draft exp 3 preparing buffer solutions

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Introduction

Preparing Buffer SolutionsTable of Contents

1. Origin of the Henderson-Hasselbalch Equation

2. References

3. Contributors

When it comes to buffer solution one of the most common equation is the Henderson-Hasselbalch approimation.!n

important point that must be made about this equation is it"s useful onl# if stoichiometric or initial concentration can be

substituted into the equation for equilibrium concentrations.

Origin of the Henderson-HasselbalchEquationWhere the Henderson-Hasselbalch approimation comes from

HA+H2OH3O++A(1)

$here%

Ais the con&ugate base HAis the $ea' acid

We 'no$ that Kais equal to the products o(er the reactants and% b# definition% H 2O is essentiall# a pure liquid that $e

consider to be equal to one.

Ka=[H3O+][A](2)

)a'e thelogof both sides*

logKa=log([H3O+][A])(3)

logKa=log[H3O+]log[A](4)

+sing the follo$ing t$o relationships*

log[Ka]=pKa(5)

log[H3O+]=pH(6)
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We can simplif# the abo(e equation*

pKa=pHlog[A](7)

,f $e addlog[A]to both sides% $e get the Henderson-Hasselbalch approximation

pH=pKa+log[A](8)

)his approimation is onl# !alid$hen*

1. )he con&ugate base acid falls bet$een the (alues of .1 and 1

2. )he molarit# of the buffers eceeds the (alue of the /ab# a factor of at least 1

Example "

0uppose $e needed to ma'e a buffer solution $ith a pH of 2.11. ,n the first case% $e $ould tr# and find a $ea' acid

$ith a p/a(alue of 2.11. Ho$e(er% at the same time the molarities of the acid and the its salt must be equal to one

another. )his $ill cause the t$o molarities to cancel lea(ing thelog[A]equal tolog(1)$hich is ero.

pH=pKa+log[A]=2.11+log(1)=2.11

)his is a (er# unli'el# scenario% ho$e(er% and #ou $on"t often find #ourself $ith Case 1

Example #

What mass of NaC7H5O2must be dissol(ed in .2 4 of .3 5 HC6H7O2to produce a solution $ith pH 8 9.6:;

SO$%TIO&

HC7H5O2+H20H3O++C7H5O2

Ka=6.3105

Ka=[H3O+][C7H5O2][HC7H5O2]=6.3105

[H3O+]=10pH=104.78=16.6106M[HC7H5O2]=0.30M[C7H5O2]=

[C7H5O2]=Ka[HC7H5O2][H3O+]

1.14M=6.31050.3016.6106

5ass 8 .2 4 1.19 mol C6H7O2- 14 1mol >aC6H7O2 1 mol C6H7O2- 199 g >aC6H7O2 1 mol >aC6H7O28 32.:32

g >aC6H7O2


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u!er_#$luti$ns

Preparing Buffer Solutions

HPLC A Basic Knowledge of Analysis

The pH of the mobile phase (eluent is ad!usted to impro"e component separation and to e#tend the

column life$ This pH ad!ustment should in"ol"e not simply dripping in an acid or al%ali but using buffer

solutions& as much as possible$ 'ood separation reproducibility (stability may not be achie"ed if

buffer solutions are not used$

A buffer solution is prepared as a combination of wea% acids and their salts (sodium salts& etc$ or of

wea% al%alis and their salts$ Common preparation methods include ) dripping an acid (or al%ali intoan a*ueous solution of a salt while measuring the pH with a pH meter and + ma%ing an a*ueous

solution of acid with the same concentration as the salt and mi#ing while measuring the pH with a pH

meter$ Howe"er& if the buffer solution is used as an HPLC mobile phase& e"en small errors in pH can

lead to problems with separation reproducibility$ Therefore& it is important to diligently inspect and

calibrate any pH meter that is used$ This page introduces a method that does not rely on a pH meter$

The method in"ol"es weighing theoretically calculated fi#ed *uantities of a salt and acid (or al%ali as

shown in the table below$ Consider the important points below$

,enoting Buffer Solutions

A buffer solution denoted& -).. m/ phosphoric acid (sodium buffer solution pH 0 +$)&- for e#ample&

contains phosphoric acid as the acid& sodium as the counterion& ).. m/ total concentration of the

phosphoric acid group& and a guaranteed buffer solution pH of +$)$

/a#imum Buffer Action Close to the Acid (or Al%ali pKa

1hen an acetic acid (sodium buffer solution is prepared from )) acetic acid and sodium acetate& for

e#ample& the buffer solution pH is appro#imately 2$3 (near the acetic acid pKa& and this is where the

ma#imum buffer action can be obtained$

Buffer Capacity 4ncreases as Concentration 4ncreases

The buffer capacity of an acetic acid (sodium buffer solution is larger at ).. m/ concentration than

at ). m/& for e#ample$ Howe"er& precipitation occurs more readily at higher concentrations$
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Beware of Salt Solubility and Precipitation

The salt solubility depends on the type of salt& such as potassium salt or sodium salt$ Salts precipitate

out more readily when an organic sol"ent is mi#ed in$

4n addition& a"oid using buffer solutions based on organic acids (carbo#ylic acid as much as possible

for highly sensiti"e analysis at short 56 wa"elengths$ Consider the "arious analytical conditions and

use an appropriate buffer solution& such as an organic acid with a hydro#yl group at the 7 position

(see Supplement to restrict the effects of metal impurity ions$ (8$/a&9$:g

http://www.shimad%u.c$m/an/hplc/supp$rt/li&/lctalk/'(/'(la&.html

A &u!er is a s$luti$n $* weak acid and

c$n+u"ate &ase $r weak &ase and c$n+u"ateacid used t$ resist pH chan"e with added

s$lute.

LEARNING OBJECTIVE

Describe the properties of a buffer solution.

KEY POINTS

u!ers$luti$nsare resistant topHchange because of the presence ofan e,uili&riumbetween the acid(HA) and its c$n+u"ate &ase(A-).

When some str$n" acidis added to a buffer, the equilibrium is shifted to theleft, and the hydrogeni$nc$ncentrati$nincreases by less than epected for the amount ofstrong acid added.

!uffer solutions are necessary in biology for "eeping the correct pH for proteinstow$rk.

!uffers can be prepared in multiple ways by creating a solution of an acid and its

con#ugate base.

TERMS

a,ue$us
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$onsisting mostly of water.

e,uili&rium

%he state of a reaction in which the rates of the forward (reactant to product) and

re&erse (product to reactant) reactions are the same.

pKa

A quantitati&e measure of the strength of an acid in solution' a wea" acid has a pa

&alue in the approimate range * to +* in water and a strong acid has a pa &alue of

less than about *.

-e"ister *$r - t$ st$p seein" ads

FULL TEXT

Buffers

A buffer is an a,ue$ussolution containing a weak acidand its con#ugate

base or a weak &aseand its c$n+u"ate acid. A buffers pH changes &ery

little when a small amount of strong acid or base is added to it. t is used to

pre&ent any change in the pH of a solution, regardless of s$lute. !uffer

solutions are used as a means of "eeping pH at a nearly constant &alue in a

wide &ariety of chemical applications. or eample, blood in the human body

is a buffer solution.

!uffer solutions are resistant to pH change because of the presence of an

equilibrium between the acid (HA) and its con#ugate base (A0). %he &alanced

e,uati$nfor this reaction is/

HA H0 0 A 1
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When some strong acid (more H1) is added to an equilibrium mi2tureof the

wea" acid and its con#ugate base, the equilibrium is shifted to the left, in

accordance with 2e $hateliers principle. %his causes the hydrogen ion (H 1)

concentration to increase by less than the amount epected for the quantity of

strong acid added. 3imilarly, if astr$n" &aseis added to the miture, the

hydrogen ion concentration decreases by less than the amount epected for

the quantity of base added. %his is because the reaction shifts to the right to

accommodate for the loss of H1in the reaction with the base.

!uffer solutions are necessary in a wide range of applications. n biology, they

are necessary for "eeping the correct pH for proteins to wor"' if the pH mo&es

outside of a narrow range, the proteins stop wor"ing and can fall apart. A

buffer of carbonic acid (H3$4') and bicarbonate (H$4'4) is needed in

blood plasmato maintain a pH between 5.67 and 5.87. ndustrially, buffer

solutions are used in fermentation processes and in setting the correct

conditions for dyes used in coloring fabrics.

