This course is designed to cover subjects in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. This course is a precursor to the Advanced Chemistry Coursera course. Areas that are covered include atomic structure, periodic trends, compounds, reactions and stoichiometry, bonding, and thermochemistry.
4.02 Aqueous solutions
Loading...
Del curso dictado por University of Kentucky
Chemistry
546 calificaciones
De la lección
Week 4
We will explore how compounds react with one another to form new substances and then write balanced chemical equations to represent what is happening in a reaction. We will explore several different types of reactions including precipitation, acid-base, oxidation-reduction, and combustion reaction.
Conoce a los instructores
Dr. Allison SoultLecturer
Chemistry
Dr. Kim WoodrumSr. Lecturer
Chemistry
In this module, we're going to talk about aqueous solutions.
By the end of this module, you should be able to describe the properties and
characteristics of aqueous solutions.
When we look at solutions,
there are two main distinctions that we're going to worry about.
One is homogeneous, the other is a heterogeneous mixture.
When we have a solution, what we see is that all of the solute, or
all of one of the substances, is dissolved in the other.
For example, if I look at the beaker on the left,
what I see is that the salt is completely dissolved in the water.
On the right side, I have a homogeneous mixture because the sand has settled to
the bottom because it does not dissolve in water.
When I create my solution, I have two things I have to worry about to
make this solution, one of which is the solute, and
that will always be the substance present in the smaller amount.
And then I have the solvent, which is the substance present in the larger amount.
We then combine both of these together, the solute and the solvent, and
together, they make up the solution.
By far, the most common type of solution we'll look at are aqueous solutions.
A few things to note here.
We have four examples of aqueous solutions.
While some of them are colored and
one of them is colorless, notice that they're all transparent.
This is a characteristic of our aqueous solutions.
Note that these are still considered clear even though they have color
because clear and colorless mean two different things.
Clear means something is transparent,
whereas colorless means something has no color.
So I see that these last three solutions are clear, but they do have color.
These are still aqueous solutions because when we talk about an aqueous solution,
we simply mean that it has been dissolved in water.
Solutions don't have to be aqueous, and nor
do they have to be in the liquid phase.
All of the substances you see here are actually solutions,
just not all aqueous solutions.
For example, first, we want to look at air.
Air is actually a solution.
The solvent is nitrogen because that's the component present in the larger amount,
and oxygen is the primary solute.
Note that there are many other solutes in the air, but
they're present in much smaller amounts.
Next, we can look at something like club soda, or carbonated water.
In this case, it is an aqueous solution because our solvent is water, but
we also see that our solute is CO2, which is a gas.
So we have a gas dissolved in water, and that also makes up a solution.
Next, another aqueous solution when we have vinegar.
We took acetic acid, which is in the liquid state, and dissolve it in water.
The acetic acid is present in the smaller amount, so
it's our solute, and our solvent is water.
Now we can move on to some things that are not aqueous and
are not in the liquid phase.
When I look at a dental amalgam or
a filling, what I see is that the solvent is actually silver.
That's the component present in the larger amount.
But I also have solute of mercury, which present in a smaller amount.
So this is considered a solution because of, it contains two components,
one of which is present in a larger amount, one in a smaller amount, and
we have a homogeneous combination of these two components.
When I look at the iodine, what I see is that the solute is iodine and
the solvent is ethanol.
And our last example is actually steel.
And this one then, it's hard, a little bit harder to see,
because we think about steel and don't think of that as a solution.
But what we have is our solvent is iron and our solute is carbon.
We tend to associate solutions with the liquid phase, but
it doesn't always have to be that way.
Now here's an example for you to identify the components of a mixture of
a 20 grams sample of NaCl with 50 grams of water.
Because water is present in the larger amount, it's going to be the solvent.
Because NaCl is present in the smaller amount, it's going to be our solute.
Together they make up a solution.
Now we want to look a little bit at how these things behave in water.
What happens for aqueous solutions when something dissolves in water?
What happens to the molecules that make up that solute?
And then how does that affect the properties of that solution when we
have different substances dissolved in water?
The first example we'll look at is our non-electrolyte.
In this case, the molecules will dissolve in water, but they don't dissociate in
water, meaning they don't break apart into their little components.
If we look at something like sucrose, where we have a large organic compound,
what we see is that it has, it remains intact.
So if we start with sugar, or sucrose in the solid form, and add it to water,
when it dissolves in water, those sugar molecules will stay intact.
They won't break apart into ions.
On the other extreme, if we look at a strong electrolyte, something like
sodium chloride, what we find is that when we add solid sodium chloride to water,
it will dissolve, and it will actually break apart into sodium ions and
chloride ions, and those will now be our solute.
The ions will stay separate from one another in the water.
We will not have units of NaCl present.
And these are strong electrolytes because they completely dissociate in the water.
In the middle, we have weak electrolytes, and
these are compounds that partially dissociate in water.
The most common weak electrolytes that we see are actually weak acids.
In this case, we're looking at acetic acid, and
what we see is that acetic acid partially dissociates in water.
A small fraction of the acetic acid molecules will break up into the ions, and
here we see a cation and an anion.
But notice that the majority of the molecules will stay intact as
the acetic acid molecules.
So we can write these equations to represent what's going on.
For a strong electrolyte, we use a regular reaction arrow that indicates it
goes completely from the left, our reactants, to the products on the right.
