chap 11 acids and bases.ppt - bakersfield college spring_10/chap_11... · chapter 3 about acids and...

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Chapter 11

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Acid–Base Concepts

1. Arrhenius Concept of Acids and Bases

2. Brønsted–Lowry Concept of Acids and Bases

3. Lewis Concept of Acids and Bases

Acid and Base Strengths

4. Relative Strengths of Acids and Bases

5. Molecular Structure and Acid Strength

Contents and Concepts

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Self-Ionization of Water and pH

6.Self-Ionization of Water

7.Solutions of a Strong Acid or Base

8.The pH of a Solution

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Learning Objectives

Acid Base Concepts

• Arrhenius Concept of Acids and Base

– a. Define acid and base according to the Arrhenius concept.

• Brønsted–Lowry Concept of Acids and Bases

– a. Define acid and base according to the Brønsted–Lowry concept.

– b. Define the term conjugate acid–base pair.

– c. Identify acid and base species.

– d. Define amphiprotic species.

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3. Lewis Concept of Acids and Bases

• a. Define Lewis acid and Lewis base.

• b. Identify Lewis acid and Lewis base species.

Acid and Base Strengths

4. Relative Strengths of Acids and Bases

– a. Understand the relationship between the strength of an acid and that of its conjugate base.

– b. Decide whether reactants or products are favored in an acid–base reaction.

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5. Molecular Structure and Acid Strength

• a. Note the two factors that determine relative

acid strengths.

• b. Understand the periodic trends in the

strengths of the binary acids HX.

• c. Understand the rules for determining the

relative strengths of oxoacids.

• d. Understand the relative acid strengths of a

polyprotic acid and its anions.

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Self-Ionization of Water and pH

6. Self-Ionization of Water

• a. Define self-ionization (or autoionization).

• b. Define the ion-product constant for water.

7. Solutions of a Strong Acid or Base

– a. Calculate the concentrations of H3O+ and

OH- in solutions of a strong acid or base

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8. The pH of a Solution

1. Define pH.

2. Calculate the pH from the hydronium-ion

concentration.

3. Calculate the hydronium-ion concentration

from the pH.

4. Describe the determination of pH by a pH

meter and by acid–base indicators.

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• When gaseous hydrogen chloride meets

gaseous ammonia, a smoke composed of

ammonium chloride is formed.

• HCl(g) + NH3(g) � NH4Cl(s)

• This is an acid–base reaction.

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• We will examine three ways to explain

acid–base behavior:

• Arrhenius Concept

• Brønsted–Lowry Concept

• Lewis Concept

H+ and OH-

H+ = protondonor

acceptordonor

acceptorelectron pair

Note: H+ in water is H3O+

acid base

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Acid-Base Concepts

• Antoine Lavoisier was

one of the first

chemists to try to

explain what makes a

substance acidic.

– In 1777, he proposed that oxygen was an essential element in acids.

– The actual cause of acidity and basicity was ultimately explained in terms of the effect these compounds have on water by Svante Arrhenius in 1884.

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Acid-Base Concepts

• In the first part of this chapter we will look at

several concepts of acid-base theory including:

– The Arrhenius concept

– The Bronsted Lowry concept

– The Lewis concept

This chapter expands on what you learned in Chapter 3 about acids and bases.

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Arrhenius Concept of Acids and Bases

• According to the Arrhenius concept of acids and

bases, an acid is a substance that, when

dissolved in water, increases the concentration

of hydronium ion (H3O+).

– Chemists often use the notation H+(aq) for the H3O

+(aq) ion, and call it the hydrogen ion.

– Remember, however, that the aqueous hydrogen ion is actually chemically bonded to water, that is, H3O

+.

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• The Arrhenius concept limits bases to

compounds that contain a hydroxide ion.

• The Brønsted–Lowry concept expands the

compounds that can be considered acids and

bases.

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The H3O+ is

shown here

hydrogen bonded to three water molecules.

• According to the Arrhenius concept of acids and

bases, an acid is a substance that, when

dissolved in water, increases the concentration

of hydronium ion (H3O+).

