO₄}, mercury(II) sulfate {HgSO₄}, and calcium sulfate {CaSO₄}) are soluble.
5. Most hydroxides are only slightly soluble (but see rule 2).
6. Sulfides (S^2-), carbonates (CO₃^2-), phosphates (PO₄^3-), and chromates (CrO₄^2-) are only slightly soluble (but see rule 2).

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.

• Oxidation refers to taking electrons away from a substance. So to oxidize means to behave like oxygen normally does and "steal" electrons.
• Reduction refers to recieving electrons - it is the opposite of oxidation.
• Note that oxidation and reduction always go together. When oxygen oxidizes it is itself reduced (it gains electrons).

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Last modified 13 June 2002