For a non-electrolyte, we use the same kind of arrow, only now what we notice is
that we're not breaking the formula down as we go from reactants to products, but
that we do change the phase of the substance.
For a weak electrolyte, we use a new kind of arrow that we haven't seen before.
This is known as a equilibrium arrow.
And in the equilibrium arrow, what we have is it indicates that the reaction's going
in both the forward and the reverse direction at the same time.
Now, if I look at acetic acid molecule here, what I see is that a small fraction
of those molecules will dissociate into the ions, and this will actually continue.
The acetic acid molecules will continue to dissociate, but what I
will also see is that some of my ions will then recombine and form the reactant.
So this arrow, this equilibrium arrow, indicates that the reaction is going in
both the forward and the reverse direction.
If I were able to tag one of these acetic acid molecules and monitored it,
what I would notice is that most of the time, it's present as a molecule.
Occasionally, it's also present as an ion.
But if I check later, it might be back in the molecular form.
Some examples of our electrolytes.
For strong electrolytes, we see anything that we recognize as a strong acid or
a strong base, such as sodium or potassium hydroxide, and our ionic compounds.
All of these substances are strong electrolytes.
Then we can see our weak electrolytes, things like acetic acid,
which is a weak acid.
HF, HNO2 are weak acids.
Water is also a very weak electrolyte.
When we look at our non-electrolytes,
what we notice is that we have larger organic molecules.
What we see is that these compounds stay intact when we add them to a solution.
In this demonstration, we're going to look at the conductivity of a few
different solutions.
What we have here is a conductivity meter, and
we have our probe attached to a very high-tech paint stirrer.
And what we're going to see is that we can actually get the conduction of
electricity through this solution.
So it's going to conduct between these two probes.
So the first thing I'm going to do I'm going to put it in water, and
what I see is I see just one light coming on.
And when that happens, I see a small amount of conduction of electricity.
In order to conduct electricity, we need ions in order to
carry that electrical current from one probe to the other.
And water only has a very small,
small amount of ions present from the dissociation of water.
Now we can look at a solution of sugar.
And what we know about sugar is that it does not dissociate in water, so
we don't get the formation of ions.
So we see a result similar to what we saw with the water.
However, when we go to the salt solution, the sodium chloride solution,
what we know is that NaCl is going to dissociate in water and
form sodium ions and chloride ions.
And when we do,
we see a lot higher levels of conductivity in solution because there are ions present
that are able to transfer that electrical current from one probe to the other.
So let's look at our experiment down here that we're demonstrating.
We have a conductivity indicator, and what it's showing is that for
the non-electrolyte, we see very little connection of electricity.
The tips of the probes are not in physical contact with one another, and so
they rely on anything in the water, in the solution, to be able to transport that
charge from one side to the other so that we get the conduction of electricity.
For non-electrolytes, we have little to no conductivity.
When I look at a weak electrolyte, now I have some ions present in solution.
So I can actually get conduction of electricity in that solution because I
have some ions.
When I go to the salt solution where it's completely dissociated because it's
a strong electrolyte, I have a lot of ions present in solution, so
we see lots of anions and lots of cations.
And in fact, we see no, none of our intact molecules.
And what we see is that we have very good conduction of electricity.
And so we're able to get very high on the conductivity meter because there's lots of
ions in this solution to transport that charge from one probe to the other.
So, when we look at the strong versus the weak electrolytes, we need to
recognize that there are certain things that are strong acids and bases.
And the reason we call them strong acids and
strong bases is because they completely dissociate in water.
So we have six strong acids that we know, HCl, HBr, HI, HNO3, H2SO4, and HClO4.
If we remember those six strong acids,
then we know that any other acid we come across is going to be a weak acid,.
Likewise with our bases, we see six examples here, lithium, sodium,
potassium hydroxide, calcium, strontium and barium hydroxide.
Other bases we can assume that they are weak bases.
And so, weak acids and bases, just like weak electrolytes, partially ionize or
dissociate, however you want to call it, in water.
These two terms actually mean the same thing.
Dissociate means break apart into ions.
Ionize means break apart into ions.
So if we have something that partially dissociates, then we have a weak acid or
a weak base.
Now, note that weak does not mean not dangerous.
Weak simply indicates the amount of dissociation of that substance in water.
For something like HF, or hydrofluoric acid, it is actually a very
dangerous acid, but it is a weak acid because it doesn't completely dissociate.
If you've ever seen etched glass, that may have been done with hydrofluoric acid
because it actually can dissolve glass.
We keep it in a plastic bottle, usually a Teflon-lined bottle, so
that it can't dissolve its container.
And we also have to be very careful when we handle these this acid because very
brief exposure can cause a lot of damage.
The HF molecules are very small, and they can have a lot of
residual effect both on the skin, as well as absorbing into the skin and
start causing problems within your bloodstream.
So some things we have to be more careful about than others, but
weak is not dangerous.
Weak acid can still be a dangerous acid, as we see with HF.
In the next module, we're going to look at the solubility of ionic compounds.
Explora nuestro catálogo
Inscríbete de manera gratuita y obtén recomendaciones personalizadas, actualizaciones y ofertas.
Coursera brinda acceso universal a la mejor educación del mundo, al asociarse con las mejores universidades y organizaciones, para ofrecer cursos en línea.
© 2019 Coursera Inc. Todos los derechos reservados.