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• A base, in the Arrhenius concept, is a substance

that, when dissolved in water, increases the

concentration of hydroxide ion, OH-(aq).

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• In the Arrhenius concept, a strong acid is a

substance that ionizes completely in aqueous

solution to give H3O+(aq) and an anion

– Other strong acids include HCl, HBr, HI, HNO3 , and H2SO4.

– An example is perchloric acid, HClO4.

)aq(ClO)aq(OH)l(OH)aq(HClO 4324

−−−−++++++++→→→→++++

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• In the Arrhenius concept, a strong base is a

substance that ionizes completely in aqueous

solution to give OH-(aq) and a cation.

– Other strong bases include LiOH, KOH, Ca(OH)2, Sr(OH)2, and Ba(OH)2.

– An example is sodium hydroxide, NaOH.

)aq(OH)aq(Na)s(NaOHOH2

−−−−++++++++→→→→

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• Most other acids and bases that you encounter are

weak. They are not completely ionized and exist

in reversible reaction with the corresponding ions.

– Ammonium hydroxide, NH4OH, is a weak base.

– An example is acetic acid, HC2H3O2.

(aq)OHC(aq)OH 2323

−−−−++++++++)l(OH)aq(OHHC 2232

++++

)aq(OH)aq(NH)aq(OHNH 44

−−−−++++++++

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Brønsted-Lowry Concept of Acids and Bases

• A base is the species accepting the protonin a proton-transfer reaction.

– In any reversible acid-base reaction, both forward and reverse reactions involve proton transfer.

• According to the Brønsted-Lowry concept, an

acid is the species donating the proton in a

proton-transfer reaction.

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• Consider the reaction of NH3 and H20.

)aq(OH)aq(NH)l(OH)aq(NH 423

−−−−++++++++++++


−−−−++++++++++++

H+

base acid

– In the forward reaction, NH3 accepts a proton

from H2O. Thus, NH3 is a base and H2O is an acid.

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• Brønsted–Lowry Concept of Acids and

Bases

• An acid–base reaction is considered a proton

(H+) transfer reaction.

H+

H+ H+

H+

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• Consider the reaction of NH3 and H2O.


−−−−++++++++++++

H+

baseacid


−−−−++++++++++++

base acid

– The species NH4+ and NH3 are a conjugate

acid-base pair.

– A conjugate acid-base pair consists of two

species in an acid-base reaction, one acid and one base, that differ by the loss or gain of a proton.

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Brønsted-Lowry Concept of Acids and Bases

• Consider the reaction of NH3 and H2O.

– The Brønsted-Lowry concept defines a species as an acid or a base according to its function in the proton-transfer

reaction.


−−−−++++++++++++

base acid

– Here NH4+ is the conjugate acid of NH3

and NH3 is the conjugate base of NH4+.

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• Substances in the acid–base reaction that differ

by the gain or loss of a proton, H+, are called a

conjugate acid–base pair. The acid is called the

conjugate acid; the base is called a conjugate

base.

•

Acid Base Conjugate acid

Conjugate base

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• What is the conjugate acid of H2O?

• What is the conjugate base of H2O?

The conjugate acid of H2O has gained a proton.

It is H3O+.

The conjugate base of H2O has lost a proton.

It is OH-.

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• Label each species as an acid or base.

Identify the conjugate acid-base pairs.

a. HCO3-(aq) + HF(aq) H2CO3(aq) + F-(aq)

b. HCO3-(aq) + OH-(aq) CO3

2-(aq) + H2O(l)

Base Acid Conjugate base

Conjugate acid

Acid Base Conjugate

acid

Conjugate

base

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• A Brønsted–Lowry acid is the species

donating a proton in a proton-transfer reaction;

it is a proton donor.

• A Brønsted–Lowry base is the species

accepting a proton in a proton-transfer

reaction; it is a proton acceptor.

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• Some species can act as an acid or a base.

– An amphoteric species is a species that can act either as an acid or a base (it can gain or lose a proton).