Preparing a Buffer Solution

%here are a couple of ways to prepare a buffer solution of a specific pH. n the

first method, prepare a solution with an acid and its con#ugate base by

dissol&ing the acid form of the buffer in about 9:; of the v$lumeof water

required to obtain the final solution &olume. %hen, measure the pH of the
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solution using a pH probe. %he pH can be ad#usted up to the desired &alue

using a strong base li"e A-?@ >HA?)

where pH is the concentration of >H0?, pKais the acid diss$ciati$nconstant,

and >A-? and >HA? are concentrations of the con#ugate base and starting acid.
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te2t&$$k/acid0&ase0e,uili&ria056/&u!er0s$luti$ns0557/preparin"0a0&u!er0s$luti$n0

with0a0speci8c0ph097907365/

uesti$ns

5a

THEORIES OF ACIDS AND BASES

This page describes the Arrhenius, Bronsted-Lowry, and Lewis

theories of acids and bases, and explains the relationshipsbetween them. It also explains the concept of a conjugate pair - an

acid and its conjugate base, or a base and its conjugate acid.

Note:Current UK A' level syllabuses concentrate on the Bronsted-

Lowry theory, but you should also be aware of Lewis acids and bases.

The Arrhenius theory is of historical interest only, and you are unlikely

to need it unless you are doing some work on the development of

ideas in chemistry.

The Arrhenius Theory of acids and bases

The theory

Acids are substances which produce hydrogen ions in

solution.

Bases are substances which produce hydroxide ions in

solution.

Neutralization happens because hydrogen ions and hydroxide ions

react to produce water.
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Limitations of the theory

Hydrochloric acid is neutralised by both sodium hydroxide solution

and ammonia solution. In both cases, you get a colourless solution

which you can crystallise to get a white salt - either sodium chloride

or ammonium chloride.

These are clearly very similar reactions. The full equations are:

In the sodium hydroxide case, hydrogen ions from the acid are

reacting with hydroxide ions from the sodium hydroxide - in line with

the Arrhenius theory.

However, in the ammonia case, there don't appear to be any

hydroxide ions!

You can get around this by saying that the ammonia reacts with the

water it is dissolved in to produce ammonium ions and hydroxideions:

This is a reversible reaction, and in a typical dilute ammonia

solution, about 99% of the ammonia remains as ammonia

molecules. Nevertheless, there are hydroxide ions there, and we

can squeeze this into the Arrhenius theory.

However, this same reaction also happens between ammonia gas

and hydrogen chloride gas.

In this case, there aren't any hydrogen ions or hydroxide ions in

solution - because there isn't any solution. The Arrhenius theory


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wouldn't count this as an acid-base reaction, despite the fact that it

is producing the same product as when the two substances were in

solution. That's silly!

The Bronsted-Lowry Theory of acids and bases

The theory

An acid is a proton (hydrogen ion) donor.

A base is a proton (hydrogen ion) acceptor.

The relationship between the Bronsted-Lowry theory and the

Arrhenius theory

The Bronsted-Lowry theory doesn't go against the Arrhenius theory

in any way - it just adds to it.

Hydroxide ions are still bases because they accept hydrogen ions

from acids and form water.

An acid produces hydrogen ions in solution because it reacts withthe water molecules by giving a proton to them.

When hydrogen chloride gas dissolves in water to produce

hydrochloric acid, the hydrogen chloride molecule gives a proton (a

hydrogen ion) to a water molecule. A co-ordinate (dative covalent)

bond is formed between one of the lone pairs on the oxygen and

the hydrogen from the HCl. Hydroxonium ions, H3O+, are produced.


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Note:If you aren't sure aboutco-ordinate bondingyou should follow

this link. Co-ordinate bonds will be mentioned several times over the

course of the rest of this page.

Use the BACK button on your browser to return quickly to this page.

When an acid in solution reacts with a base, what is actually

functioning as the acid is the hydroxonium ion. For example, a

proton is transferred from a hydroxonium ion to a hydroxide ion tomake water.

Showing the electrons, but leaving out the inner ones:

It is important to realise that whenever you talk about hydrogen

ions in solution, H+(aq), what you are actually talking about are

hydroxonium ions.

The hydrogen chloride / ammonia problem

This is no longer a problem using the Bronsted-Lowry theory.

Whether you are talking about the reaction in solution or in the gas

state, ammonia is a base because it accepts a proton (a hydrogen

ion). The hydrogen becomes attached to the lone pair on the

nitrogen of the ammonia via a co-ordinate bond.
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If it is in solution, the ammonia accepts a proton from a

hydroxonium ion:

If the reaction is happening in the gas state, the ammonia accepts

a proton directly from the hydrogen chloride:

Either way, the ammonia acts as a base by accepting a hydrogen

ion from an acid.

Conjugate pairs

When hydrogen chloride dissolves in water, almost 100% of it

reacts with the water to produce hydroxonium ions and chloride

ions. Hydrogen chloride is a strong acid, and we tend to write this

as a one-way reaction:

Note:I am deliberately missing state symbols off this and the next

equation in order to concentrate on the bits that matter.

You will find more aboutstrong and weak acidson another page in this

section.
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In fact, the reaction between HCl and water is reversible, but only to

a very minor extent. In order to generalise, consider an acid HA,

and think of the reaction as being reversible.

Thinking about theforward reaction:

The HA is an acid because it is donating a proton (hydrogen

ion) to the water.

The water is a base because it is accepting a proton from

the HA.

But there is also aback reactionbetween the hydroxonium ion and

the A-ion:

The H3O+is an acid because it is donating a proton

(hydrogen ion) to the A-ion.

The A-ion is a base because it is accepting a proton from

the H3O+.

The reversible reaction containstwoacids andtwobases. We think

of them in pairs, calledconjugate pairs.

When the acid, HA, loses a proton it forms a base, A-. When the

base, A-, accepts a proton back again, it obviously refoms the acid,HA. These two are a conjugate pair.

Members of a conjugate pair differ from each other by the presence

or absence of the transferable hydrogen ion.

If you are thinking about HA as the acid, then A-is its conjugate


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base.

If you are thinking about A-as the base, then HA is its conjugate

acid.

The water and the hydroxonium ion are also a conjugate pair.

Thinking of the water as a base, the hydroxonium ion is its

conjugate acid because it has the extra hydrogen ion which it can

give away again.

Thinking about the hydroxonium ion as an acid, then water is its

conjugate base. The water can accept a hydrogen ion back again

to reform the hydroxonium ion.

A second example of conjugate pairs

This is the reaction between ammonia and water that we looked at

earlier:

Think first about the forward reaction. Ammonia is a base because

it is accepting hydrogen ions from the water. The ammonium ion is

its conjugate acid - it can release that hydrogen ion again to reform

the ammonia.

The water is acting as an acid, and its conjugate base is the

hydroxide ion. The hydroxide ion can accept a hydrogen ion to

reform the water.

Looking at it from the other side, the ammonium ion is an acid, and

ammonia is its conjugate base. The hydroxide ion is a base and

water is its conjugate acid.


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Amphoteric substances

You may possibly have noticed (although probably not!) that in one

of the last two examples, water was acting as a base, whereas in

the other one it was acting as an acid.

A substance which can act as either an acid or a base is described

as beingamphoteric.

Note:You might also come across the termamphiproticin thiscontext. The two words are related and easily confused.

Anamphiproticsubstance is one which can both donate hydrogen ions

(protons) and also accept them. Water is a good example of such a

compound. The water acts as both an acid (donating hydrogen ions)

and as a base (by accepting them). The "protic" part of the word refers

to the hydrogen ions (protons) either being donated or accepted. Other

examples of amphiprotic compounds are amino acids, and ions like

HSO4-(which can lose a hydrogen ion to form sulphate ions or accept

one to form sulphuric acid).

But as well as being amphiprotic, these compounds are

alsoamphoteric. Amphoteric means that they have reactions as both

acids and bases. So what is the difference between the two terms?

All amphiprotic substances are also amphoteric - but the reverse isn't

true. There are amphoteric substances which don't either donate or

accept hydrogen ions when they act as acids or bases. There is a

whole new definition of acid-base behaviour that you are just about to

meet (the Lewis theory) which doesn't necessarily involve hydrogen

ions at all.

A Lewis acid is an electron pair acceptor; a Lewis base is an electron

pair donor (see below).

Some metal oxides (like aluminium oxide) are amphoteric - they react

both as acids and bases. For example, they react as bases because

the oxide ions accept hydrogen ions to make water. That's not a


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problem as far as the definition of amphiprotic is concerned - but the

reaction as an acid is. The aluminium oxide doesn't contain any

hydrogen ions to donate! But aluminium oxide reacts with bases like

sodium hydroxide solution to form complex aluminate ions.

You can think of lone pairs on hydroxide ions as forming dative

covalent (coordinate) bonds with empty orbitals in the aluminium ions.

