Background Research on Sodium Chloride:
Sodium chloride, also known as
common salt,
table salt, or
halite, is a
chemical compound with
formula NaCl.
Sodium chloride is the
salt most responsible for the salinity of the
ocean and of the
extracellular fluid of many multicellular
organisms. It is commonly used as a
flavour enhancer and
preservative for
food.
Sodium chloride forms
crystals with cubic
symmetry. In these, the larger chloride
ions are arranged in a cubic
close-packing,
while the smaller sodium ions fill the octahedral gaps between them.
Each ion is surrounded by six of the other kind. This same basic
structure is found in many other
minerals, and is known as the
halite structure.
http://www.chemistrydaily.com/chemistry/Sodium_chloride
Background on Carbon Dioxide:
Carbon dioxide (also known by its chemical formula of
CO2) is a chemical compound made up of two
oxygen atoms bonded to one
carbon
atom. It is a gas at standard temperature and pressure and exists in
Eart's atmosphere in this state. Carbon dioxide is one of the greenhouse
gases.
Carbon dioxide, in room temperature, is a colorless, odorless gas. Its solid state is kown as "dry ice."
Carbon Dioxide is an important gas on Earth. It makes up a fraction of
the atmosphere, and is inhaled by plants during photosynthesis. Its
abundance in the atmosphere is increasing due to the burning of fuels
that emit CO
2. Unfortunately, carbon dioxide has been
demeaned by the theory of how humans have contributed to global warming
through the burning of fossil fuels and whatnot. Regardless, carbon
dioxide is a very important gas and is one of the main components of air
(the others being gases like
nitrogen,
oxygen,
argon, etc.)
http://chemistry.wikia.com/wiki/Carbon_Dioxide
Solubility:
solubility, Degree to which a substance dissolves in a
solvent to make a
solution (usually expressed as grams of
solute per litre of solvent). Solubility of one fluid (liquid or gas) in another may be complete (totally miscible; e.g.,
methanol and
water)
or partial (oil and water dissolve only slightly). In general, “like
dissolves like” (e.g., aromatic hydrocarbons dissolve in each other but
not in water). Some separation methods (
absorption, extraction) rely on differences in solubility, expressed as the
distribution coefficient (ratio of a material’s solubilities in two solvents). Generally, solubilities of solids in liquids increase with
temperature and those of gases decrease with temperature and increase with pressure. A
solution in which no more solute can be dissolved at a given temperature and pressure is said to be saturated (
see saturation).
http://www.britannica.com/EBchecked/topic/553675/solubility
Solubility
Rules
A basic knowledge of which compounds are soluble in aqueous solutions is
essential for predicting whether a given reaction might involve formation of a
precipitate (an insoluble compound).
The following guidelines are generalizations.
A substance is classified as insoluble if
it precipitates when equal volumes of 0.1
M solutions of its components are mixed.
Keep in mind, however, that no substance
is completely insoluble. Substances listed
as insoluble are, at some level, partially
soluble. The magnitude of the ion product
constant (
Ksp)
for the appropriate solubility equilibrium
should be examined. Larger
Ksp
values indicate greater solubility; smaller
Ksp values
indicate lesser solubility.
The symbol "<=>" is used here to signify
the 'double-arrrow' symbol for a chemical
equilibrium. The symbol "=>"
is used here to signify the 100% dissociation
of a compound into its electrolyte ions
in aqueous solution. The subscript "(s)"
following a species indicates that it is
a solid. The subscript "(aq.)"
following a species indicates that it is
in aqueous solution.
Rule 1. All compounds of Group IA
elements (the alkali metals) are
soluble.
For example, NaNO
3,
KCl, and LiOH are all soluble compounds.
This means that an aqueous solution
of KCl really contains the predominant
species K
+ and Cl
-
and, because KCl is soluble, no
KCl is present as a solid compound
in aqueous solution:
KCl
(s) => K
+(aq.) + Cl
-(aq.)
Rule 2. All ammonium salts (salts of NH
4+) are
soluble.