– For example, HCO3- acts as a proton donor (an acid) in

the presence of OH-

)l(OH)aq(CO)aq(OH)aq(HCO 2

2

33++++→→→→++++

−−−−−−−−−−−−

–H+

)l(OH)aq(CO)aq(OH)aq(HCO 2

2

33++++→→→→++++

−−−−−−−−−−−−

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– An amphoteric species is a species that can act either as an acid or a base (it can gain or lose a proton).

– Alternatively, HCO3- can act as a proton acceptor

(a base) in the presence of HF.

)aq(F)aq(COH)aq(HF)aq(HCO 323

−−−−−−−−++++→→→→++++

H+

• The amphoteric characteristic of water is

important in the acid-base properties of

aqueous solutions.

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– Water reacts as an acid with the base NH3.


−−−−++++++++→→→→++++

H+

– Water can also react as a base with the acid HF.

)aq(OH)aq(F)l(OH)aq(HF 32

++++−−−−++++→→→→++++

H+

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• In the Brønsted-Lowry concept:

1. A base is a species that accepts protons; OH-

is only one example of a base.

2. Acids and bases can be ions as well as molecular substances.

3. Acid-base reactions are not restricted to aqueous solution.

4. Some species can act as either acids or bases depending on what the other reactant is.

Look at Example 15.1 Do Exercise 15.1 See Problems 15.35-36

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Lewis Concept of Acids and Bases

• The Lewis concept defines an acid as an electron pair acceptor and a base as an electron pair donor.

– This concept broadened the scope of acid-

base theory to include reactions that did not

involve H+.

– The Lewis concept embraces many

reactions that we might not think of as acid-base reactions.

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• The reaction of boron trifluoride with

ammonia is an example.

+ N

H

H

H:

::

: B

F

F

F

: :

::

::

::

: B

F

F

F

: :

::

:: N

H

H

H

– Boron trifluoride accepts the electron pair, so it is a Lewis acid. Ammonia donates the electron pair, so it is the Lewis base.

See Example 15.2 and Problems 15.39-42Do Exercise 15.2

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•

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Relative Strength of Acids and Bases

• The Brønsted-Lowry concept introduced

the idea of conjugate acid-base pairsand proton-transfer reactions.

– We consider such acid-base reactions to

be a competition between species for

hydrogen ions.

– From this point of view, we can order acids by their relative strength as hydrogen ion donors.

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– The stronger acids are those that lose their hydrogen ions more easily than other acids.

– Similarly, the stronger bases are those that hold onto hydrogen ions more strongly than other bases.

– If an acid loses its H+, the resulting anion is now in a position to reaccept a proton, making it a Brønsted-Lowry base.

– It is logical to assume that if an acid is considered strong, its conjugate base (that is, its anion) would be weak, since it is unlikely to accept a hydrogen ion.

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• Consider the equilibrium below.

– In this system we have two opposing Brønsted-Lowry acid-base reactions.

– In this example, H3O+ is the stronger of the two

acids. Consequently, the equilibrium is skewed toward reactants.

(aq)OHC(aq)OH 2323

−−−−++++++++)l(OH)aq(OHHC 2232

++++

acid acidbase base

conjugate acid-base pairs

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• Consider the equilibrium below.

(aq)OHC(aq)OH 2323

−−−−++++++++)l(OH)aq(OHHC 2232

++++

acid acidbase base

conjugate acid-base pairs

– Table 15.2 outlines the relative strength of some common acids and their conjugate bases.

– This concept of conjugate pairs is fundamental to understanding why certain salts can act as acids or bases.

Do Exercise 15.2 See Example 15.3 and Problems 15.45-48

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Acetate Ion

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Molecular Structure and Acid Strength

• Two factors are important in determining

the relative acid strengths.

– One is the polarity of the bond to which the hydrogen atom is attached.

– The H atom should have a partial positive charge:

XH −−−−δ+δ+δ+δ+ δδδδ−−−−

– The more polarized the bond, the more easily the proton is removed and the greater the acid strength.

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– The second factor is the strength of the bond. Or, in other words, how tightly the proton is held.

– This depends on the size of atom X.

XH −−−−δδδδ+ δδδδ-

– The larger atom X, the weaker the bond and the greater the acid strength.

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• For a binary acid, as the size of X in HX

increases, going down a group, acid strength

increases.