The aluminium ions are accepting lone pairs (acting as a Lewis acid).

So aluminium oxide can act as both an acid and a base - and so is

amphoteric. But itisn'tamphiprotic becausebothof the acid reaction

and the base reaction don't involve hydrogen ions.

I have gone through 40-odd years of teaching (in the lab, and via

books and the internet) without once using the term amphiprotic! I

simply don't see the point of it. The term amphoteric takes in all the

cases of substances functioning as both acids and bases without

exception. The term amphiprotic can only be used where both of these

functions involve transference of hydrogen ions - in other words, it can

only be used if you are limited to talking about the Bronsted-Lowry

theory. Personally, I would stick to the older, more useful, term

"amphoteric" unless your syllabus demands that you use the word

"amphiprotic".

The Lewis Theory of acids and bases

This theory extends well beyond the things you normally think of as

acids and bases.

The theory

An acid is an electron pair acceptor.

A base is an electron pair donor.

The relationship between the Lewis theory and the Bronsted-

Lowry theory


17/94

Lewis bases

It is easiest to see the relationship by looking at exactly what

Bronsted-Lowry bases do when they accept hydrogen ions. Three

Bronsted-Lowry bases we've looked at are hydroxide ions,

ammonia and water, and they are typical of all the rest.

The Bronsted-Lowry theory says that they are acting as bases

because they are combining with hydrogen ions. The reason they

are combining with hydrogen ions is that they have lone pairs ofelectrons - which is what the Lewis theory says. The two are

entirely consistent.

So how does this extend the concept of a base? At the moment it

doesn't - it just looks at it from a different angle.


18/94

But what about other similar reactions of ammonia or water, for

example? On the Lewis theory,anyreaction in which the ammonia

or water used their lone pairs of electrons to form a co-ordinate

bond would be counted as them acting as a base.

Here is a reaction which you will find talked about on the page

dealing with co-ordinate bonding. Ammonia reacts with BF3by

using its lone pair to form a co-ordinate bond with the empty orbital

on the boron.

As far as the ammonia is concerned, it is behaving exactly the

same as when it reacts with a hydrogen ion - it is using its lone pair

to form a co-ordinate bond. If you are going to describe it as a base

in one case, it makes sense to describe it as one in the other caseas well.

Note:If you haven't already read the page aboutco-ordinate

bondingyou should do so now. You will find an important example of

water acting as a Lewis base as well as this example - although the

termLewis baseisn't used on that page.


Lewis acids

Lewis acids are electron pair acceptors. In the above example, the
http://www.chemguide.co.uk/atoms/bonding/dative.html#tophttp://www.chemguide.co.uk/atoms/bonding/dative.html#tophttp://www.chemguide.co.uk/atoms/bonding/dative.html#tophttp://www.chemguide.co.uk/atoms/bonding/dative.html#tophttp://www.chemguide.co.uk/atoms/bonding/dative.html#top


19/94

BF3is acting as the Lewis acid by accepting the nitrogen's lone

pair. On the Bronsted-Lowry theory, the BF3has nothing remotely

acidic about it.

This is an extension of the termacidwell beyond any common use.

What about more obviously acid-base reactions - like, for example,

the reaction between ammonia and hydrogen chloride gas?

Whatexactlyis accepting the lone pair of electrons on the nitrogen.

Textbooks often write this as if the ammonia is donating its lone

pair to a hydrogen ion - a simple proton with no electrons around it.

That is misleading! You don't usually get free hydrogen ions in

chemical systems. They are so reactive that they are always

attached to something else. There aren't any uncombined hydrogen

ions in HCl.

There isn't an empty orbital anywhere on the HCl which can accept

a pair of electrons. Why, then, is the HCl a Lewis acid?

Chlorine is more electronegative than hydrogen, and that means

that the hydrogen chloride will be a polar molecule. The electrons in

the hydrogen-chlorine bond will be attracted towards the chlorine

end, leaving the hydrogen slightly positive and the chlorine slightly

negative.

Note:If you aren't sure aboutelectronegativity and bond polarityit

might be useful to follow this link.

http://www.chemguide.co.uk/atoms/bonding/electroneg.html#tophttp://www.chemguide.co.uk/atoms/bonding/electroneg.html#tophttp://www.chemguide.co.uk/atoms/bonding/electroneg.html#top


20/94

The lone pair on the nitrogen of an ammonia molecule is attracted

to the slightly positive hydrogen atom in the HCl. As it approaches

it, the electrons in the hydrogen-chlorine bond are repelled still

further towards the chlorine.

Eventually, a co-ordinate bond is formed between the nitrogen and

the hydrogen, and the chlorine breaks away as a chloride ion.

This is best shown using the "curly arrow" notation commonly used

in organic reaction mechanisms.

Note:If you aren't happy about the use ofcurly arrowsto show

movements of electron pairs, you should follow this link.


The whole HCl molecule is acting as a Lewis acid. It is accepting a

pair of electrons from the ammonia, and in the process it breaks

up.Lewis acids don't necessarily have to have an existing empty

orbital.

A final comment on Lewis acids and bases

If you are a UK A' level student, you might occasionally come

across the termsLewis acidandLewis basein textbooks or other

sources. All you need to remember is:
http://www.chemguide.co.uk/basicorg/conventions/curlies.html#tophttp://www.chemguide.co.uk/basicorg/conventions/curlies.html#tophttp://www.chemguide.co.uk/basicorg/conventions/curlies.html#top


21/94

A Lewis acid is an electron pair acceptor.

A Lewis base is an electron pair donor.

Note:Remember this by thinking of ammonia acting as a base. Most

people at this level are familiar with the reactive lone pair on the

nitrogen accepting hydrogen ions. Ammonia is basic because of its

lone pair. That means that bases must have lone pairs to donate.

Acids are the opposite.

?or all general purposes% stic' $ith the @ronsted-4o$r#

theor#.

http://www.chem"uide.c$.uk/physical/acid&asee,ia/the$ries.html

Br'nsted Concept of (cids andBasesTable of Contents

1. @rAnsted-4o$er# Befinition

2. !cids are roton Bonors and @ases are roton !cceptors3. Duestions

9. !ns$ers

7. Outside 4in's

. 0ources

6. Contributors

,n 1F23% chemists Gohannes @rAnsted and 5artin 4o$r# independentl# de(eloped definitions of acids and bases

based on compounds abilities to either donate or accept protons . @rAnsted and ).5. 4o$r# independentl# de(eloped the theor# of proton donors and proton acceptors in acid-

base reactions% coincidentall# in the same region and during the same #ear. )he !rrhenius theor# $here acids and

bases are defined b# $hether the molecule contains h#drogen and h#droide ion is too limiting. )he main effect of the
http://www.chemguide.co.uk/physical/acidbaseeqia/theories.htmlhttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Br.C3.B8nsted-Lowery_Definitionhttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Acids_are_Proton_Donors_and_Bases_are_Proton_Acceptorshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Questionshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Answershttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Outside_Linkshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Sourceshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Contributorshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Arrhenius_Concept_of_Acids_and_Baseshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Arrhenius_Concept_of_Acids_and_Baseshttp://www.chemguide.co.uk/physical/acidbaseeqia/theories.htmlhttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Br.C3.B8nsted-Lowery_Definitionhttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Acids_are_Proton_Donors_and_Bases_are_Proton_Acceptorshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Questionshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Answershttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Outside_Linkshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Sourceshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases#Contributorshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Arrhenius_Concept_of_Acids_and_Bases


22/94

@rAnsted-4o$r# definition is to identif# the proton


23/94

?or a reaction to be in equilibrium a transfer of electrons needs to occur. )he acid $ill gi(e an electron a$a# and the

base $ill recei(e the electron. !cids and @ases that $or' together in this fashion are called a conjugate pairmade up

of conjugate acidsand conjugate bases.

HA+;A+H;+

! stands for an !cidic compound and I stands for a @asic compound

! Bonates H to form HI.