For example, NH
4OH is a soluble compound. Molecules
of NH
4OH completely dissociate to give ions of
NH
4+ and OH
- in aqueous solution.
Rule 3. All nitrate (NO
3-),
chlorate (ClO
3-),
perchlorate (ClO
4-),
and acetate (CH
3COO
-
or C
2H
3O
2-,
sometimes abbreviated as Oac
-)
salts are
soluble.
For example, KNO
3 would be classified
as completely soluble by rules 1 and 3.
Thus, KNO
3 could be expected
to dissociate completely in aqueous solution
into K
+ and NO
3-
ions: KNO
3 => K
+(aq.)
+ NO
3-(aq.)
Rule 4. All chloride (Cl
-), bromide (Br
-), and
iodide (I
-) salts are
soluble except for those of
Ag
+, Pb
2+, and Hg
22+.
For example,
AgCl is a classic insoluble chloride
salt:
AgCl
(s) <=> Ag
+(aq.) +
Cl
-(aq.) (
Ksp =
1.8 x 10
-10).
Rule 5. All sulfate ( SO
4=) compounds are
soluble except those of Ba
2+, Sr
2+,
Ca
2+, Pb
2+, Hg
22+, and
Hg
2+, Ca
2+ and Ag
+ sulfates are only
moderately soluble.
For example, BaSO
4
is insoluble (only soluble to a very
small extent):
BaSO
4(s) <=> Ba
2+(aq.) + SO
42-(aq.)
(
Ksp = 1.1 x 10
-10).
Na
2SO
4 is completely
soluble:
Na
2SO
4(s) => 2 Na
+(aq.)
+ SO
42-(aq.).
Rule 6. All hydroxide (OH
-) compounds are
insoluble
except those of Group I-A (alkali metals) and Ba
2+, Ca
2+,
and Sr
2+.
For example, Mg(OH)
2 is insoluble
(
Ksp = 7.1 x 10
-12).
NaOH and Ba(OH)
2 are soluble, completely dissociating in aqueous
solution:
NaOH
(s) =>
Na
+(aq.) + OH
-(aq.), a
strong base
Ba(OH)
2(s) => Ba
2+(aq.)
+ 2OH
-(aq.) (
Ksp
= 3 x 10
-4)
Rule 7. All sulfide (S
2-) compounds are
insoluble
except those of Groups I-A and II-A (alkali
metals and alkali earths).
For example, Na
2S
(s) <=> 2Na
+(aq.)
+ S
2-(aq.)
MnS is insoluble (
Ksp
= 3 x 10
-11).
Rule 8. All sulfites (SO
3=),
carbonates (CO
3=),
chromates (CrO
4=),
and phosphates (PO
43-)
are
insoluble except for those of
NH
4+ and Group I-A
(alkali metals)(see rules 1 and 2).
For example, calcite, CaCO
3(s)
<=> Ca
2+(aq.)
+ CO
3=(aq.)
(
Ksp = 4.5 x 10
-9).
Solubility:
-
-
-
The average
kinetic energy of the
solute molecules increases with
temperature, destabilizing the
solid state. The increased vibration (
kinetic energy) of the molecules causes them to be less able to hold together, and thus they dissolve more readily.
-
solubility
The amount of a
substance that will dissolve in a given amount of a solvent to give a saturated
solution under specified conditions.
-
kinetic energy
The
energy possessed by an object because of its motion, equal to one half the mass of the body times the square of its velocity.
Solid Solubility and Temperature
The
solubility of a given
solute in a given solvent typically depends on
temperature. For many solids dissolved in
liquid water, the
solubility increases with
temperature up to 100 °C. In
liquid water at high temperatures (e.g., that approaching the
critical temperature), the
solubility of ionic solutes tends to decrease due to the change of properties and structure of
liquid water; the lower dielectric constant results in a less
polar solvent.
