4.01 Buffers and the Common Ion Effect

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From the course by University of Kentucky

Advanced Chemistry

354 ratings

University of Kentucky

Advanced Chemistry

354 ratings

A chemistry course to cover selected topics covered in advanced high school chemistry courses, correlating to the standard topics as established by the American Chemical Society. Prerequisites: Students should have a background in basic chemistry including nomenclature, reactions, stoichiometry, molarity and thermochemistry.

From the lesson

Aqueous Equilibria

This unit continues and expands on the theme of equlibria. You will examine buffers, acid/base titrations and the equilibria of insoluble salts.

4.01 Buffers and the Common Ion Effect10:01

4.02 pH of Buffer Solutions14:04

4.02a Aqueous Equilibria Worked Example 13:58

4.03 Buffer Action25:21

4.03a Aqueous Equilibria Worked Example 25:22

4.03b Aqueous Equilibria Worked Example 35:11

4.04 Buffer: Preparation and Capacity8:45

4.05 Strong Acid - Strong Base Titration17:20

4.06 Titrations Involving Either a Weak Acid or a Weak Base36:45

4.06 Part 1 Aqueous Equilibria Worked Example3:46

4.06 Part 1.a - Aqueous Equilibria Worked Example4:06

4.06 Part 1.b - Aqueous Equilibria Worked Example4:55

4.06 Part 1.c - Aqueous Equilibria Worked Example4:21

4.06 Part 1.f Aqueous Equilibria Worked Example8:01

4.06 Part 1.h Aqueous Equilibria Worked Example3:57

4.06 Part 2.a Aqueous Equilibria Worked Example8:25

4.06 Part 2.b Aqueous Equilibria Worked Example5:53

4.07 Polyprotic Acid Titrations5:53

4.08 Indicators10:19

4.09 Solubility Equilibria4:50

4.10 Molar Solubility and the Solubility Product10:37

4.10.a Aqueous Equilibria Worked Example2:20

4.11 Molar Solubility and the Common Ion Effect7:20

4.11 2.b - Aqueous Equilibria Worked Example4:42

4.11 2.c Aqueous Equilibria Worked Example3:45

4.12 The Effect of pH on Solubility7:07

4.12.b Aqueous Equilibria Worked Example3:42

4.12c Aqueous Equilibria Worked Example3:30

4.13 Precipitation Reaction and Selective Precipitation15:03

4.13a Aqueous Equilibria Worked Example9:05

Meet the Instructors

Dr. Allison Soult
Lecturer
Chemistry
Dr. Kim Woodrum
Sr. Lecturer
Chemistry

You have been through several units in which you have

dealt with equilibria.

First, we introduced the concept of equilibria

and then we applied it in the environment of

acid solutions or base solutions

or salt solutions.

That have acidic and basic properties.

For there we are going to go into this chapter

and this unit where we are going to be looking at

other environments in which aqueous and this unit where we are going to be looking at

other environments in which aqueous

solutions play a big role.

This unit is divided into two main

categories. We have got

homogeneous equilibrium

and that will be our first homogeneous equilibrium

and that will be our first

8 learning objectives and that will be our first

8 learning objectives

and then we will look at heterogeneous equilibrium.

We will see environments where a solid

will form from solution.

Our first learning objective

is to learn about buffers.

We need to identify what a buffer is

and the necessary

component that make up a buffer.

This is what a buffer is.

It is a solution which resists change to pH. This is what a buffer is.

It is a solution which resists change to pH.

So you create this solution

and if you add an acid or you add

base to the solution, the

pH will not change very much.

It will resist the change.

This is a recipe in a sense.

What is in the solution that makes it a buffer?

It needs to have a week acid

and its salt. It needs to have a week acid

and its salt.

Or a weak base and its salt.

But lets look at what I mean by that.

If it is a weak acid and it is But lets look at what I mean by that.

If it is a weak acid and it is

salt, the salt is

[take a weak acid, and I am going to take a weak acid

and write it in this space right here [take a weak acid, and I am going to take a weak acid

and write it in this space right here

HF for example.] and write it in this space right here

HF for example.]

I take take weak acid

and I were to

think about what its

conjugate base is, that would be F-. think about what its

conjugate base is, that would be F-.

And then if I were to create the salt from that

I would need to add and ion with it.

So a counter ion, like cation

NaF

would be a salt

containing that weak base.

So a buffer consists of a weak acid.

and the salt of its conjugate base.

It could also consist of a weak base

and its salts. So lets think of a weak base.

Ammonia is a weak base.

If we were to think about its conjugate acid

that would be NH4+.

Well, I cannot go and obtain some NH4+

off a stockroom shelf I need to

obtain a salt that contains that ion

so I put a counter ion with it.

A negatively charged ion for example

and we would have A negatively charged ion for example

and we would have

[and I should not say for example, it has to be negatively charged]

chloride for example there would be a salt

so we have a weak base.

And its conjugate acid

salt, that would be a buffer.

So here is some examples besides the ones I already gave you.

If I took acetic acid

and that is what this substance is acetic acid.

And I were to consider its conjugate base

the conjugate base would be removing that proton right there.

and that would become

anion, which is its conjugate base

and the salt, letsput a counterion with that

we put a sodium for example, so that is a weak acid and the salt, letsput a counterion with that

we put a sodium for example, so that is a weak acid

and its conjugate base salt.