• For a binary acid, going across a period, as the

electronegativity increases, acid strength

increases.

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• Which is a stronger acid: HF or HCl?

• Which is a stronger acid: H2O or H2S?

• Which is a stronger acid: HCl or H2S?

HF and HClThese are binary acids from the same group, so we compare the size of F and Cl. Because Cl is larger, HCl is the stronger acid.

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• H2O and H2S

• These are binary acids from the same group, so

we compare the size of O and S. Because S is

larger, H2S is the stronger acid.

• HCl and H2S

• These are binary acids from the same period,

but different groups, so we compare the

electronegativity of O and S. Because Cl is

more electronegative, HCl is the stronger acid.

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• Consider a series of binary acids from a

given column of elements.

– As you go down the column of elements, the radius increases markedly and the H-X bond strength decreases.

– You can predict the following order of acidic strength.

HIHBrHClHF <<<<<<<<<<<<

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• As you go across a row of elements, the

polarity of the H-X bond becomes the dominant factor.

– As electronegativity increases going to the right, the polarity of the H-X bond increases and the acid strength increases.


HFOHNH 23<<<<<<<<

HFOHNH 23<<<<<<<<

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• Consider the oxoacids. An oxoacid has

the structure:

– The acidic H atom is always attached to an O atom, which in turn is attached to another atom Y.

– Bond polarity is the dominant factor in the relative strength of oxoacids.

– This, in turn, depends on the electronegativity of the atom Y.

−−−−−−−−−−−− YOH

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• Consider the oxoacids. An oxoacid has

the structure:

– If the electronegativity of Y is large, then the O-H bond is relatively polar and the acid strength is greater.

−−−−−−−−−−−− YOH


HOIHOBrHOCl >>>>>>>>

– Other groups, such as O atoms or O-H groups, may be attached to Y.

– With each additional O atom, Y becomes effectively more electronegative.

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• For oxoacids, several factors are relevant: the

number and bonding of oxygens, the central

element, and the charge on the species.

• For a series of oxoacids, (OH)mYOn, acid

strength increases as n increases.

(OH)Cln = 0

(OH)ClOn = 1

(OH)ClO2

n = 2

(OH)ClO3

n = 3

Weakest Strongest

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– As a result, the H atom becomes more acidic.

– The acid strengths of the oxoacids of chlorine increase in the following order.

432 HClOHClOHClOHClO <<<<<<<<<<<<

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• For a series of oxoacids differing only in the

central atom Y, the acid strength increases with

the electronegativity of Y.

Stronger Weaker

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• Consider polyprotic acids and their

corresponding anions.

– Each successive H atom becomes more difficult to remove.

– Therefore the acid strength of a polyprotic acid and its anions decreases with increasing negative charge.

4342

2

4 POHPOHHPO <<<<<<<<−−−−−−−−

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• The acid strength of a polyprotic acid and its

anions decreases with increasing negative

charge.

•H2CO3 is a stronger acid than HCO3-.

•H2SO4 is a stronger acid than HSO4-.

•H3PO4 is a stronger acid than H2PO4-.

•H2PO4- is a stronger acid than HPO4

2-.

• A reaction will always go in the direction from

stronger acid to weaker acid, and from stronger

base to weaker base.

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•Decide which species are favored at the completion of the following reaction:

•HCN(aq) + HSO3-(aq)

•CN-(aq) + H2SO3(aq)

We first identify the acid on each side of the

reaction: HCN and H2SO3.

Next, we compare their acid strength:H2SO3 is stronger.

This reaction will go from right to left (), and the reactants are favored.

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Do Exercise 15.4

See Problems 15.51-52

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• Self-Ionization of Water

• H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

• Self-Ionization of Water

• H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

Base Acid Conjugate base

Conjugate acid

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Self-ionization of Water

• Self-ionization is a reaction in which two like

molecules react to give ions.

– In the case of water, the following equilibrium is established.

– The equilibrium-constant expression for this system is:

)aq(OH)aq(OH)l(OH)l(OH 322

−−−−++++++++++++

2

2

3c

]OH[

]OH][OH[K

−−−−++++

====

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– The concentration of ions is extremely small, so the concentration of H2O remains

essentially constant. This gives:

constant

• Self-ionization is a reaction in which two like

molecules react to give ions.