The chart (Figure 0) shows
solubility curves for some typical
solid inorganic salts. Many salts behave like barium nitrate and disodium
hydrogen arsenate, and show a large increase in
solubility with
temperature. Some solutes (e.g., sodium chloride in water) exhibit
solubility that is fairly independent of
temperature. A few, such as cerium(III) sulfate, become less soluble in water as
temperature increases. This
temperature dependence is sometimes referred to as retrograde or inverse
solubility. Occasionally, a more complex pattern is observed, as with sodium sulfate, where the less soluble decahydrate
crystal loses water of crystallization at 32 °C to form a more soluble anhydrous phase. The following figure shows
solubility curves for some typical
solid inorganic salts: Figure 0
THEORETICAL PERSPECTIVE
As the
temperature of a
solution is increased, the average
kinetic energy of the molecules that make up the
solution also increases. This increase in
kinetic energy allows the solvent molecules to more effectively break apart the
solute molecules that are held together by
intermolecular attractions.
The average
kinetic energy of the
solute molecules also increases, destabilizing the
solid state. The increased vibration (
kinetic energy) of the molecules causes them to be less able to hold together, and thus they dissolve more readily.
A useful heuristic is to consider that as the
temperature of a
solid is increased, it is closer to its
melting point, and thus closer to a
liquid, and it is easier to dissolve something that is closer to a
liquid.
APPLICATION IN RECRYSTALLIZATION
One application of the differential
solubility of compounds at different temperatures is to perform a recrystallization. An impure
substance is taken up in a
volume of solvent at a
temperature at which it is insoluble, but is heated until it becomes soluble. The impurities dissolve as well, but when the
solution is cooled it is often possible to selectively crystallize (precipitate) the desired
substance in a purer form.
https://www.boundless.com/chemistry/solutions/factors-affecting-solubility/solid-solubility-and-temperature/#key_term_glossary_temperature
Gas Solubility and Temperature The Effects of Temperature on the Solubility of Gases in the Universal Solvent (Water)
The
solubility of gases is dependent on
temperature. An increase in
temperature results in a decrease in
gas solubility in water, while a decrease in
temperature results in an increase of
gas solubility in water. To comprehend this phenomenon, one must consider the two processes that occur when a non-polar
gas is added to water. When adding the
solute
to the solvent, a type of cavity initially develops; this cavity is
representative of the conformation and overall size of the added
gas. In turn, a successive process occurs in which attractive forces between the
gas
and water molecules are stimulated. This dual process induces the water
to produce both attractive and repulsive forces. By examining the water
on a microscopic level and the components of the water that exert
attractive forces to the non-polar gases,
temperature dependencies become observable.
The Solubility of Gases in Organic Solvents
Gases
assimilated with organic solvents become more soluble at higher
temperatures. Le Chatelier's Principle can be applied to the
determination of why the
solubility of gases increases with inclining temperatures. The principle proclaims that when a system is placed under stress, an
equilibrium shift will occur in the direction that will most relieve the stress. In relation to the principle, adding
heat to a
solution will induce a shift in the
equilibrium that favors
dissolution in order to
reduce heat. On a molecular stance, because organic solvents are incapable of forming
hydrogen bonds with gases (in contrast to water), more
heat is released when a
gas is placed in water than in an organic solvent. A synthesized conclusion is that stronger attractions between a solvent and
solute and
entropy contribute to a greater transfer of
heat, or enthalpy. Thus, increasing the
temperature of the
solution will favor an
endothermic shift in the
equilibrium of the system. As a result, the
solubility of gases increases with increasing
temperature in organic solvents.
Making Connections: The Phenomena of Gas Solubility
The reasoning for this relationship between
temperature and
gas solubility is similar to that of
temperature and
vapor pressure. An increase in
temperature causes an increase in
kinetic energy, resulting in a more rapid motion of molecules and the breaking of
intermolecular bonds, which enables molecules to escape from the
solution. This is an example of the Second Law of Thermodynamics.
https://www.boundless.com/chemistry/solutions/factors-affecting-solubility/gas-solubility-and-temperature/
