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Chapter 5

CHM238Y1 Chapter Notes - Chapter 5: Sulfuric Acid, Formaldehyde, Equivalence Point


Department
Chemistry
Course Code
CHM238Y1
Professor
F
Chapter
5

Page:
of 4
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.
HC2H3O2H+ (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***;
Silver, lead and mercury (Ag+, Pb2+, Hg22+) ***Except nitrates, acetates and
perchlorates***
Problem: AgNO3(aq) + NaI(aq) AgI(s) + NaNO3(aq). Write this a net ionic equation.
Solution:
Complete ionic equation:
Ag+(aq) + NO3-(aq) + Na+ + I-(aq) AgI(s) + Na+(aq) + NO3-(aq)
Net ionic equation:
Ag+(aq) + I-(aq) AgI(s)
Problem: Write the net ionic equation for the reaction that occurs when aqueous
solutions of magnesium sulphate (MgSO4) and sodium oxalate (Na2C2O4) are mixed.
Solution:
(PUT A LINE HERE)
Writing Net Ionic Equations
1. Write the balanced molecular equation.
2. Write the ionic equation showing the strong electrolytes completely dissociated
into cations and anions.
3. Cancel the spectator ions on both sides of the ionic equation.
4. Check that charges and number of atoms are balanced in the net ionic
equation.
Acid Base Reactions
Acids are identified by their sour tastes, their ability to react with a variety of metals
and carbonate minerals. From a chemist’s standpoint, an acid can be defined as a
substance that provides hydrogen ions (H+) in aqueous solutions. A strong acid
completely ionizes in a water solution:
HCl (aq) (H20) H+ (aq) + Cl- (aq)
A weak acid, on the other hand, does not go towards completion and is also a weak
electrolyte. A example is acetic acid: HC2H3O2 (aq) H+ (aq) + C2H3O2-(aq).
Types of Acids
Monoprotic Acids
HCl H+ + Cl- : Strong electrolyte, strong acid
HNO3 H+ + NO3- : Strong electrolyte, strong acid
CH3COOH H+ + CH3COO- : Weak electrolyte, weak acid