I !ccepts H from ! $hich forms HI

!-becomes con&ugate base of H! and in the re(erse reaction it accepts a H from HI to recreate H! in order

to remain in equilibrium

HIbecomes a con&ugate acid of I and in the re(erse reaction it donates a H to ! -recreating I in order to

remain in equilibrium

,uestions1. Wh# is HAan !cid;2. Wh# is ;a @ase;

3. Ho$ can !-be a base $hen H! $as and !cid;

9. Ho$ can HIbe an acid $hen I used to be a @ase;

7. &o) that )e understand the concept let.s loo/ at an an example )ith actual compounds0

HCl+H2OH3O++Cl

HC4 is the acid because it is donating a proton to H2O

H2O is the base because H2O is accepting a proton from HC4

H3Ois the con&ugate acid because it is donating an acid to C4 turn into it"s con&ugate acid H2O

ClJ is the con&ugate base because it accepts an H from H3O to return to it"s con&ugate acid HClHo$ can H2O be a base; , thought it $as neutral;

(ns)ers1. It has a proton that can be transferred

2. It receives a proton from HA

3. A-is a conjugate base because it is in need of a H in order to remain in equilibrium and return to HA

9. HZ+is a conjugate acid because it needs to donate or give away its proton in order to return to it's previous

state of Z

7. In the @rAnsted-owry !heory what ma"es a compound an element or a base is whether or not it donates or

accepts protons# If the H$% was in a different problem and was instead donating an H rather than accepting an H it

would be an acid&

http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_ases/Acid/r$nsted_C$

ncept_$*_Acids_and_ases
http://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Baseshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Baseshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Baseshttp://chemwiki.ucdavis.edu/Physical_Chemistry/Acids_and_Bases/Acid/Bronsted_Concept_of_Acids_and_Bases


24/94

5&

Expressing Concentrations of Solutions

A complete description of a solution states what the solute is and how much solute isdissolved in a given amount of solvent or solution. he !uantitative relationship

"etween solute and solvent is the concentration of the solution. his concentrationma# "e e$pressed using several different methods% as discussed ne$t.

A. Concentration by Masshe concentration of a solution ma# "e given as the mass of solute in a given amountof solution% as in the following statements& he northern part of the 'acific Oceancontains 3(.) g salt in each * g seawater. he North Atlantic Ocean has a higher

salt concentration% 3,.) g salt* g seawater.

B. Concentration by Percenthe concentration of a solution is often e$pressed as percent concentration "# mass or

percent "# volume of solute in solution. 'ercent "# mass is calculated from the massof solute in a given mass of solution. A (-"#-mass a!ueous solution of sodiumchloride contains ( g sodium chloride and )( g water in each * g solution.

'ercent "# mass /mass of solute

mass of solution0 *

1$ample&

How man# grams of glucose and of water are in ( g of a (.3 "#-mass glucose solution

2olution


25/94

e now that (.3 of the solution is glucose&

he remainder of the ( g is water&

5f "oth solute and solvent are li!uids% the concentration ma# "e e$pressed as percent"# volume. 6oth eth#l alcohol and water are li!uids7 the concentration of alcohol-water solutions is often given as percent "# volume. 8or e$ample% a )( solution of

eth#l alcohol contains )( m9 eth#l alcohol in each * m9 solution.

'ercent "# volume /volume of solute

volume of solution0 *

1$ample&

:u""ing alcohol is an a!ueous solution containing , isoprop#lalcohol "# volume. How would #ou prepare ;( m9 ru""ing alcoholfrom pure isoprop#l alcoholpp"? are encountered more and more fre!uentl# aswe "ecome aware of the effects of su"stances present in trace amounts in water andair% and as we develop instruments sensitive enough to detect su"stances present insuch low concentrations. 5n discussing mass% parts per million means concentration ingrams per *@grams% or micrograms per gram. 5n discussing volume% parts per millionma# mean milliliters per cu"ic meter% or the mi$ed designation of milligrams percu"ic meter. 8or parts per "illion% the general trend is toward the use of micrograms

per liter when discussing water contaminants% micrograms per cu"ic meter for air% andmicrograms per ilogram for soil concentrations.

D. Concentration in !erms of Moleshe concentration of a solution ma# "e stated as molarit# >?% which is the num"er ofmoles of solute per liter of solution or the num"er of millimoles >mmol? >* millimole/ *-3mole? per milliliter of solution.

olarit# >? /moles solute

volume >liter? solution/

millimoles solute

milliliter solution

A @ >sa# Bsi$ molarB? solution of h#drochloric acid contains @ mol h#drochloricacid in * 9 solution.

he molarit# of a solution gives a ratio "etween moles of solute and volume ofsolution. 5t can "e used as a conversion factor "etween these two units in calculationsinvolving solutions. As a conversion factor% it can "e used two wa#s&


27/94

*. olesvolume >9? states the num"er of moles in one liter of solution. hisconversion factor is used in calculating the num"er of moles of solute in agiven volume of solution.

;. olume >9?moles states that one liter contains some num"er of moles of

solution. his conversion factor is used to calculate the volume of a solutionthat contains a given !uantit# of solute.

1$ample&

How man# moles of h#drochloric acid are in ; m9 of .*( HClmore dilute? solution will "e the same as the moles ofsulfuric acid in the portion of the more concentrated solution. e cancalculate the moles of sulfuric acid in the final dilute solution&

his answer gives the moles of acid needed. e can calculate thevolume of 3.;( H;2O4that would contain .@( mol H;2O4.


30/94

his answer gives the volume of concentrated acid that conatins themoles of acid needed for the dilute solution. his volume of 3.;( H;2O4would "e dissolved in 4E m9 >( m9 - ; m9? water toprepare .( 9 of .*3 H;2O4. his pro"lem is diagramed in thefigure.


31/94

1$ample&

hat volume of @.3) sodium chloride contains (*.; mmol sodiumchloridepp"?

Gg g

https://www.chem.wisc.edu/dept8les/"enchem/sstut$rial/


34/94

concentrated& a solution that contains a large proportion of solute relative to

solvent.

icroscopic view of a dilutesolution of li!uid 6r;dissolvedin li!uid water.

icroscopic view of aconcentrated solution of li!uid6r;dissolved in li!uid water.

Semi'&uantitati%e Expressions of Concentration

A solution can "e semi-!uantitativel# descri"ed as

unsaturated& a solution in which more solute will dissolve% or

saturated& a solution in which no more solute will dissolve.

he solubilityof a solute is the amount of solute that will dissolve in a given amountof solvent to produce a saturated solution. 8or e$ample% at oC% we can dissolve ama$imum of 3(., g of solid NaCl in * m9 of water >a saturated solution?. An#additional solid NaCl that we add to the saturated solution simpl# falls to the "ottom

of the container and does not dissolve.

&uantitati%e Expressions of Concentration

here are a num"er of wa#s to e$press the relative amounts of solute and solvent in asolution. hich one we choose to use often depends on convenience. 8or e$ample% itis sometimes easier to measure the volume of a solution rather than the mass of thesolution.

Note that some e$pressions for concentration are temperature-dependent >i.e.% the

concentration of the solution changes as the temperature changes?% whereas others arenot. his is an important consideration for e$periments in which the temperature doesnot remain constant.

!emperature Dependence of Se%eral Concentration Expressions

concentration expression measurements reuired temperature dependent)


35/94

percent composition>"# mass?

mass of solutemass of solution

no

>mass does not change withtemperature?

molarit# moles of solutevolume of solution

yes

>volume changes withtemperature?

molalit# moles of solutemass of solvent

no

>neither mass nor moleschanges with temperature?

mole fraction moles of solutemoles of solvent

no

>moles does not changewith temperature?

Percent Composition (by mass

e can consider percent "# mass >or weight percent% as it is sometimes called? in twowa#s&

he parts of solute per * parts of solution.

he fraction of a solute in a solution multiplied "# *.

e need two pieces of information to calculate the percent "# mass of a solute in asolution&

he mass of the solute in the solution.

he mass of the solution.

=se the following e!uation to calculate percent "# mass&

Molarity

olarit# tells us the num"er of moles of solute in e$actl# one liter of a solution. >Notethat molarit# is spelled with an BrB and is represented "# a capital .?

e need two pieces of information to calculate the molarit# of a solute in a solution&


36/94

he moles of solute present in the solution.

he volume of solution >in liters? containing the solute.

o calculate molarit# we use the e!uation&

Molality

olalit#% m% tells us the num"er of moles of solute dissolved in e$actl# one ilogramof solvent. >Note that molalit# is spelled with two BlBs and represented "# a lowercase m.?

e need two pieces of information to calculate the molalit# of a solute in a solution&

he moles of solute present in the solution.

he mass of solvent >in ilograms? in the solution.

o calculate molalit# we use the e!uation&

Mole *raction

he mole fraction%X% of a component in a solution is the ratio of the num"er of molesof that component to the total num"er of moles of all components in the solution.

o calculate mole fraction% we need to now&

he num"er of moles of each component present in the solution.

he mole fraction of A%XA% in a solution consisting of A% 6% C% ... is calculated usingthe e!uation&


37/94

o calculate the mole fraction of 6%X6% use&

http://www.chem.purdue.edu/"chelp/s$luti$ns/character.html

Different ways of expressing the concentration of solutions

ass Bercentage

%he mass percentage of a component in a gi&en solution is the mass of the componentper +::g of the solution. or e.g., if WAis the mass of the component A, W!is the massof the component ! in a solution. %hen,

Cample/ A +:; solution of sodium chloride in water (by mass) means that +:g ofsodium chloride are present in +::g of the solution.

olume percentage

%his unit is used in case of a liquid dissol&ed in another liquid. %he &olume percentage isdefined as the &olume of the solute per +:: parts by &olume of solution.

or e.g., f Ais the &olume of component A present is sol&olume of the solution.

