The Copper Cycle
Most of the background material for this laboratory will be covered in greater detail in the lecture course later in the semester. Here is some background information so you will understand the chemistry behind the reactions you will perform. Many aspects of our lives involve chemical reactions—from the batteries that power our cars and cell phones to the thousands of processes occurring within our bodies. Most of these reactions can be classified into one of three main types of chemical reactions: precipitation reactions, acid-base neutralization reactions, and oxidation-reduction (also called “redox”) reactions.
Many reactions occur in an aqueous environment (i.e., in a solution where ions and compounds are dissolved in water). When we indicate that a reactant or product has the physical state (aq), we mean the substance is dissolved in water. When an ionic compound is in aqueous solution, the individual ions are present in solution; for example, NaCl(aq) exists as Na+ and Cl– ions moving around in water.
Many ionic compounds are soluble—i.e., they dissolve in water. Others generally do not dissolve in water and are considered insoluble. To determine if an ionic compound is soluble—i.e., will dissolve—in water, we use the Solubility Rules:
Solubility Rules for Ionic Compounds in Water
The compound is SOLUBLE if it has:
2. C2H3O2–, NO3–, ClO4–
3. Cl–, Br–, or I–, except compounds
with Ag+, Pb+2, and Hg2+2 are
4. SO42- except compounds with
Ag2SO4, CaSO4, SrSO4, BaSO4,
PbSO4, and Hg2SO4 are insoluble
The compound is INSOLUBLE if it has:
5. CO32–, CrO42–, PO43–, except compounds
with Li+, Na+, K+, NH4+ are soluble
6. S2–, except compounds with Li+, Na+, K+,
NH4+, Ca+2, Sr+2, Ba+2 are soluble
7. Hydroxide ion, OH–, except compounds
with Li+, Na+, K+, NH4+ are soluble
The Solubility Rules indicate which compounds are soluble, and thus are represented as aqueous: e.g., KI(aq), BaCl2(aq), NaOH(aq), etc. The Solubility Rules also indicate which compounds are insoluble—i.e., do not dissolve in water and remain as solids: e.g. BaSO4(s), AgCl(s), CaCO3(s), etc.
Double Replacement/Precipitation Reaction
For example, consider the reaction between aqueous lead(II) nitrate with aqueous potassium bromide, as shown below:
Note that the chemical formulas for the products formed are based on their charges, not how they appear on the reactant side of the chemical equation.
GCC CHM 151LL: The Copper Cycle
© GCC, 2013
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Based on Solubility Rules #4 and #1, we find that PbBr2 is insoluble and KNO3 is soluble. Thus, the complete, balanced equation is:
Pb(NO3)2(aq) + 2 KBr(aq)
We can cancel the spectator ions from the ionic equation and write the net ionic equation: Pb2+(aq) + 2 Br -(aq) PbBr2(s)
This reaction produces a cloudy mixture with small particles of the solid suspended in the solution. When enough solid has formed, it will begin to settle at the bottom of the beaker. Thus, a clear solution becoming cloudy when another solution is added is often taken as experimental evidence of a solid or precipitate forming.
Acids and Bases
Acids can be defined as substances that produce hydronium ions (H3O+) when they are dissolved in water. A hydronium ion is the product of a hydrogen ion that reacts with a water molecule: H+(aq) + H2O(l) H3O+(aq). A hydrated hydrogen ion (H+(aq)) is equivalent to an aqueous hydronium ion. The two equations below both represent the ionization of hydrochloric acid, HCl(aq), but the second one shows a particular water molecule explicitly.
HCl(aq) H+(aq) + Cl–(aq)
HCl(aq) + H2O(l) H3O+(aq) + Cl–(aq)
Acids are usually easy to recognize since their formulas start with H and contains nonmetal elements other than H—e.g. HCl(aq),...
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