Introduction: During this experiment, the solubility of different ions in water is being evaluated. Solubility is a phys

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Introduction:
During this experiment, the solubility of different ions in
water is being evaluated. Solubility is a physical property that
defines how much of a solute, the substance in the smaller amount,
can be dissolved in the solvent, substance in the larger amount.
From both predictions and observations, some assumptions about this
solubility are being made, namely that the ionic compound is
classified as either soluble or not soluble in water. However, this
is an oversimplification of the property of solubility. Solubility
is better represented by considering it as a range with some ionic
compounds being more soluble in water than other ionic compounds.
Most ionic compounds have at least some solubility in water, albeit
it may be very, very small. Some of these differences should be
noticeable during this experiment and should be logged as
qualitative data. However, this simplified classification system
has an advantage in that it helps to organize and better understand
the relationship between different ions and their solubility in
water.
Figure 1: Solubility is best considered as a
range. An ionic compound that is water soluble would
be considered to have a high solubility; whereas an ionic compound
that is water insoluble would have a low solubility. The instrument
used to detect the solubility will limit the precision of such a
classification, the most simply of which is either soluble or not
soluble.
Ions are atoms that carry an overall charge because they have
different numbers of electrons and protons. There are two types of
ions, cations and anions. Cations are positively charged as
they have lost electrons to obtain a full outer shell of valence
electrons. Anions are negatively charged as they have gained
electrons to obtain a full outer shell of valence electrons.
Ionic compounds, in their simplest form, are comprised of a
single type of cation and a single type of anion. These oppositely
charged ions are held together by electrostatic interactions in a
similar manner as to how two magnets are held together, with
opposite poles being attracted to one another. All ionic compounds
have these electrostatic interactions. When two ions of opposite
charge come close enough together, these attractive forces will
draw them even closer together. Many ionic compounds are found as
solid salts.
When solid ionic compounds are put into water, there are two
main outcomes. Either the ions will separate from each other and
form an aqueous solution, or the ions will remain together in a
solid form.
When the ions separate from each other, or dissociate, and form
an aqueous solution, these compounds are considered water soluble.
During this process, the ions in the solid ionic compound
dissociate from each other to instead interact with the water
molecules through electrostatic interactions that are similar to
those interactions between the ions themselves but are weaker than
those found between the ions. However, with enough of these weaker
interactions, the ions can be persuaded to interact with the water
molecules instead of with each other. This is a result of the
ion to water interactions being a more energetically favorable
situation than with the ions interacting just with each other. (See
Figure 2A).
When the ions remain together in a solid form, these compounds
are considered water insoluble. In these cases, the ions
interacting with each other are more energetically favorable than
the individual ions interacting with the water molecules. (See
Figure 2B)
Figure 2: Ionic compounds can be water soluble or
water insoluble. The combination of water soluble
ions results in a clear aqueous solution (A); whereas, the
combination of water insoluble ions results in an opaque solution
that over time will result in a solid precipitate settling to the
bottom of the container (B).
When an ion can either interact with another ion or with water
molecules, the option that is more energetically favorable will be
the preferred option, thus making these types of interactions
predictable. Scientists have discovered this predictable nature by
combining two aqueous ionic solutions together and observing the
outcomes of lots of different aqueous ionic solution pairs. By
keeping track of the outcomes of these combinations of solutions,
they have been able to put together tables of data for the
solubility of lots of different ions in water. Figure 3 is a
condensed version of these solubility rules.
Figure 3: Condensed solubility rules
chart.
You may also refer to this document here.
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Precipitation reactions are reactions that involve the
combination of two aqueous ionic compound solutions that then form
a solid, or a precipitate, as one or more of the products.
During this process, the products themselves can be predicted based
on the ions that are involved in the original solutions. Again, two
different outcomes are often observed, either no precipitate will
form, and the solution will remain clear, or a precipitate will
form, which will obscure the mixture so that it is now considered
opaque, or cloudy.
You will be provided links to videos to watch where experiments
have been performed to determine whether a precipitate will
form. Before you watch the videos, you will predict what will
occur by writing the balanced equations for the reactions using the
solubility rules. You will indicate which product will be a
solid (s) or remain as ions in solution (aq). To evaluate the
solubility of different ions in water, the people in the videos
combine two aqueous ionic solutions together. You will record
the results, and then you repeat this process for several other
solutions. After watching each experiment, you will be able
to check if your predictions are correct.
Examples:
Here are two examples to better understand the rules of
solubility in precipitation reactions. There are several steps that
can be followed in order to write a complete balanced precipitation
reaction.
Example 1: Consider the combination of aqueous barium
chloride and aqueous sodium bromide.
Figure 4: Ions Barium, Chloride,
Sodium, and Bromide in solution.
Step 1: Write the correct ionic compound formulas for
the reactants.
Aqueous barium chloride =
BaCl2 (aq)
Aqueous sodium bromide = NaBr (aq)
Step 2: Determine the products and the correct ionic
compound formulas for the products.