We often write it with this

slash like this, where you write the base first and the acid second

that is a standard way of writing slash like this, where you write the base first and the acid second

that is a standard way of writing

a buffered solution.

I have already given you that example in the previous slide

and it to, would you write the base first and the acid second

it would be a standard way of writing and it to, would you write the base first and the acid second

it would be a standard way of writing

that buffered solution. it would be a standard way of writing

that buffered solution.

Here is an example of something that is not a buffer

but often student think that is might be.

If I had a solution containing HCl

and sodium chloride. Well, this is an acid.

This is its conjugate base.

this is its salt.

So it is an acid with a conjugate base salt

but this is not a weak acid.

It is a strong acid.

And as a strong acid

it cannot be used to make a buffer.

Lets have you consider it cannot be used to make a buffer.

Lets have you consider

this solution. We have a solution containing

this substance this solution. We have a solution containing

this substance

and this substance and it is not a buffer.

It has a weak base in it, it has a salt.

containing an acid. It has a weak base in it, it has a salt.

containing an acid.

Why is that not a buffer?

Think about it and choice your answer.

Well if you said, number 3 then you are correct.

There are not conjugates of each other Well if you said, number 3 then you are correct.

There are not conjugates of each other

while CH_3NH_2

methylamine

is a weak base.

And while this is a

acid, conjugate acid of a weak base.

It is not the conjugate acid

of this weak base.

Therefore, this cannot be a buffer.

Now, this is a buffer

and this one throws students very often

we have sodium hydroxide, well what is sodium hydroxide?

It is a strong base.

and we have HF

it is a weak acid. and we have HF

it is a weak acid.

It certainly does not constitute

the components of a buffer It certainly does not constitute

the components of a buffer

but acid and bases react with each other

lets consider the reaction that will take place.

The OH- of the base

will react with the acid.

A base reacts with an acid will react with the acid.

A base reacts with an acid

to form F- and water.

So if we look at the amount we have

we have 10 ml of .1 molar

and 25 ml of .1 molar.

We have a little bit of this comparatively

a lot of this.

So what will happen? Well this

OH- will react with the HF. So what will happen? Well this

OH- will react with the HF.

And produce some of this.

It will all go away.

And we will have a little less of this

and some of that, and what will we have?

We have in solution and some of that, and what will we have?

We have in solution

a weak acid

and a conjugate base.

Therefore we have a buffer.

So when you have reaction between a strong and a weak

and undergoes a neutralization reaction

if the amounts are just right, it turns into a buffer.

Now we will talk about this a whole lot

more when we get into titrations. Now we will talk about this a whole lot

more when we get into titrations.

So this is just a reminder, when we get to titrations

we will talk about creating buffer solutions in the So this is just a reminder, when we get to titrations

we will talk about creating buffer solutions in the

course of that, and that won't be

a few learning objectives down the road.

So buffers are what we call a

common ion situation. So buffers are what we call a

common ion situation.

What does that mean? Well a common ion common ion situation.

What does that mean? Well a common ion

is when you have dissolved two substances

into a solution

and the two

things you are dissolving, both contain

the same ion. That would be a common ion solution.

So if we were to take HF, an acid,

and dissolved in there NaF.

So we have placed into a flask both of these things

what would be the common ion?

What do they have in common? what would be the common ion?

What do they have in common?

Well, they both have F- in them.

So we know that when we place HF in water

it, to some extend because it is weak,

will produce some of the F- ion,

in solution.

We, know that when you dissolve sodium fluoride

it 100% dissolves

it is a nice soluble salt. 100% break apart

and we end up with some F-. it is a nice soluble salt. 100% break apart

and we end up with some F-.

So this is a common ion situation.

So lets thing about this, this is a question for you to answer.

If you have this reaction happening

and you add some sodium fluoride,

in which there is some F-, and you add some sodium fluoride,

in which there is some F-,

the question is: What does that added

F-, that we see right here,

what does that added F- do

to this equilibrium?

Does it shift it to the right?

And when you shift this reaction here

to the right, we would say it increases its ionization. And when you shift this reaction here

to the right, we would say it increases its ionization.

Or does it shift it to the left? to the right, we would say it increases its ionization.

Or does it shift it to the left?

And if it shifts it to the left

we would say we are decreasing its ionization.

Well, if you said decreasing the ionization you are correct.

You are increasing the amount of this

Le Chatelier's principle says that when you add a product You are increasing the amount of this

Le Chatelier's principle says that when you add a product

the reaction will shift left

to try and use up that added product.

So this shift is known as a common ion effect.

The F- is the common ion So this shift is known as a common ion effect.

The F- is the common ion

shifting it to the left is a common ion effect.

And this is, in a buffer always what happens

when you add

the salt that contains the conjugate base,

for example. You will decrease

the ionization of that acid

and you will have in solution

a fair amount of both

the acid form, and the base form

of that buffer.

So this ends our first learning objective.

We have completed the idea of what a buffer is,

what it does, what its components are.

And then we are going to start going into

calculating the pH with our next learning objective. And then we are going to start going into

calculating the pH with our next learning objective.

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