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Chapter 5

CHEM 4AA3 Chapter Notes - Chapter 5: Strong Electrolyte, Magnesium Chloride, Alkali Metal


Department
Chemistry
Course Code
CHEM 4AA3
Professor
Mescudi
Chapter
5

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Chemistry Chapter Five
The Nature of Aqueous Solutions
Reactions in aqueous (water)solution are important because: (1) water is
inexpensive and is able to dissolve a vast number of substances; (2) in such
solutions, many substances are dissociated into ions which can participate
chemical reactions; and (3) these solutions are found everywhere, from seawater to
living systems.
Unlike metallic conductors in which electrons carry the electric charge, the
electricity conducted in aqueous solutions is carried by the ions. When a solute
dissociates into ions in an aqueous solution and becomes and electric conductor, it
is known as an electrolyte. ***Pure water contains so few ions that it does not
conduct an electric current. *** Based on how well a solution conducts electricity,
we can deduce the strength of the presence of ions. We can label a solute as a non-
electrolyte, strong electrolyte, or weak electrolyte. A non-electrolyte is a
substance that is not ionized and does not conduct electric current (e.g. the lamp
fails to light up). Therefore, there are no ions or extremely low concentration of
ions. A strong electrolyte is a substance that is essentially completely ionized in
aqueous solution, and the solution is a good electrical conductor (e.g. the lamp
lights up brightly) and thus, has a high concentration of ions. A weak electrolyte is
partially ionized in aqueous solution and the solution is only a fair conductor of
electricity, thus, the concentration o f ions in the solution is low (e.g. the lamp lights
up only dimly). When determining if a solution is more likely to be a strong
electrolyte, weak electrolyte or non-electrolyte, it is best to remember this
generalization:
Essentially all soluble ionic compounds and only a relatively few molecular
compounds are strong electrolytes.
Most molecular compounds are either non-electrolytes or weak electrolytes.
Some examples of a strong electrolyte are: HCl, NaOH and KBr. Some examples of a
weak electrolyte are: HF, CH3COOH. Some examples of non-electrolytes are: H2O
and CH3OH.
If a solution contains strong electrolytes, the equation is written with the arrow of
the reaction going in one direction, usually right. This indicates that the ionization in
water is complete.
MgCl2(s) (H20) Mg 2+ (aq) + 2Cl-(aq)
In a situation where the solution is characterized as a weak electrolyte is best
described as a reaction that does not go to completion. In these cases, only a
portion of the solute molecules in the solution are ionized. The double arrows
indicate that the process is reversible. This means that while the forward reaction is
taking place, the reverse action is also occurring and its products are the reactants
of the forward reaction.
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HC2H3O2↔H+ (aq) + C2H3O2- (aq)
As for a non-electrolyte solution, we would imply write the molecular formula (e.g.
CH3OH (aq))
Relative Concentration in Solutions
In fully dissociated reactions such as the decomposition of magnesium chloride as
state above, the concentration is derived from the number of ions present in the
equation for each element based on one elements statistics. For example, as in the
decomposition of magnesium chloride, assume that there is 0.0050 M of MgCl2 to
decompose and we were to find the concentration of its products. We would write it
as the following: [Mg 2+] = 0.0050 M and [Cl-] = 0.0100 M. The square brackets
indicate concentration. The reason [Cl-] is twice the M of [Mg2+] is because there are
2 ions of Cl- for every Mg2+ ion.
Problem: What are the aluminum and sulfate ion concentrations in 0.0165 M
Al2(SO4)3(aq)?
Solution: Identify the solute as a strong electrolyte and write an equation to
represent its dissociation.
Al2(SO4)3(s) (H2O) 2 Al 3+(aq) + 3SO42-(aq)
[Al3+] = [(0.0165 mol Al2(SO4)3/ 1 L) x (2 mol Al3+/ 1 mol Al2(SO4)3)] = 0.0330 mol
Al3+/ 1 L = 0.0330 M
[SO42-] = [(0.0165 mol Al2(SO4)3/ 1 L) x (3 mol SO42-/ 1 mol Al2(SO4)3)] = 0.0495 mol
SO42-/ 1 L = 0.0495 M
Precipitation Reactions
Precipitation reactions occur when certain cations and anions combine to produce
an insoluble solid called precipitate.
Solubility Rules:
Compounds that ARE soluble: Alkali metal ions and ammonium salt (Li+, Na+,
K+, Rb+, Cs+ and NH4+); Nitrates (NO3-), perchlorate (ClO4-) and acetates
(CH3CO2-).
Compounds that ARE mostly soluble: Chlorides, bromides and iodides (Cl-,
Br-, I-)***Except those of Pb2+, Ag+ and Hg22+***; Sulfates (SO42-) ***Except
those of Pb2+, Sr2+, Ba2+ and Hg22***
Compounds that ARE insoluble: Hydroxides and sulfides (OH- and S2-)
***Except alkali metal and ammonium salts, sulfides of alkaline earths are
soluble, hydroxides of Sr2+ and Ca2+ are slightly soluble***; Carbonates and
phosphates (CO32- and PO43-) ***Except alkali metal and ammonium salts***;
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