%hen,

or e.g., a +:; solution of ethanol $*H74H, in water (by &olume) means that +:cm6ofethanol is present in +::cm6of the solution.

3trength of a solution is defined as the amount of the solute in gms, present in one litreof the solution. t is epressed as g2-+.
http://www.chem.purdue.edu/gchelp/solutions/character.htmlhttp://www.chem.purdue.edu/gchelp/solutions/character.html


38/94

athematically,

olarity

olarity of a solution is defined as the number of moles of solute dissol&ed per litre ofsolution.

athematically,

or e.g., f a is the weight of the solute (in gms) present in $$&olume of the solution.

%hen,

olarity is epressed by the symbol . t can also be epressed as,


39/94

Eelationship between molarity and normality

%he molarity and normality of a solution is related to each other as follows/

olality

olality of a solution is defined as the number of moles of solute dissol&ed in +:::g of asol&ent. athematically, it is epressed as

olality is epressed by the symbol m.

olality does not change with temperature.

ormality

n case of ionic compounds li"e $l, $a$46etc. ormality is used in place of molarity.

t is the number of gram formula masses of solute dissol&ed per liter of the solution. t is

denoted by the symbol . athematically it is gi&en as,

ole raction

t is the ratio of number of moles of one component (solute or sol&ent) to the totalnumber of moles of all the components (solute and sol&ent) present in the solution. t isdenoted by the symbol F. 2et us suppose that a solution contains two components A and! and suppose that nAmoles of A and n!moles of ! are present in the solution then,


40/94

Adding eq (i) and (ii) we get

A0 != +

Barts per million (ppm)

When a solute is present in &ery small amounts, its concentration is epressed in partsper million. t is defined as the amount of the solute present in one million parts of thesolution.

t may be noted that the concentration units li"e molarity, mole fraction etc. arepreferred as they in&ol&e the weight of the solute and sol&ent, which is independent oftemperature. !ut units li"e, molarity,


41/94

+efinition of pH pOH p3) p3a p3b

)he pK factorK is defined as the log of the $hate(er quantit# that follo$s the s#mbol. )heKpK is an operator. ,t communicates the instruction to calculate the negati(e log of an#quantit# that follo$s the s#mbol. )he definition of pH in equation form is

pH 8 -logLH1M $here LH1M means the molar concentration of h#dronium ions% 5 8 moles liter

)his allo$s the definition of the follo$ing series of quantities.

pOH 8-logLOH-M

the negati(e log of the h#droide ion molarit#

p/$ 8 -log/$

the negati(e log of the $ater ion product % /$

p/a 8 -log /athe negati(e log of the acid dissociation

constant% /a

p/b 8 -log /bthe negati(e log of the base dissociationconstant% /b


42/94

The relationship pH 4 pOH 5 "6

,n a $ater solution the ion product for $ater is*

7H48 7OH-8 5 3)5 " 9 ":-"6

)a'e the -log of both sides of the equation

- log 7H48 4;- log 7OH-8< 5 - log 7" 9 ":-"68

pH 4 pOH 5 "6

Calculations of pH

?or strong acids li'e HCl the molar concentrations are essentiall# the h#dronium ionconcentration. )hese strong acids can produce solutions $here the pH can be equal to orless than 1% the pH (alue $ould ha(e a (alue from -19.

Example +etermination of pH from 7H=O48

What is the pH of a solution $hose LH3OM 8 1 1-95

pH 8 -logLH3OM

pH 8 - logL1 1-9M

pH 8 - L log 1 log 1-9M

>ote* When #ou multipl# numbers #ou al$a#s !BB their log forms

log 1 is al$a#s ero

log 18 so log 1-98 -9

pH 8 - L log 1 log 1-9M 8 - L


43/94

What is the pH of a solution $hose LH3OM 8 2.7 1 -75

pH 8 -logLH3OM

pH 8 - logL2.7 1 -7M

pH 8 - L log 2.7 log 1-7M

>ote* When #ou multipl# numbers #ou al$a#s !BB their log forms

log 18 so log 1-78 -7

log 2.7 can be determined using a calculator ha(ing thelog function 'e#*

Enter the number in this case 2.7

depress the log 'e#

Read the displa# $hich should be .3F6F for thisproblem

pH 8 - L.3F6F - 7M 8 9.21 or46>?:#

!lternatel# if #ou can enter a number in scientific notation into #our calculator'e# in 2.7 1 -7

depress the log 'e#

Read the displa# $hich should be -9.2 for this problem

5ultipl# b# -1 to get 4 6>?:#


44/94

Example +etermination of pH from 7OH"- 8 using defintion pOH and equation pH 4pOH 5 "6

Calculate the pH of a solution that has a LOH1-M 8 1 1-75

Betermine pOH 8 -logLOH1- M 8 -log L1 1-7M 8 7

+se the relationship pH pOH 8 19

pH 7 8 19

pH 5 "6 -@ 5 A

http://www.(==mainstreet.c$m/acid_&ase/de8niti$ns0ph.html

Calculations in"ol"ing acids and bases

);$)$) State the e#pression for the ionic product constant of water (Kw$

Water equilibrium

1ater is in e*uilibrium with its dissociated ions (hydrogen and hydro#ide$

The e*uilibrium

H+< H==

Can be e#pressed according to the e*uilibrium law

Kc 0

?H=@?@

?H+


45/94

Howe"er& as the concentration of the water effecti"ely remains constant on both sides of the

e*uilibrium then the ?H+)2mol+dm>D

As the concentration of the hydrogen ions e*uals the concentration of the hydro#ide ions (see

note )then the concentration of hydrogen ions in pure water at +C 0 the s*uare root of

the ionic product of water

0 ) # ).>3 mol dm>E

All e*uilibrium constants are temperature dependent (and this one is no e#ception

The dissociation of water molecules into ions is bond brea%ing and is therefore an endothermic

process(energy must be absorbed to brea% the bonds$ :ndothermic processes are fa"oured

by an increase in temperature and so as the temperature rises the e*uilibrium mo"es further

to the right hand side and Kw gets larger$

As Kw gets larger so do the "alues of the hydrogen ion concentration and the hydro#ide ion

concentration$

As pH is a measure of the hydrogen ion concentration (pH 0 >log?H=@ then as the

temperature increases the pH gets lower > i$e$ the water becomes more acidic$

This is calculated in the following section$

);$)$+ ,educe ?H=(a*@ and ?(a*@ for water at different temperatures gi"en Kw "alues$

Variation of Kw with temperature

The e*uilibrium

H+< H==
http://ibchem.com/IB/ibnotes/18.1.htm#n1http://ibchem.com/IB/ibnotes/18.1.htm#n1http://ibchem.com/IB/ibnotes/18.1.htmhttp://ibchem.com/IB/ibnotes/18.1.htmhttp://ibchem.com/IB/ibnotes/18.1.htm#n1http://ibchem.com/IB/ibnotes/18.1.htm#n1http://ibchem.com/IB/ibnotes/18.1.htmhttp://ibchem.com/IB/ibnotes/18.1.htm


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in"ol"es the brea%ing of bonds and is therefore endothermic > energy must be applied to

brea% one of the the H>H bonds to gi"e the ions$ Conse*uently& according to Le Chatelier& an

increase in temperature fa"ours the forward reaction > i$e$ the position of e*uilibrium shifts

towards the right hand side and Kw becomes larger$

Howe"er& as the ratio of hydrogen ions to hydro#ide ions in pure water must remain ))& then

if we %now the "alue of Kw& it is a simple matter to calculate the "alue of either H=orFand

to obtain the concentrations and hence the "alues of pH and p)2mol+dm>D

As$$$

Kw 0 ?H=@?@

and$$$

?H=@ 0?@

Then$$$

Kw 0 ?H=@+

Therefore$$$

?H=@ 0 G Kw

?H=@ 0 G D$ # ).>)2

?H=@ 0+$ # ).>3

pH 0 D$

The p


47/94

Ior e#ample at +C the hydrogen ion concentration of pure water is ) # ). >3 mol dm>E

