This is an introductory course for students with limited background in chemistry; basic concepts such as atomic and molecular structure, solutions, phases of matter, and quantitative problem solving will be emphasized with the goal of preparing students for further study in chemistry.
2.1 Introduction to Solutions Part I
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En provenance du cours de Duke University
Introduction to Chemistry: Structures and Solutions
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Solutions
<p>Welcome to week 6! We have reached the last week of video lectures and exercises for the course! Are you ready to make the final push through the rest of the material? After that, you should be ready to tackle the final exam which opens next week. I have enjoyed working on this course; and I hope you have too!</p>
<p>This week, we will first release a couple of review videos related to solution and solubility from <i>Introduction to Chemistry: Reactions and Ratios</i>. Please review these important concepts before starting Lesson Two - Solutions. Then we will start with a review of dissolution and covers some concentration units. This is accompanied by solution-related demonstrations and another bonus interview with a chemist who also works at Duke. The interviews were conceived and filmed by Abdul Latif, an undergraduate in the humanities who is interested in the history of chemistry and how people choose their professions. Enjoy!</p>
Rencontrer les enseignants
Prof. Dorian A. CanelasAssistant Professor of the Practice
Chemistry
[MUSIC]
In this video, I'll be talking about solutions, and
more specifically, aqueous solutions, which are compounds dissolved in water.
Now the biggest aqueous solution on Earth is of course, the oceans,
the oceans contain almost entirely water, and that's a chemical.
But there are also other compounds dissolved in the water, such as oxygen, so
that the fish can breathe, and all kinds of salts.
We know that the ocean is salt water,
those salts produce a wide variety of ions when they dissolve.
Here's a table that shows some of the ions that are present in typical samples of sa,
ocean water.
You can actually taste these compounds if you get a mouthful of water in the ocean,
you can taste the dissolved Sodium and chloride ions.
Let's look at the structure of water first, and then we'll move on to look at
the structures of the ionic compounds that can dissolve in water.
Obviously water is the most abundant component of the ocean, it's the solvent,
the species that's present in the greatest amount in a solution.
So I think starting to talk about water is a good place to start.
First, I'm going to show you some cartoons that give
the structure of a water molecule.
Water, of course, has the chemical structure, H-2-O.
And I can draw a little cartoon that shows, in this case,
the oxygen is shown as yellow, and that's just my choice to pick yellow.
And the two hydrogens are shown as small pink spheres, and
again, pink was just my choice.
In this case we know that the oxygen has more electro-negatives than the hydrogens.
So water molecule has part of the molecule that has
a greater density of electrons than the other part, in other words,
the oxygen, which here is yellow is, pardon me, partially negative.
[SOUND] And the two hydrogens which are not as electronegative,
are not as attractive to the electrons and so they're each partially positive.
This little delta symbol a greek lower case d means partially,
it's showing a partial charge.
And bulk water at room temperature and normal atmospheric pressure,
there are billions, and billions of water which are very loosely organized.
We can look at the sample over here on the left,
if we were able to zoom in to look at just a few water molecules.
We would see that they're grouped together and they're not particularly well
organized, they can move around, they're very fluid, the waters can spin in
place and they can easily slide past each other, because it's in the liquid state.
Sodium chloride on the other hand is a solid salt, and the particles
in sodium chloride are extremely highly organized into a lattice as a solid state.
Here are some cartoons depicting the Sodium ion and
chloride ions in a salt crystal.
This is just what a crystal of table salt would look
like if you zoomed in with a microscope.
So, if I take a little,
tiny bit of that and I am able to use a microscope to see it better,
I can see that there are alternating Chloride ions and Sodium ions, but
they're very well organized into a lattice where I alternate between chloride,
Sodium, Chloride, Sodium, Chloride, as I move through the arrangement in space.
I'm going to show these throughout this presentation as these spheres,
the Sodium's going to be a pink sphere or kind of a reddish sphere with a positive
charge, and the Chloride's going to be a blue sphere with a negative charge.
They aren't exactly drawn to scale.
A Sodium cat ion actually has a radius of 116 picometers while a Chloride
ion has a radius of 167 picometers.
But you get the idea,
the chloride is slightly larger and it's negatively charged.
Now, this picture is a little bit misleading because there's not
exactly water in these empty spaces and
in fact, there's not as much empty space as is shown.
The Chloride ions and the Sodium ions are touching-
Along-
These axes like this, so
they're shown much smaller than they really are in this picture.
