general chemistry ii 2302102

General Chemistry II2302102

Acid and Base Equilibria

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Lecture 1

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mailto:[email protected]

Acids and Bases - 3 Lectures

• Autoionization of Water and pH

• Defining Acids and Bases

• Interaction of Acids and Bases with Water

• Buffer Solutions

• Acid-Base Titrations

Outline - 5 Subtopics

By the end of this lecture AND completion of the set problems, you should be able to:

• Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O

+] and [OH-].

• Know the Arrhenius and Brønsted definitions of acids and bases.

• Understand and know examples of conjugate acid-base pairs.

• Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids.

• Understand the common ion effect in acid ionization.

Objectives - Lecture 1

Acids and Bases

Acids & Bases

Gold mining frequently uses the base, cyanide (CN-), in the extraction process. Tailings dams sometimes have high concentrations of CN-. The conjugate acid of CN-, HCN, is extremely toxic.

“Acid mine drainage” (AMD) is naturally produced by the exposure of sulfide ores to water. This process is greatly exacerbated by mining the sulfide ores (for Cu, Pb, Zn etc). AMD results in streams with elevated H+ concentrations and pH values in the range 1-3.

A gold mining tailings dam leaks into an AMD-affected stream - would you evacuate?

Would you evacuate?

Acids & Bases - 2 Lectures

Introduction to acids and basesStrong & Weak AcidsConjugate acid-base pairsCommon Ion Effect

BasesBuffersIndicatorsTitrations

• Strong Acids• Weak Acids

Outline

Acids & Bases - Lecture 1

By the end of this lecture AND completion of the set problems, you should be able to:

• Define BrØnsted acids and bases• Calculate [H+], [OH-] and pH• Distinguish between strong & weak acids• Calculate equilibrium concentrations of acids &

bases using acidity constants• Identify conjugate acid-bases pairs• Determine the effect of adding a common ion on

equilibrium concentrations

Objectives

Acids & Bases - Why are we interested?

Acids - solvents (dissolve other materials)- ores, food (stomach contains HCl)

Bases - solvents - cleaners like “Draino”, bleach

Large Range of Industrial Processes:

Acids - wine, beer, citrus fruits, vinegar, coffeeFoods:

Blood - pH 7.3-7.5 falls below 6.8 will be fatal (acidosis)Most bodily functions under ‘circumneutral’ conditions

Physiology:

‘Natural’ Erosion of Limestone Caves“acid rain” - dissolved H2SO4 & HNO3 in upper atmosphere

- rainfall runoff into acid lakes - devoid of life

Environment:

Acids & Bases

1. ACID Species which can donate a hydrogen ion (H+).

2. BASE Species which can accept a hydrogen ion.

3. AMPHOTERIC SPECIES Species which can act as both an acid and a

base. e.g. HCO3- (CO3

2-, H2CO3), HSO4-

BrØnsted-Lowry Definitions:

• Hydrogen Ion

What is the Hydrogen Ion?• Hydrogen Atom Atomic weight = 1

1 proton + 1 electron

• Hydrogen Ion in Water

Represented as:

H+(aq) or H3O+

Atomic weight = 1

1 proton

Acid-Base ReactionIn fact a proton-transfer reaction in which the proton is transferred from the acid to the base.

Acidity - pH

pH convenient measure of H3O+ concentration

H O

concentration of H O M

pH

pH

3

3

10

10

e.g. pH = 8.5

H O

concentration of H O M

38

9

39

10

3 16 10

3 16 10

.5

.

.