The logarithm of ) # ).>3mol dm>E0 >3

The negati"e of >3 0 =3

Therefore the pH of pure water at +C is 3

,efinition of p


48/94

As sulphuric acid dissociates ).. according to the e*uation

H+S


49/94

?H=@ 0 G E$D # ).>D

?H=@ 0)$; # ).>E

The pH of this solution is

pH 0 >log )$; # ).>E0 +$3

);$)$ Sol"e problems in"ol"ing solutions of wea% acids and bases using the e#pressions Ka #

Kb 0 Kw& pKa = pKb 0pKw& pH = p


50/94

CHEC


51/94

?CHEC i$e$ the e*uilibrium constant of the products of acid

dissociation di"ided by the acid concentration at e*uilibrium (howe"er the appro#imation that

the acid concentration at e*uilibrium is the same as the original acid concentation is usually

used for con"enience$

Ka is usually a "ery small number (for e#ample )$3; # ). > for ethanoic acid$ 4t is more

con"enient to use the logarithm of this Ka "alue to gi"e number that are handled more easily$

Howe"er ta%ing logs of "ery small number produces a negati"e "alue$ To a"oid this the

negati"e of the logarithm is used and called the pKa "alue$

Hence

>log Ka 0 pKa


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4f we are dealing with bases then Kb again is "ery small and so pKb is used to define base

strength where

>log Kb 0 pKb

As shown in section );$E$D abo"e

Ka # Kb 0?H=@

#?@

Conse*uently at +C

Ka # Kb 0 ) # ).>)2

And

pKa = pKb 0)2

Ka !alue

5sing the typical wea% acid (HA e*uation& this is represented by the e*uilibrium

HA H== A>

Irom which& by the e*uilibrium law

4t may be seen that an increase in the components of the right hand side of the e*uilibrium

will gi"e rise to a greater "alue for Ka$

Hence the stronger the acid the larger the "alue of Ka

pKa !alue

The relationship between pKa and Ka is one of an in"erse log and so the larger the "alue of Ka

the smaller the "alue of pKa$

Hence the stronger the acid the smaller the "alue of pKa

This may be illustrated by some Ka and pKa "alues


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Acid or base Ka pKa acid strength

Trichloroethanoic acid $). # ).>+ )$+

decreasing acid strength

Chloroethanoic acid )$E; # ).>E +$;D

/ethanoic acid )$33 # ).>2 E$3

:thanoic acid )$3; # ).> 2$3

Propanoic acid )$+D # ).> 2$.

Carbonic acid E$; # ).>3 D$2.

1ater )$.. # ).>3 3$..

Ammonia $+D # ).>). $+

/ethylamine +$+2 # ).>)) ).$D

emember that Ka = Kb 0 Kw

And so& pKb 0 )2 > pKa for the bases

:#ample Calculate the pH of .&+/ ethanoic acid (pKa 0 2$3

Ior the e*uilibrium CHEC


54/94

and ?H=@ 0 ?CHEC# ?CHEC# .$+

Therefore ?H=@ 0 +$)) # ).>E

pH 0 >log ?H=@

Therefore pH 0 +$D;

http://i&chem.c$m/>/i&n$tes/5(.5.htm

3

u!er #$luti$ns


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he "lood is a natural "uffer% and so are other "od# fluids and plant fluids due to

mi$tures of wea acids and "ases present in them.

On this page% we e$plore the reasons wh# the pH of "uffer solutions resists to change.

6uffer solutions are re!uired for man# chemical e$periments. he# are also useful to

standardiIe pH meters. hus% there are man# suppliers of "uffer solutions.

ncpH u!ers.

#ens$re2C$l$r c$ded &u!ers.

here are also computer programs availa"le to help design and mae "uffer solutions on the

internet. 8or e$ample&

u!er Makerusin" the Henders$n0Hassel&alch e,uati$n.

#># #cienti8c #$*tware&u!er maker.

Titration o a Weak Acid by a Strong Base$u have investi"ated h$w the pH varies in a str$n"0acid and str$n"0&ase


56/94

x2+ Kax- CaKa= 0

-Ka+ (Ka2+ 4 CaKa)

1/2

x= ---------------------

2

pH = -log(x)

Discussion

he method has "een full# discussed in ea acids and "asese!uili"rium. 2#m"ols are

used here% "ut appro$imations ma# "e applied to numerical pro"lems.

Example +.

!et us "ake a bufer solution by "i#ing Va"! o acid HA and Vs"!

o its salt $aA% &or si"plicity' let us assu"e both the acid and the

salt solutions hae the sa"e concentration CM% What is the pH o

the so prepared bufer solution? The acid dissociation constant is Ka%

Solution

After mi$ing% the concentrations Caand Csof the acid HA and its salt NaA respectivel#

are

CaB C Va/ Va1VsD

CsB C Vs/ Va1VsD

AssumexM $* the acid is i$ni%ed.


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Discussion

he formulas forxand the pH derived a"ove can "e used to estimate the pH of an# "uffer

solution% regardless how little salt or acid is used compared to their counter part.

hen the ratio Ca Csis "etween .* and *% the Henderson-Hassel"alch e!uition is a

convenient formula to use.

[H+] [A-]

Ka= ----------

[HA]

[A-]

pKa= pH - log (----)

[HA]


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A. ecause the c$ncentrati$n is hi"h@ we use the appr$2imati$nEH1F B CaKaD5/3

B =.=='56pH B 3.==

Notethe sharp increase in pH when .* m9 >3 drops? of "asic solution is added to

the solution.

. ?hen =.5 mG NaOH is added@ the c$ncentrati$n $* salt CsD@ andc$ncentrati$n $* acid Caare:CsB =.55.= M / 5=.5 B =.==II MCaB I.I5.= M / 5=.5 B =.I(

C. HA = H+ + A-

D. Ca-x x x

E.

F. [A-] = x + 0.0099 (= Cs)

G.

H. x (x + 0.0099)

I. Ka= -------------- = 1e-

!. 0.9" - x

#.

$. x2+ 0.0099 x= 9."e-% - 1e x

&. x2+ (0.0099 + 1e) x- 9."e-% = 0

N.

'. x = (-0.0099 + (0.00992+ 40.9"1e-)1/2) / 2

. = 0.00090*

.

,. pH = .042

Notethat using the Henderson-Hassel"alch e!uation will not #ield the correct

solution. Jo #ou now wh#Another wa# of asing the same !uestion.?


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cchiehuwaterl$$.ca

http://www.science.uwaterl$$.ca/cchieh/cact/c53'/&u!er.html

Hydr$lysis $* salts

:eturn to the Acid 6ase menu

A 6rief 5ntroduction to H#drol#sis Calculations

H#drol#sis happens when a su"stance chemicall# reacts with water. H#drol#sis should

"e distinguished from solvation% which is the process of water molecules associating

themselves with individual solute molecules or ions.

,. Salts of -ea Acids

5n general% all salts of wea acids "ehave the same% therefore we can use a generic salt

to represent all salts of wea acids. 9et NaA "e a generic salt of a wea acid and A

its anion. Here are two specific e$amples of salts of wea acids&

#u&stance $rmula and several others? could also "e used a"ovewithout affecting an# discussions of this topic. As a practical matter% onl# Na+and

P+tend to get used in e$amples.

he generic chemical reaction >in net ionic form? for h#drol#sis ma# "e written

thusl#&

A 1 H3O 00Q HA 1 OH

his reaction is of a salt of a wea acid >NO the acid? undergoing h#drol#sis% the

name for a chemical reaction with water. he salt is NaAc and it is reacting with thewater. Peep in mind that the acid >HAc? does not undergo h#drol#sis% the salt does.

5t is ver# important that #ou notice several things&
http://www.science.uwaterloo.ca/~cchieh/cact/c123/buffer.htmlhttp://www.chemteam.info/AcidBase/AcidBase.htmlhttp://www.chemteam.info/AcidBase/HydrolysisCalcsIntro.htmlhttp://www.science.uwaterloo.ca/~cchieh/cact/c123/buffer.htmlhttp://www.chemteam.info/AcidBase/AcidBase.htmlhttp://www.chemteam.info/AcidBase/HydrolysisCalcsIntro.html


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*? he Na+>notice onl# OH is written? 52 NO involved. 5ts source is the salt >NaA?

that is dissolving in the water and it JO12 NO affect the pH. 5ts presence in "oth

writing the chemical reactions and doing the calculations is deleted. However% eep in

mind that Na+is present in the solution. 2ome teacher might want to as a Bsnea#B

!uestion on a test.

;? HA is the =NJ522OC5A1J acid. Peep in mind that it is not the acid that maes

the acidic pH of a solution% it is the amount of h#drogen ion >or h#dronium ion% H 3O+%

if #ou wish?. 5n order to produce the h#drogen ion% the acid must dissociate.