]OH][OH[K]OH[ 3c2

2

−−−−++++====

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– We call the equilibrium value for the ion product [H3O

+][OH-] the ion-product constant for water, which is written Kw.

– At 25 oC, the value of Kw is 1.0 x 10-14.

– Like any equilibrium constant, Kw varies with temperature.

]OH][OH[K 3w

−−−−++++====

– Because we often write H3O+ as H+, the ion-

product constant expression for water can be written:

]OH][H[Kw

−−−−++++====

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• These ions are produced in equal numbers in pure

water, so if we let x = [H+] = [OH-]

– Thus, the concentrations of H+ and OH- in pure water are both 1.0 x 10-7 M.

– If you add acid or base to water they are no longer

equal but the Kw expression still holds.

C25at )x)(x(100.1o14 ====××××

−−−−

714100.1100.1x

−−−−−−−−××××====××××====

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• H2O(l) + H2O(l) H3O+(aq) + OH-(aq)

• We call the equilibrium constant the ion-

product constant, Kw.

Kw = [H3O+][OH-]

At 25°C, Kw = 1.0 × 10-14

• As temperature increases, the value of Kw

increases.

See Example 15.4 and Problems 15.53-54Do Exercise 15.5

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Solutions of Strong Acid or Base

• In a solution of a strong acid you can

normally ignore the self-ionization of water

as a source of H+(aq).

– The H+(aq) concentration is usually

determined by the strong acid

concentration.

– However, the self-ionization still exists and

is responsible for a small concentration of OH- ion.

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• As an example, calculate the concentration of OH-

ion in 0.10 M HCl.

Because you started with 0.10 M HCl (a

strong acid) the reaction will produce 0.10

M H+(aq).

)aq(Cl)aq(H)aq(HCl−−−−++++

++++→→→→

– Substituting [H+]=0.10 into the ion-product expression, we get:

]OH)[10.0(100.114 −−−−−−−−

====××××

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• As an example, calculate the concentration of OH-

ion in 0.10 M HCl.

Because you started with 0.10 M HCl (a strong acid) the reaction will produce 0.10 M H+(aq).

– Substituting [H+]=0.10 into the ion-product expression, we get:

)aq(Cl)aq(H)aq(HCl−−−−++++

++++→→→→

MOH13-

14-

101.010.0

101.0][ ×=

×=

−

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• Similarly, in a solution of a strong base you

can normally ignore the self-ionization of

water as a source of OH-(aq).

– The OH-(aq) concentration is usually determined by the strong base

concentration.

– However, the self-ionization still exists and is responsible for a small concentration of

H+ ion.

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• As an example, calculate the concentration of H+

ion in 0.010 M NaOH.

Because you started with 0.010 M NaOH (a strong base) the reaction will produce 0.010

M OH-(aq).

– Substituting [OH-]=0.010 into the ion-product expression, we get:


−−−−++++++++→→→→

)010.0](H[100.114 ++++−−−−

====××××

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Because you started with 0.010 M NaOH (a strong base) the reaction will produce 0.010 M OH-(aq).


−−−−++++++++→→→→

– Substituting [OH-]=0.010 into the ion-

product expression, we get:

MH12-

14-

101.0010.0

101.0][ ×=

×=

+

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• By dissolving substances in water, you can

alter the concentrations of H+(aq) and OH-

(aq).

– In a neutral solution, the concentrations of H+(aq) and OH-(aq) are equal, as they are in pure water.

– In an acidic solution, the concentration of H+(aq) is greater than that of OH-(aq).

– In a basic solution, the concentration of OH-(aq) is greater than that of H+(aq).

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• At 25°C, you observe the following

conditions.

– In an acidic solution, [H+] > 1.0 x 10-7 M.

– In a neutral solution, [H+] = 1.0 x 10-7 M.

– In a basic solution, [H+] < 1.0 x 10-7 M.

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The pH of a Solution

• Although you can quantitatively describe

the acidity of a solution by its [H+], it is

often more convenient to give acidity in

terms of pH.