When ions dissociate, they are considered to be separate from
each other, or floating around in the water. This leaves these ions
available to interact with other ions that they come across.
In this case, the barium cation may come in contact with the
chloride anion, the sodium cation, or the bromide anion. When the
barium cation contacts the chloride ion, as the original solution
is comprised of these two ions and the solution is aqueous, so the
ions will pass by each other, still happy to interact with the
water molecules. When the barium cation contacts the sodium cation,
they will repel each other, as they are both positively charged and
have no interest in each other. However, when the barium cation
contacts the bromide ion, then there are two options. Either the
ions stay separated, causing the solution to remain clear, or they
join together as an ionic compound and form a solid precipitate
causing an opaque solution. The solubility rules in step 4 will be
used to determine which is the case for this set of ions. But this
means that there are two potential products, determined by the
cation of the first compound combining with the anion of the second
compound to form the first product and the anion of the first
compound combing with the cation of the second compound to form the
second product.
Note that when combing the anions and cations, do NOT carry the
subscript. Each product formula should be written properly for an
ionic compound, with the sum of the charges on the anions and
cations in that compound equal to zero.
BaCl2 (aq) + NaBr (aq)
—> BaBr2 (?) + NaCl
(?)
Step 3: Balance the chemical reaction using
coefficients. Usually these balances are reasonably
straightforward. If you are having difficulty balancing these
equations, go back and check that all of the ionic compound
formulas for your products and your reactants are written
correctly.
BaCl2 (aq) + 2 NaBr (aq)
—> BaBr2 (?) + 2 NaCl (?)
Step 4: Use solubility rules for ionic compounds (Figure
3) to “predict” whether the products are likely to be clear (water
soluble, represented by “aq” for aqueous) or opaque (water
insoluble, represented by “s” for solid). One
rule indicates that bromide anions are soluble in water unless
combined with silver ions, mercury (II) ions, or lead (II) ions.
For the product barium bromide, none of the exceptions are met,
thus barium bromide would be predicted to be water soluble, or
aqueous.
BaBr2 (aq)
Another rule indicates that sodium ions are soluble with no
exceptions and that chloride anions are soluble in water unless it
is combined with silver ions, mercury (II) ions, or lead (II) ions.
For the product sodium chloride, none of the exceptions are met,
thus sodium chloride would be predicted to be water soluble, or
aqueous.
NaCl (aq)
Step 5: Make conclusions based on your complete balanced
chemical reaction for this combination of aqueous ionic compound
solutions.
BaCl2 (aq) + 2 NaBr (aq)
—> BaBr2 (aq) + 2 NaCl
(aq)
No precipitate will form since neither of the products are water
insoluble.
Example 2: Consider the combination of aqueous barium
chloride and aqueous silver nitrate.
Figure 5: Ions Barium, Chloride,
Silver, and Nitrate (NO3- = polyatomic
ion) in solution.
Step 1: Write the correct ionic compound formulas for
the reactants.
Aqueous barium chloride =
BaCl2 (aq)
Aqueous silver nitrate = AgNO3 (aq)
Step 2: Determine the products and the correct ionic
compound formulas for the products.
There are two potential products, determined by the cation of
the first compound combining with the anion of the second compound
to form the first product and the anion of the first compound
combing with the cation of the second compound to form the second
product.
Note that when combing the anions and cations, do NOT carry the
subscript. Each product formula should be written properly for an
ionic compound, with the sum of the charges on the anions and
cations in that compound equal to zero.
BaCl2 (aq) +
AgNO3 (aq) —>
Ba(NO3)2 (?) + AgCl
(?)
Step 3: Balance the chemical reaction using
coefficients. Usually these balances are reasonably
straightforward. If you are having difficulty balancing these
equations, go back and check that all of the ionic compound
formulas for your products and your reactants are written
correctly.
BaCl2 (aq) + 2
AgNO3 (aq) —>
Ba(NO3)2 (?) + 2 AgCl
(?)
Step 4: Use the ionic compound solubility rules in your
textbook to “predict” whether the products are likely to be clear
(water soluble, represented by “aq” for aqueous) or opaque (water
insoluble, represented by “s” for solid). One rule
indicates that nitrate anions are soluble in water and that the
combination with barium is not an exception, thus barium nitrate
would be predicted to be water soluble, or aqueous.
Ba(NO3)2 (aq)
Another rule indicates that chloride anions are soluble in water
unless it is combined with silver ions, mercury (II) ions, or lead
(II) ions. Therefore, the combination of the chloride anions with
the silver cations, means that these ions prefer interacting with
each other rather than interacting with the water molecules. Thus,
silver chloride would be expected to form a solid in water causing
the solution to be opaque.
AgCl (s)
Step 5: Make conclusions based on your complete balanced
chemical reaction for this combination of aqueous ionic compound
solutions.
BaCl2 (aq) + 2
AgNO3 (aq) —>
Ba(NO3)2 (aq) + 2 AgCl
(s)
A precipitate of silver chloride will form while barium ions and
nitrate ions remain in solution.