3? here is free h#dro$ide ion >OH? in the solutionKK his is the thing that maes the

pH greater than ,.

Now% 5 can see a !uestion forming in #our mind. 5f there is acid >HA? and "ase >OH?%wh# dont the# Qust react and give "ac the reactants on the left side< Now% that reall#

is a good !uestion.

he answer< his reaction is an e!uili"rium. Now% if #ou are taing chemistr# for the

first time% #ou pro"a"l# Qust got done with e!uili"rium a few wees ago and it might

have "een hard to understand. hats understanda"le% "ut please realiIe that

e!uili"rium is one of more important concepts in chemistr#. Peep up the worKK

hen a chemical reaction comes to e!uili"rium% there is a mi$ture of all involvedsu"stances in the reaction vessel. his mi$ture is characteriIed "# a constant

composition. >Peep in mind that constant composition JO12 NO impl# e!ual

composition.? he e# point that maes a reaction come to e!uili"rium is that it is

reversi"le. his means that "oth the forward reaction and the reverse reaction can

happen% althought NO initiall# with e!ual pro"a"ilit#. he reaction comes to

e!uili"rium when the rates of the two reactions >forward and reverse? "ecome e!ual.

2o% while it is true that the HA and OH will react in the reverse direction% so can the

A and the H;O in the forward direction. he e# point is that the reaction happens insuch a wa# that a small amount >as opposed to Iero? of HA and OH are present at

e!uili"rium.

hen calculations are done% the important points will "e >*? how much OH is

formed and >;? what is the pH of the solution


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Ruic answers& >*? the amount of OH formed will "e greater than the *, value

present in pure water and >;? the pH will "e greater than ,% so the solution of the salt

of a wea acid will "e "asic.

,,. Salts of -ea Bases

5n general% all salts of wea "ases "ehave the same% therefore we can use a generic salt

to represent all salts of wea "ases. 9et 6 "e a generic "ase and H6+its salt.

>Compare how this is worded compared to the Bsalt of wea acidB discussion.? H6+is

a cation% "ut that word is not used as much in discussions as is BanionB is a"ove. Here

are two specific e$amples of salts of wea "ases&

#u&stance $rmula in net ionic form? for h#drol#sis reaction ma# "e

written thusl#&

H11 H3O 00Q 1 H'O1

his reaction is of a salt of a wea "ase >NO the "ase? undergoing h#drol#sis% the

reaction with water. he salt in this case is H6+Cl and it is reacting with the water.

:emem"er% the most common specific e$ample would "e ammonium chloride


64/94

>NH4+Cl?. Peep in mind that the "ase >generic e$ample / 6% specific e$ample /

ammonia or NH3? does not undergo h#drol#sis% the salt does.

5t is ver# important that #ou notice several things&

*? here is an anion involved% "ut it is usuall# not written. 8or e$ample Cl could "e

the anion% "ut it 52 NO involved. 5ts source is the salt >H6+Cl? that is dissolving in

the water and it JO12 NO affect the pH. 5ts presence in writing the appropriate

chemical reactions and doing the calculations is deleted. However% eep in mind that

Cl is present in the solution. 2ome teacher might want to as a Bsnea#B !uestion on

a test.

;? 6 is the =N':OONA1J "ase. Peep in mind that it is not the "ase that maes

the "asic pH of a solution% it is the amount of h#dro$ide ion >OH?. 5n order toproduce it% the "ase must protonated "# the water.

3? here is free h#dronium ion >H3O+? in the solutionKK his is the thing that maes the

pH less than ,.

Now% 5 can see a !uestion forming in #our mind. 5f there is "ase >6? and acid >H3O+?%

wh# dont the# Qust react and give "ac the reactants on the left side< Now% that reall#

is a good !uestion.

he answer% of course% is given in a"ove in the discussion of salts of wea acids. 5t

would "e the same e$planation here% so 5 wont repeat it. hat #ou might want to do%

however% is loo at the different phrasing in part 5 as compared to part 55.

Of course% when calculations are done% the important points will "e >*? how much

H3O+is formed and >;? what is the pH of the solution*? the amount of H3O+formed will "e greater than the *, value

present in pure water and >;? the pH will "e less than ,% so the solution of the salt of a

wea "ase will "e acidic.

http://www.chemteam.in*$/Acidase/Hydr$lysis.html

'
http://www.chemteam.info/AcidBase/Hydrolysis.htmlhttp://www.chemteam.info/AcidBase/Hydrolysis.html


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What is pH and Why is It

Important?The pH of normal human blood and tissues

is about !.". #f this pH is changed by $.% or

more& either up or down& it is a life'

threatening situation. ind out why here.

Donald Reinhardt

1$ months ago
https://suite.io/donald-reinhardthttps://suite.io/donald-reinhardt


66/94

pH - The Hydrogen Ion concentration in common situation from alkaline to neutral to acid

4ften pH is measured for soil, water, blood, urine andrelated clinical specimens, and many chemical reactions' pH

is an important chemical condition and pH &alues aresignificant and ha&e chemical consequences. C&en spas andswimming pools require pH chec"s, otherwise disinfectantsmay not be acti&e.

Basic Concepts of pH are Related to

Water Ionization%he pH is a measure of hydrogen ion concentration. Asample of absolutely pure water has a pH &alue of 5.:. %hepH scale ranges from : (acid) to +8 (basic). A pH of 5.: isconsidered neutralpH.

Bure water ioniGes to a limited degree to form H 0 (protons)

and (4H) - hydroyl ions/ H4H (water molecule) I H0and (4H)-.


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concentration. %hus, + in +: million reciprocated is +:million, and the log +: &alue of that is 5.:.

Whene&er pH is measured, hydrogen ion concentration is

determined. %hus, pH &alues of 7.:, 6.: and +.: epress,respecti&ely, that + in +::,:::, + in +,::: and + in +:concentrations of hydrogen ions (protons) are present.%hese acidities occur in wea"er, organic acids (citric, aceticand lactic acids) and strong, inorganic acids (H$l,hydrochloric acid' H


68/94

pH Litmus Paper to measure pH from 1 to 14 photo credit: Amazon.com


69/94

pH Meter - Small, Compact and Accurate pH MeterPhoto credit: ibchem.com

Hydrogen ion concentration and correlation to pH values in common chemical

environments.Photo Credit:Woods Hole Oceanographic Institute


70/94

Importance of pH in Living Systems,

Chemistry and Biochemistry

%he pH figures here re&eal se&eral typical pH &alues. %hephoto immediately abo&e this section clearly indicates thehydrogen ion concentration and pH relationship.


71/94

n summary, pH which is a measure of hydrogen ionconcentration, is critical to life and biochemistry and manyimportant chemical reactions.

Eesources

Alters, 3 and !. Alters. *::9,+iology. Lohn Wiley M 3onsnc., Hobo"en,


72/94

turn determines pr$perties $* the m$lecule. a!? + H;O>l? --S H3O+>a!? + A->a!?

Pa/ LH3O+MLA-M

LHAM

A "uffer s#stem can "e made "# mi$ing a solu"le compound that contains the

conQugate "ase with a solution of the acid such as sodium acetate with acetic acid orammonia with ammonium chloride. he a"ove e!uation for Pacan "e rearranged tosolve for the h#dronium ion concentration. 6# nowing the Paof the acid% the amountof acid% and the amount of conQugate "ase% the pH of the "uffer s#stem can "ecalculated.

LH3O+M / PaLHAM LA-M

pH / -logLH3O+M

Calculation of the pH of a 6uffer 2olution

Calculation of the pH of a 6uffer 2olution after Addition of a 2mall Amount of

2trong Acid
http://www.answers.com/Q/What_is_Biological_importance_of_pHhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#BufferpHhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddacidhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddacidhttp://www.answers.com/Q/What_is_Biological_importance_of_pHhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#BufferpHhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddacidhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddacid


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Calculation of the pH of a 6uffer 2olution after Addition of a 2mall Amount of

2trong 6ase

Calculation of the 6uffer Capacit#

Calculation of the p/ of a Buffer Solution

5n order to calculate the pH of the "uffer solution #ou need to now the amount ofacid and the amount of the conQugate "ase com"ined to mae the solution. heseamounts should "e either in moles or in molarities. he Paof the acid also needs to "enown.

Example0 A "uffer solution was made "# dissolving *. grams of sodium acetate in;. m9 of *. acetic acid. Assuming the change in volume when the sodiumacetate is not significant% estimate the pH of the acetic acidsodium acetate "uffersolution. he Pafor acetic acid is *., $ *-(.

