2.2 Introduction to Solutions Part II

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Из курса от партнера Duke University

Introduction to Chemistry: Structures and Solutions

79 оценки

Duke University

Introduction to Chemistry: Structures and Solutions

79 оценки

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.

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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>

2.1 Introduction to Solutions Part I 17:41

2.2 Introduction to Solutions Part II 11:45

2.3 [Demo] Electrolytes10:30

2.4 Dissolution and the Solubility Product Constant Part I 15:20

2.5 Dissolution and the Solubility Product Constant Part II11:58

2.6 Solubility Rules 13:20

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Prof. Dorian A. Canelas
Assistant Professor of the Practice
Chemistry

So this is what happens for the cases of the acetate and

the sulfate in the previous examples.

They stay together, but the cation dissociates from the anion.

When an ionic compound dissolves in water, the cation and

the anion are separated from each other by these layers of water.

They get in between them, this is called dissociation.

Not all ionic compounds are highly soluble in water and

we'll talk about how we can measure the solubility and

express it with a number, in one of the later lectures this week.

In addition, some groups of

primarily non-metals like to stick together, and those are those groups of

polyatomic ions that we learned when we were learning how to name ionic compounds.

In this case, the compound that dissolved with silver nitrate.

And when it broke apart the nitrate,

the NO3 one minus, groups of polyatomic ions stay together.

But the silver separated from the nitrate didn't it?

So the cation always separates from the anion but

sometimes these polyatomic cations or polyatomic anions stay together.

This is, this type of sticking together of groups is

what was happening with the acetate and the sulfate in the previous examples.

But in all of these examples we've seen the metals ions completely separated and

end up being by themselves.

Completely dissociated from anything else.

Let's try one more of these just to make sure you're getting comfortable with

the concept.

On some paper, let's write out the dissolution equation for

sodium phosphate in water.

How many moles of ions are produced for

each mole of sodium phosphate that dissolves?

So, I want you to think about dissolving one mole of sodium phosphate,

write the dissolution equation, and

answer with how many moles of total number of ions are present in that solution.

Here's how we would do that problem.

Sodium phosphate solid has the equation Na3PO4.

We'll dissolve that in water.

In this case,

we make three sodium cations, because there was a three after the metal.

We want to make sure that's balanced.

And only one phosphate polyatomic anion, which has a charge of -3.

So, each sodium phosphate that dissolves, each unit makes 4 ions.

So, one mole of sodium phosphate dissolving would make four moles of ions.

So far, what we've been doing is counting the number of ions that

are forming relative to the solid that we've dissolved.

But we also need to know how much is dissolved in a given solution.

Is the salt water have just a tiny little bit of salt, or

is a lot of salt dissolved, for example?

The solute, which is present in the lesser amount is dissolved in the solvent.

In all of the past examples, water was the solvent, but

there are other liquids that could be the solvent.

You could use gasoline as a solvent, for example.

That would be a completely different type of solution than these aqueous solutions,

but you could still use it as a solvent

The solvent is always the item that is present in the greater amount.

The relative amounts of species in a solution are expressed in

units of concentration.

All concentrations are ratios of the solute to either the solvent or

the solution.

But there are many subtle difference in the ways the concentration could be

expression and some of these were shown here.

We can talk about the relative masses of the solute and the solution.

Or the relative number per million or per billion for less concentrated solutions.

So PPM and PPB, which are used a lot in environmentals sampling, for

example, are for, are very useful for

less concentrated solutions where there's not as much solute in the solution.

We could also look at the amount of solute present.

Relative to the volume of the solute and the volume of the solution.

One that gets used a lot in chemistry is molarity.

And that's the one I'll focus on the, the most but there's others and

I'm not even showing all the different ways you can express concentration here.

You might say things like mole fraction, and that is just.

What is the relative number of moles of solute to solution?

There's also molality, which gets used a lot in engineering.

And that is the moles of solute per kilogram of solvent in that case.

So, for all of these, the solute is in the numerator.

And there's a fraction where either the solvent or

the solution is in the denominator.

Let's concentrate on the molarity.

That's the one that I use the most in chemistry.

Molarity is defined as the moles of solute per liter of solution.

So, I'm using N here as the symbol for the number of moles and V as the symbol for

the volume.

For example, I could write something like this.

In square brackets I could write a chloride ion.

Now, if you're trying to show molar concentrations you have to use.

Square brackets, like that.

You can't use parenthesis and you can't use little squirrely brackets.

Those are not correct for showing molarity or

molar concentration, another thing we call molarity.