– The pH of a solution is defined as the

negative logarithm of the molar hydrogen-ion concentration.

]Hlog[pH++++

−−−−====

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• For a solution in which the hydrogen-ion

concentration is 1.0 x 10-3, the pH is:

– Note that the number of decimal places in

the pH equals the number of significant figures in the hydrogen-ion concentration.

00.3)100.1log(3

=×−=−

pH

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• In a neutral solution, whose hydrogen-ion concentration

is 1.0 x 10-7, the pH = 7.00.

• For acidic solutions, the hydrogen-ion concentration is greater than 1.0 x 10-7, so the pH is less than 7.00

• .

• Similarly, a basic solution has a pH greater than 7.00.

• Figure 15.8 shows a diagram of the pH scale and the pH values of some common solutions.

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Figure 15.8: The pH Scale

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•Calculate the hydronium and hydroxide ion concentration at 25°C in

• a. 0.10 M HCl

• b. 1.4 × 10-4 M Mg(OH)2

a. When HCl ionizes, it gives H+ and Cl-. So [H+] = [Cl-] = [HCl] = 0.10 M.

a. When Mg(OH)2 ionizes, it gives Mg2+ and 2 OH-.

So [OH-] = 2[Mg2+] = 2[Mg(OH)2] = 2.8 × 10-4 M.

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A Problem to Consider

• A sample of orange juice has a hydrogen-ion

concentration of 2.9 x 10-4 M. What is the pH?

54.3pH ====

)109.2log(pH4−−−−

××××−−−−====

]Hlog[pH++++

−−−−====

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A Problem to Consider

• The pH of human arterial blood is 7.40. What

is the hydrogen-ion concentration?

)pHlog(anti]H[ −−−−====++++

)40.7log(anti]H[ −−−−====++++

M100.410]H[840.7 −−−−−−−−++++

××××========

See Example 15.5 and Problems 15.67-68 Do Exercise 15.7

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• A has 5 H3O+ and 5 OH-. It is neutral.

• B has 7 H3O+ and 3 OH-. It is acidic.

• C has 3 H3O+ and 7 OH-. It is basic.

• Listed from most acidic to most basic: B, A, C.

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The pOH of a Solution

• A measurement of the hydroxide ion

concentration, similar to pH, is the pOH.

– The pOH of a solution is defined as the

negative logarithm of the molar hydroxide-

ion concentration.

]OHlog[pOH−−−−

−−−−====

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The pOH of a Solution



– Then because Kw = [H+][OH-] = 1.0 x 10-14

at 25 oC, you can show that

00.14pOHpH ====++++

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– Then because Kw = [H+][OH-] = 1.0 x 10-14

at 25 oC, you can show that

00.14pOHpH ====++++

See Exercise 15.6 and Problems 15.75-76 Do Exercise 15.9-10

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http://www.quia.com/rr/4051.html

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• The pH of a solution can accurately be

measured using a pH meter (see Figure

15.9).

– Although less precise, acid-base indicators are often used to measure pH because they usually change color within a narrow pH range.

– Figure 15.8 shows the color changes of various acid-base indicators.

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Figure

15.9:

A digital

pH meter.Photo

courtesy of

American

Color.

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Operational Skills

1 Identify acid and base species.

2 Identify Lewis acids and bases.

3 Decide whether reactants or products are favored in an

Acid-base reaction.

4 Calculate Concentrations of H3O+ and OH-.

5 Calculate the pH from the hydronium concentration

and vise versa.

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Figure 15.12: Preparation of Sodium Hydroxide by Hydrolysis

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Problem 15.27

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Problem 15.28

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Problem 15.37

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Problem 15.38

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Operational Skills

• Identifying acid and base species

• Identifying Lewis acid and base species

• Deciding whether reactants or products are favored in an acid-base reaction

• Calculating the concentration of H+ and OH- in solutions of strong acid or base

• Calculating the pH from the hydrogen-ion concentration, and vice versa

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chap 11 acids and bases.ppt - bakersfield college spring_10/chap_11... · chapter 3 about acids and...

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