8irst% write the e!uation for the ioniIation of acetic acid and the Pae$pression.

:earrange the e$pression to solve for the h#dronium ion concentration.

CH3COOH>a!? + H;O>l? --S H3O+>a!? + CH3COO

->a!?

LH3O+M / PaLCH3COOHM

LCH3COO-M

2econd% determine the num"er of moles of acid and of the conQugate "ase.

>*. CH3COOH?>;. m9?>* 9* m9? / .; mol CH3COOH

>*. g NaCH3COO?>* molE;.3 g? / .*;; mol NaCH3COO

2u"stitute these values% along with the Pavalue% into the a"ove e!uation and

solve for the h#dronium ion concentration. Convert the h#dronium ionconcentration into pH.

LH3O+M / >*., $ *-(?>.;.*;;? / ;.,) $ *-(

pH / 4.(@

Example0 Calculate the ratio of ammonium chloride to ammonia that is re!uired tomae a "uffer solution with a pH of ).. he Pa for ammonium ion is (.@ $ *-*.
http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddbasehttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddbasehttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Buffercapacityhttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddbasehttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Bufferaddbasehttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Buffercapacity


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8irst% write the e!uation for the ioniIation of the ammonium ion in water and

the corresponding Pa e$pression. :earrange the e!uation to solve for theh#dronium ion concentration.

NH4+>a!? + H;O>l? --S H3O+>a!? + NH3>a!?

Pa/ LH3O+MLNH3M

LNH4+M

LH3O+M / PaLNH4

+M

LNH3M

2econd% convert the pH "ac into the h#dronium ion concentration and then

su"stitute it into the a"ove e!uation along with the Pa. 2olve for the ratio ofammonium ion to ammonia.

LH3O+M / * $ *-)

* $ *-)/ (.@ $ *-*>NH4+NH3?

>NH4+NH3? / *.,E@*

A ratio of *.,@E moles of ammonium ion for ever# * mole of ammonia or *.,@E ammonium ion to * ammonia.

op

Calculation of the p/ of a Buffer Solution after Addition of a Small Amount of

Acid

hen a strong acid >H3O+? is added to a "uffer solution the conQugate "ase present inthe "uffer consumes the h#dronium ion converting it into water and the wea acid ofthe conQugate "ase.

A-

>a!? + H3O+

>a!? --S H;O>l? + HA>a!?

his results in a decrease in the amount of conQugate "ase present and an increase inthe amount of the wea acid. he pH of the "uffer solution decreases "# a ver# smallamount "ecause of this > a lot less than if the "uffer s#stem was not present?. AnB5C1B chart is useful in determining the pH of the s#stem after a strong acid has "eenadded.
http://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Tophttp://www.chem.purdue.edu/gchelp/howtosolveit/Equilibrium/Buffers.htm#Top


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Example0 (. m9 of .* HCl was added to a "uffer consisting of .;( molesof sodium acetate and .3 moles of acetic acid. hat is the pH of the "uffer afterthe addition of the acid< Pa of acetic acid is *., $ *-(.

8irst% write the e!uation for the ioniIation of acetic acid in water and the related

Pae$pression rearranged to solve for the h#dronium ion concentration.

CH3COOH>a!? + H;O>l? --S H3O+>a!? + CH3COO

->a!?

LH3O+M / PaLCH3COOHM

LCH3COO-M

2econd% mae an B5C1B chart. 9et B$B represent the h#dronium ion

concentration once e!uili"rium has "een re-esta"lished. e will assume that

all of the added acid is consumed.

CH3COOH>a!? H3O+>a!? CH3COO->a!?

5nitial Amount .3 moles >.( 9?>.* ? / .( moles .;( moles

Change in Amount + .( moles -.( moles - .( moles

1!uili"rium Amount .3( moles $ .; moles

2u"stitute into the Pae$pression and solve for the h#dronium ionconcentration. Convert the answer into pH.

LH3O+M / >*., $ *-(?>.3(.;? / ;.),( $ *-(pH / 4.(3

op

Calculation of the p/ of a Buffer Solution after Addition of a Small Amount of

Strong Base

hen a strong "ase >OH-? is added to a "uffer solution% the h#dro$ide ions areconsumed "# the wea acid forming water and the weaer conQugate "ase of the acid.

he amount of the wea acid decreases while the amount of the conQugate "aseincreases. his prevents the pH of the solution from significantl# rising% which itwould if the "uffer s#stem was not present.

OH->a!? + HA>a!? --S H;O>l? + A->a!?


76/94

he process for finding the pH of the mi$ture after a strong "ase has "een added issimilar to the addition of a strong acid shown in the previous section.

Example0 Calculate the pH of a "uffer solution that initiall# consists of .4 molesof ammonia and .;( moles of ammonium ion% after ;. m9 of .,( NaOH has

"een added to the "uffer. Pa for ammonium ion is (.@ $ *-*.

8irst% write the e!uation for the ioniIation of the ammonium ion and the related

Pae$pression solved for the h#dronium ion concentration.

NH4+>a!? + H;O>l? --S H3O

+>a!? + NH3>a!?

LH3O+M / PaLNH4

+M

LNH3M

2econd% mae an B5C1B chart. 9et B$B "e the concentration of the h#droniumion at e!uili"rium. he change in the amount of the ammonium ion will "ee!ual to the amount of strong "ase added >,( $ .; 9 / .*( mol?.

NH4+>a!? H3O+>a!? NH3>a!?

5nitial Amount .;( moles not needed .4 moles

Change in Amount - .*( moles not needed + .*( moles

1!uili"rium Amont .;3( moles $ .4*( moles

hird% su"stitute into the Pae$pression and solve for the h#dronium ionconcentration. Convert the answer into pH.

LH3O+M / >(.@ $ *-*?>.;3(.4*(? / 3.*, $ *-*pH / ).(

op

Calculation of the Buffer Capacity

he "uffer capactit# refers to the ma$imum amount of either strong acid or strong"ase that can "e added "efore a significant change in the pH will occur. his is simpl#a matter of stoichiometr#. he ma$imum amount of strong acid that can "e added ise!ual to the amount of conQugate "ase present in the "uffer. he ma$imum amount of

"ase that can "e added is e!ual to the amount of wea acid present in the "uffer.


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Example0 hat is the ma$imum amount of acid that can "e added to a "uffer made"# the mi$ing of .3( moles of sodium h#drogen car"onate with .( moles of sodiumcar"onate< How much "ase can "e added "efore the pH will "egin to show asignificant changeafter reading and stud#ingman# aspects of health and nutrition? that B9ow pHB alone% isthe :oot Cause for at least ( of diseases leading tohospitaliIation% cancer on down the line% etc% etc% includingps#chological distur"ances% and that people who J:5NPA1: are generall# H1A9HU '1O'91 >#ou need a half

gallon to *.( gallons ma$ per da# as per he erc anual%standard medical reference?. hose people who drin aterdont have the same pro"lems that ever#one else does. he#dont have to Qoin A.A.% the# dont have to "e hospitaliIed for's#chological distur"ances >unless depressed% in which case *mg of 6 itamins dail# will fi$ that !uic% 6* hiamine% 6;:i"oflavin% 6@ '#rido$ine% and 6*; Co"alium% in the form of a

6 Comple$ or 6 2tress or 6 * or 6 * sold at mostgroceries?.

ANU 611:AD1 provides #ou with necessar# 89=5J2 forsurvival% "ut ON9U A1: will help #ou maintain a properBpHB.


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2OJA% W=5C12% CO8811 V 1A are NO a 2u"stitute forA1:&as 2oda >and other "everages? have a dangerousl# low pH. 5

shae hands with people% and 3 to @ seconds later m# hand is"urning due to their low pH% 5 can imagine how much the# aresuffering >"ut the# are used to it% "ut if the# onl# new% and 5 dotr# to help them when 5 can?. 5 have a health food e$pert friendwho wors in the medical industr#% and she tells me of patientswho come in with maQor pro"lems% and their "lood test resultsshow a ver# low level of pH >near death?. No wonder. he#dont drin water. he# thin other "everages will do.

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1p/ factor1% is a N=61:% which represents the range orscale of AC5J5U >which eat awa# at things? to A9PA95N1 or6A21 >which "uilds up?.

http://www.archure.net/salus/ph.html

!cidic and !l'aline ?ood 4ist

hat ou &eed To 3no) (bout Being Health* D (l/aline

The benefits of being al/aline is an opportunit* not ust to tal/ about but

to experience an extraordinar* health* lifest*le>

eople (ar#% but for most the ideal diet is :Q al'aliing and 2Q acidif#ing f

draft exp 3 preparing buffer solutions

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