Acidity measure of the H3O+ concentration

AQUEOUS SYSTEMS

2H2O() H3O+(aq) + OH-(aq)

Equilibrium between H3O+ and OH-:

at equilibrium [H3O+][OH-] = Kw

at 25 °C Kw = 1.00 x 10-14

(c.f. Ksp = [Ag+][Cl-])

AQUEOUS SYSTEMS

2H2O() H3O+(aq) + OH-(aq)

Kw = [H3O+][OH-]

at 25 °C Kw = 1.00 x 10-14

Pure water at 25 °C. If 2z mol/L of water react:

[H3O+] = z and [OH-] = z

z2 = 1.00 x 10-14 z = 1.00 x 10-7

and pH = 7.00

DEFINITIONS

“acid” pH < 7.00 [H3O+(aq)] > [OH-(aq)]

“basic” pH > 7.00 [H3O+(aq)] < [OH-(aq)]

“neutral” pH = 7.00 [H3O+(aq)] = [OH-(aq)]

Acid-Base Reaction: Proton-Transfer between an Acid and a Base (Water)

• According to the Brönsted-Lowry concept, when HCl gas is dissolved in water to form the solution of hydrochloric acid, a proton-transfer reaction occurs:

H2O (l) + HCl (g) H3O+ (aq) + Cl- (aq)

BaseBase(Proton(Proton

Acceptor)Acceptor)HydroniumHydronium

IonIonA hydrated proton = HA hydrated proton = H++ (aq) = H (aq) = H33OO++

AcidAcid(Proton(ProtonDonor)Donor)

ClClOOHH ::

HH

++ HH :: HH OO HH

HH:: ClCl ::

--++

++::

:::: :: ::

::

ACIDSACIDS

ACIDS

Note HA may be a molecule anion

or cation

HCl aq HCOOH aq

HCO aq

NH aq

( ), ( )

( )

( )

3

4

Reaction between acids and water

H2O() + HA(aq) H3O+(aq) + A-(aq)

Equilibrium constant Ka

At equilibrium

A aq H O aq

HA aqKa

( ) ( )

( )

3

ACIDSReaction between acids and water

Equilibrium constant Ka


pK K

so that K

a a

apK

a

log10

10

Usually tabulated as pKa:

Note - this is exactly the same relationship as between [H+] and pH (pH = - log10 [H+])

Strength of Acids

• Strong versus Weak acids

• The strength of an acid is related to the position of the equilibrium above

• It is NOT related to the corrosive ability (this often causes confusion)

• As we shall see, HF is a weak acid, yet it is one of the most corrosive acids known.


STRONG ACID

Position of equilibrium almost completely to the right (acid is almost totally ionized)

e.g. 0.1 M HCl(aq) [H3O+] = 0.1 [Cl-] = 0.1 [HCl] not known ( < 10-10)

pKa < -10



Acids are ElectrolytesNon

No ionsin solution

Strong

Many ionsin solution -

Weak

Few ionsin solution -

Strong Acids Weak Acids

WEAK ACIDS

1. Definition of weak acid Position of equilibrium lies to the left. Only

a small fraction of the acid reacts with the water and is ionized. i.e. Acid is weakly ionized

Ka small < 10-3

pKa large > 3



WEAK ACIDS

initial 0.5 M 0 0equilibrium (0.5 - x) M x M x M

2. Dissociation in water


eg. 0.5 M HF(aq) Ka = 6.8 x 10-4

If x mol/L react then we have

H2O() + HF(aq) H3O+(aq) + F-(aq)


At equilibrium:

MHF

M

M

OH

F

48.0][

108.1][

108.1][

23

2


So HF is ca. 4% ionized

WEAK ACIDS

Diprotic & Triprotic Acids• Acids that we’ve considered thus far have

been monoprotic (donate 1 proton) e.g. HCl, HNO3

• Other acids can donate 2 or 3 protons - diprotic or triprotic acids e.g. H2SO4, H3PO4H2O() + H2A(aq) H3O+(aq) + HA-(aq)

]A[

][ ][

HHAOH

2

31

K a

H2O() + HA-(aq) H3O+(aq) + A2-(aq)