I just want people to understand that most of this space is filled up
by the electrons of the ions.
It's not water and it's, some of it's empty space, but
most of it is filled up by electrons.
What is happening in the dissolution process of ionic compound?
We can draw the chemical reaction equation for
this process and from this chemical equation it looks like the ions
are separating when the solid Sodium chloride is added to the liquid water.
And that's exactly what happens, this is called a dissolution equation.
It's the first type of chemical equation that we're learning to write in this
class, a dissolution equation shows a species that is solid dissolving and
turning into aqueous species.
In this case, because the solid was an ionic compound, when it dissolves, it,
those ionic compounds dissociated from each other to make two separates ions,
the cations and the anions.
So, the Sodium and the chloride ions remember, initially were touching,
if we just to show two of them we could show them touching,
they're touching in that lattice, in a very large three dimensional array.
And once we dissolve them in water they are no longer touching each other,
they dissociate from each other.
So now the Sodium cation and the chlorine anion are separate they can move around
separately, and they can react separately on the product side of this equation.
So the reactant was the Sodium chloride, and the products here, are these
anions and cations which are both aqueous because they're now dissolved in water.
Each Sodium chloride unit makes 2 ions, those are the ions that form.
And what happens is that those ions don't just move apart from each other and
dissociate, but water gets in between the ions and
keeps them from going back together.
We know from Coulomb's law that the opposite charges are attracted to
each other, and it's only this layer of water,
which might have even more water molecules than is shown there probably has many more
actually, it's that layer of water, that is able to prevent the cation and
the anion from going back together and making the solid again.
In fact, it's not just the pair of ions that separate.
In fact, many, many ions that were in that piece of salt, separate
during the dissolution process, and I've shown just a few of these in a cartoon so
you can imagine this happening with billions of these Sodium cations and
Chloride anions.
I start with them all touching each other and that crystal lattice, but then once I
add water to the salt the ions start to break apart, here they are broken apart.
And the reason they're breaking apart is
that the water can get in between the ions.
And there's some calculations you can do with chemistry math called thermodynamics
that you can calculate the heat that is either absorbed or
evolved by this process of performing a solution, the solvation process.
To summarize, the water molecules go between the cations and the anions and
keep them separated, and that is the case as long as the compound is dissolved.
There's a limit to how much will dissolve in the solution, and
we'll talk about that a little bit later.
There's an animation video that one of my colleagues made of this.
And I hope you'll check it out because it really shows the process better than I
can on these two dimensional slides.
Let's do more examples of ionic compounds dissolving in water.
Let's write the dissolution equations.
Let's start with Barium acetate, as the first example that we do.
Now, in order to do this problem,
we need to know what the chemical formula of Barium acetate is.
That means we either need to look up the formula of acetate, or
we need to have it memorized.
It also means we need to know the charge on both the Barium cation and
the acetate anion.
So Barium is an alkaline earth metal, it's in group two, it has a preferred charge of
plus two and acetate anion has a charge of minus one each.
Barium Acetate is this compound, so that's the reactant, Barium acetate.
If I dissolve it in the presence of water, what will the products be?
Remember when ionic compounds dissolve in water, they dissociate, the split apart.
So the Barium is going to split apart from the acetates to make
a Barium two plus cation, and then the acetate anions will also separate.
They will not only separate from the Barium, but
they will separate from each other.
However, that chunk of polyatomic ion, the C2H3O2,
will stay together, so we will show that like this.
So here is the acetate polyatomic ion that has stayed in tact, and
here is the Barium cation that has separated,
there is water in between those now, because it's aqueous.
How would we balance this equation?
Is the equation balanced right now?
Do I have the same number of each type of atom on both the product and
reactant side here?
In other words where does the two go?
I have a two right here don't I.
I have four oxygens on the left side but only two oxygens on the right side.
I can balance it by putting the two out in front of the acetate as a coefficient.
So the subscript of the solid becomes the coefficient of the aqueous ion,
and the only thing you need to be careful about is recognizing those
polyatomic ions that don't break apart, remember most of them are anions.
We learned a couple of cations but most of those that we learned were anions like
Sulfate, Carbonate, Phosphate, Nitrite, those types of things.
Those are the polytonic anines that stay together, when the species dissolves.
So when 1 molecule of Barium acetate dissolves, 3 ions were produced.
So, a little bit different than the Sodium chloride case because in that case
the Sodium and the chloride broke apart to give two ions per unit of Sodium chloride,
and in this case we're making three ions.