So you have to use these very clearly drawn square brackets if you're trying to

show that something's got a specific molar concentration.

And here this expression that I've written can be read as the chloride ion

concentration is 1 molar.

Sometimes the solution is more dilute.

This happens a lot, for example, in biological testing or

in environmental testing.

And in that case, it might be more convenient to express the concentration in

millimolar, micromolar or nanomolar for example.

And in those cases, the concentration is more dilute, so we need to use these

prefixes which mean, 10 to the minus 3, 10 to the minus 6, and 10 to the minus 9.

We can convert those to regular molarity, using these conversion factors over here.

For example, what if I had a concentration of formaldehyde.

Formaldehyde CH2O.

That's in water to concentration of 1.64 times10 to the minus for molar.

And that's its concentration.

I could say that it's the 164 micromolar concentration.

And then I don't have to use a scientific notation,

I don't have to write quite so much.

It gets a little simpler to express that micromolar.

So that's why you might see it expressed that way.

Now that we've determined all sorts of ways we can

express the concentration of our solution,

we need to think about the differences in solubility between different compounds.

For example, sodium chloride dissolves really readily in water.

We can dissolve a lot of salt in water.

But there are other chemicals,

such as sand, that don't dissolve very well in water.

If we took a container of water and

dumped in a teaspoon of sand, it wouldn't look like any of it dissolves at all.

Where as if we took a container of water and we dumped in a teaspoon of salt,

as long as we had enough water we could stir it up and the salt would dissolve.

So there's different.

Amount of solute that can dissolve in solvent.

There's limits to the amount of a solute that will dissolve in a given type of

solvent, and that depends on the chemical composition of the solute,

the chemical composition of the solvent, the temperature

of the reaction mixture, and even the pressure of the reaction mixture.

For the most part we're going to assume that we're at room temperature and

normal atmospheric pressure at sea level.

So not everything dissolves in water to the same extent.

We us the example of putting sand in water and seeing that it doesn't dissolve

at least to the naked eye it doesn't look like it dissolves at all.

So we use different words to express how much has

solute has dissolved in the solvent?

If there's a limit to how much goes in solution,

we might want to use different moler concentrations as a guide.

And these are guides I got from my good friend Denis Wurtz's book.

He says if we can make a solution of solute that is greater than 0.10 moler

in concentration, then we can fairly say, with fair certainty, that,

that compound is soluble.

We would describe it as being soluble.

So we can say that sodium chloride is soluble in water.

Some species, particularly ionic compounds, dissolve but

have a limit to their solubility.

And this is another, this is an arbitrary range that was set in a particular book.

But, in that book they said if the molar concentration that will dissolve,

the maximum amount that will dissolve before solid chunks start to form on

the bottom of the container.

Is between 0.01 molar and

0.1 molar then we say that that compound is moderately soluble in that solvent.

And if the amount of solute that were dissolved before it just continues to

fall to the bottom of the container as a solid is less than 0.1

molar in concentration then we say.

That particular compound is insoluble in that particular solvent.

Now we might say something is insoluble and

there's still a very, very tiny bit that does dissolve but

as far as we can tell with our naked eye most of it is not dissolving and

if something's not dissolving well in the solvent we say it's insoluble.

This brings up another word that we need to define and that word is saturated.

Saturated is when we have added enough solute that we have

that we have reached the limit of how much will dissolve in the solvent.

So.

If we added any more of the solute, let's say we're putting sodium chloride into

water, we add it, stir it, and it dissolves.

We add some more and stir, it dissolves.

We keep adding sodium chloride and stirring and it keeps dissolving.

And eventually, the next little bit that we add can no longer dissolve.

There is already so much sodium chloride in the water that the next little bit just

falls to the bottom as a solid.

It doesn't look like it dissolves.

This happens, this is easy to observe in sugar.

I have lots of friends being in the South who like to drink sweetened iced tea and.

Often they put enough sugar in the iced tea that the sugar does not all

completely dissolve.

And its there's some sugar flurry at the bottom.

Sitting at the bottom.

Sometime people observe this in lemonade as well.

So something saturated when this exactly at that point where

the next little tiny bit that you add will no longer dissolve.

Everything you added up till now is dissolved.

And then the next little bit will not dissolve and

we say that that solution is saturated.

One way you can make a completely saturated solution is put in

more than will dissolve, stir it readily and then run it through a filter so

that the solid is trapped and the liquid that

contains the solute passes through; and that gives you a saturated solution.

Thanks for watching this video on introduction to solution.

In the next video we'll talk about electrolytes.

Let's check that out.

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