][

][ ][

HA

AOH 23

2

K a

Diprotic & Triprotic Acids

]A[

][ ][

HHAOH

2

31

K a

][

][ ][

HA

AOH 23

2

K a

Acid Formula Ka1 pKa1 Ka2 pKa2 Ka3 pKa3

Sulfuric H2SO4 ca. 100 ca. -2 1.2 x 10-21.92

Oxalic H2C2O4 5.9 x 10-21.23 6.4 x 10-5

4.19

Phosphoric H3PO4 7.52 x 10-32.12 6.23 x 10-8

7.21 2.2 x 10-1312.67

Successive pKas increase in magnitude

- removal of successive protons results in weaker acids

Conjugate Acid-Base Pairs

- Reverse reactionH2O (l) + NH3 (aq) NH4

+ (aq) + OH- (aq)

Consider the proton-transfer reaction:– Forward reaction

H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq)

Acid(ProtonDonor)

Base(Proton

Acceptor)

Acid(ProtonDonor)

BaseBase(Proton(Proton

Acceptor)Acceptor)


H2O (l) + NH3 (aq) NH4+ (aq) + OH- (aq)

Acid 1 Base 1 Acid 2 Base 2

Conjugate Acid-Base Pair

Conjugate Acid-Base Pair

For the general equation

B + HA HB+ + A-


Acid(Proton (Proton Donor)Donor)

Base(Proton (Proton

Acceptor)Acceptor)

Acid(Proton (Proton Donor)Donor)

Base(Proton (Proton

Acceptor)Acceptor)

The acid HB+ results from the base B gaining a proton. HB+-B is a conjugate acid-base pair.

The base A- results from the acid HA losing a proton. HA-A- is a conjugate acid-base pair, too.


• Any two substances that differ by one proton

are a conjugate acid-base pair

• Can write the formula of the conjugate base of

any acid simply by removing the proton:If HCO3

- is the acid, CO32- is the conjugate base

If HCO3- is the base, H2CO3 is the conjugate acid

(HCO3- is amphoteric)


• The conjugate (base) of a strong acid is a weak base

• The conjugate (base) of a weak acid is a strong base

• The conjugate (acid) of a strong base is a weak acid

• The conjugate (acid) of a weak base is a strong acid

Direction of Acid-Base Reactions

NH4+ is a stronger acid than H2O, and

OH- is a stronger base than NH3

So reaction proceeds spontaneously to the left

• Acid-Base Reactions proceed spontaneously with the strongest acid and strongest base forming the weakest acid and the weakest base

Returning to:H2O (l) + NH3 (aq) NH4

+ (aq) + OH- (aq)

The Common Ion Effect

Recall that with sparingly soluble salts:The presence of an ion in solution which is common to the electrolyte will decrease the solubility

(adding KF to a solution of PbF2 reduced the Pb2+ concentration)

The presence of an ion in solution which is common to the weak acid will suppress it’s ionization (decrease the concentration of the free ion)




At equilibrium, [F-] = [H3O+] = 1.8 x 10-2 Mso pH = 1.74

What happens if we add 0.05 M NaF?

initial 0.5 M 0 0.05equilibrium (0.5 - x) M x M 0.05 + x M



initial 0.5 M 0 0.05equilibrium (0.5 - x) M x M 0.05 + x M


So (0.05 + x) x = (0.5 - x) 6.8 x 10-4

x = 6.00 x 10-3 M So [H3O+] = 6.00 x 10-3 M

Thus pH = 2.22 c.f. 1.74 without the NaF

Acids and Bases - End of Lecture 1

After studying this lecture and the set problems, you should be able to:

• Understand strong and weak electrolytes, Kw and the autoionization equilibrium in water, definition of the pH scale and its relationship to [H3O

+] and [OH-].

• Know the Arrhenius and Brønsted definitions of acids and bases.

• Understand and know examples of conjugate acid-base pairs.

• Be familiar with the ionization reactions for strong and weak acids, know examples of typical monoprotic, diprotic and triprotic acids.

• Understand the common ion effect in acid ionization.

Objectives Covered in Lecture 1

general chemistry ii 2302102

Documents

equilibrium h3o oh

brnsted acids

bases interaction of

base species

acid species

phdefining acids

triprotic acids

h3o acidbase reactionin