That becomes important later when you do some calculations of
calligative properties for example.
And we can use the stoichiometry to do more calculations, in other words,
what if I dissolved four molecules of Barium acetate?
How many ions would be produced,
if I dissolved four molecules of Barium acetate?
And I'm looking for
total number of ions, the number of Barium ions plus the number of acetate ions.
How many total ions would that be?
That's great,
you were able to use the stoichiometry, to calculate that it would be 12 ions.
You could also say if four moles of Barium acetate dissolve that would make 12 moles
of ions, 4 moles of ions would be Barium, cation, Barium 2+,
and the other eight moles of ions would be the acetate anions.
So you can do this by either thinking about individual molecules and
individual ions, or groups of molecules, as in moles of molecules, or
you can say dozen molecules, if you want.
But we're going to say moles of molecules in chemistry.
Here is one for you to try on your own.
Go ahead and pause the video, oh I guess an in video question will pop up for
you and pause it for you.
Write the solution equation for solid Copper (II) chloride in water and
that is show here in these photographs, here is what Copper (II) chloride looks
like, it is kind of this beautiful blue green, crystal and solid.
We dissolve it in water and it makes a similarly colored aqueous solution.
Hopefully what you wrote started with the correct chemical formula for
Copper two chloride which is CuCl two.
You do need to write the phase.
This is a solid, that's the reactant.
It dissolves in the presence of water, so we need to have an arrow and
have water shown because that's where, what you are using as a solvent.
And for products, the cation is Copper 2 plus, and the anion is chloride 2 minus.
But to balance it, we need to put the 2 as a coefficient in front of the chloride.
So, again, in this case,
each unit of Copper (II) chloride that dissolves gives three units of ions.
One of the things I can do is depict the product of this reaction as a cartoon,
analogous to the way I showed a cartoon for the Sodium and
the Chloride dissolved in the water.
In this case I have a Copper 2 cation right there,
dissolved in water, and I have two chloride anions.
So, if I just dissolved one unit,
of Copper chloride, I've been able to make three ions.
One of the things that's interesting to look at in this particular cartoon that
was shown fairly carefully is the way that
the water is organized around each item type.
So if I draw a little quick sketch of water up here, here's the oxygen,
here's the two hydrogens, drawing them a, not quite to scale, but you get the idea.
The oxygen is more electronegative so
remember it's got a partial negative charge on it and the two hydrogens [SOUND]
are less electronegative, so they're each partially positive in charge.
That's because the oxygen hogs the electrons of the oxygen hydrogen bond.
So then when waters are near a Copper 2 cation, for
example, they want to be organized so that the negative part of the water, which is
the oxygen, is closest to the copper 2+ cation, because opposites attract.
You can see that here in, in each case,
the oxygen is actually touching the copper 2 plus.
We say that this is hydrated,
it actually has six waters directly around that copper 2 plus.
And there's one that's out in front of the screen that you can't see and
one behind, and then there's the four that are in the plane of the screen.
And all of those are organized so
that the oxygen is closer to the Copper 2 plus than the hydrogens are.
In contrast, look at how the water is organized around the chloride ions.
In that case, the hydrogens, which are the more positive part of the molecule,
prefer to be closer to the negative charge of the chloride ion.
So even in the solution which we think of being highly disorganized and
it is because of course the cat ions and n ions can move around freely and the waters
can slide past each other, there is this level of organization that comes from
Coulomb's law from the attraction of positive and negative charges.
In this case between the solute, which here is the Copper (II) chloride.
That's the thing being dissolved which is in the lesser amount.
And the solvent, which for aqueous solutions is water.
Here's another one for you to try.
Let's practice by writing the dissolution equation, for
making a solution of copper II sulfate in water.
So here I have the same cation but a different anion.
Go ahead and try to write that on your paper now, the dissolution equation for
Copper II sulfate.
Hopefully, what you wrote on your paper began with a reactant that is
CuSO4 with a solid sign next to it, that's Copper II sulfate, you dissolve in water.
That produced on 1 Copper 2 plus cation, and one sulfate polyatomic anion.
Do we need to have any kind of coefficient in front of a sulfate?
We don't actually need that in this case, do we?
It's already balanced.
Each Copper sulfate makes two ions; one Copper 2 plus ion and one sulfate anion.
In all these cases when we dissolve the ionic compound,
the cation dissociates from the anion.
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