Chapter 19

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ACID-BASES

Self-Ionization of Water

• Two water molecules produce a hydronium ion & a hydroxide ion by the transfer of a proton.

• H2O (l) + H2O (l) H3O+(aq) + OH- (aq)

• In pure water, every time

you make one H3O+

you get one OH-

• That is, [H3O+] = [OH-]

Hydronium Concentration

• Symbolized

[H3O+]

In Water at 25°C

[H3O+] = 1.0 x 10-7 M

And

[OH-] = 1.0 x 10-7 M

IONIZATION CONSTANT OF WATER at 25°C

• Kw = [H3O+] [OH-]

• Kw = (1.0 x 10-7 M) x (1.0 x 10-7 M)

= 1.0 x 10-14 M2

ION PRODUCT CONSTANT FOR WATER at 25°C

• Kw = [H+] [OH-]

• Kw = (1.0 x 10-7 M) x (1.0 x 10-7 M)

= 1.0 x 10-14 M2

Neutral solutions

• Pure water is Neutral and

[H3O+] = [OH-]

• In any solution that is neutral,

[H3O+] = [OH-]

• In any solution that is acidic,

[H3O+] > [OH-]

• [H3O+] > 1.0 x 10-7 M

• In any solution that is basic,

[H3O+] < [OH-]

• [OH-] > 1.0 x 10-7 M

• [H3O+] < 1.0 x 10-7 M

STRONG ACIDS & BASES

• NEARLY COMPLETELY IONIZE OR DISSOCIATE IN AQUEOUS SOLUTION

STRONG ACID Solutions

HClO4 (aq)

H2SO4 (aq)

HNO3 (aq)HCl (aq)HBr (aq)HI (aq)

MEMORIZE!!!

Strong Bases

Group 1 HydroxidesNaOHKOHLiOHRbOHCsOH

Group 2 HydroxidesCa(OH)2

Ba(OH)2

Sr(OH)2

NaOH Na+ (aq) + OH- (aq)

Therefore,1 mole of NaOH will yield 1 mole of OH- in an aqueous solution.

Ca(OH)2 Ca2+ (aq) + 2OH- (aq)

Therefore,1 mole of Ca(OH)2 will yield 2 mole of OH- in an aqueous solution.

HCl H+ (aq) + Cl- (aq)

Therefore,1 mole of HCl will yield 1 mole of H+

in an aqueous solution.

H2SO4 2H+ (aq) + SO 42- (aq)

Therefore,1 mole of H2SO4 will yield 2 mole of H+

in an aqueous solution.

For any aqueous solution at 25°C

• Kw = [H3O+] [OH-] = 1.0 x 10-14 M2

• Using this equation,

if the concentration of one of the ions is known,

then, the concentration of the other can be calculated.

Exercise

• Determine the hydronium and hydroxide ion concentrations in a solution that is

1 x 10 -4 M HCl.

Answer:

[H3O+] = 1 x 10-4 M

[OH-] = 1 x 10-10 M

Exercise

• Determine the hydronium and hydroxide ion concentrations in a solution that is

1.0 x 10 -3 M HNO3.

Answer:

[H3O+] = 1.0 x 10-3 M

[OH-] = 1.0 x 10-11 M

Exercise

• Determine the hydronium and hydroxide ion concentrations in a solution that is

1.0 x 10 -4 M Ca(OH)2

Answer:

[H3O+] = 5.0 x 10-11 M

[OH-] = 2.0 x 10-4 M

pH

• Convenient way to express numbers that tend to be very small.

• pH = - log[H3O+]

• [H3O+] = 1 x 10-pH

• If [H3O+] = 1 x 10-7

then pH = 7

Figure 14.8The pH Scale and pH Values of Some

Common Substances

pOH

• pOH = - log[OH-]

• If [OH-] = 1 x 10-7

then pOH = 7

• pH + pOH = 14

Calculating pH & pOH

• Determine the pH & pOH of the following solutions.

• 1) 1 x 10-3 M HClpH = 3.0 pOH = 11.0

• 2) 1 x 10-5 M HNO3

pH = 5.0 pOH = 9.0• 3) 1 x 10-4 M NaOH

pH = 10.0 pOH = 4.0

Calculating the pH & pOH

• What is the pH of a solution if the hydronium ion conc. is 6.7 x 10-4 M?

• pH = 3.17• What is the pH of a solution if the

hydronium ion conc. is 2.5 x 10-2 M?• pH = 1.60 pOH = 12.40• Determine the pH of a 2.0 x 10-2 M

Sr(OH)2 solution.• pH = 12.60

Calculating the hydronium and hydroxide ion concentrations

• The pH of a solution is 5.0. What is the hydronium ion concentration?

• 1 x 10-5 M

• The pH of an aqueous solution is measured to be 1.50. Calculate the hydronium ion and hydroxide ion concentrations.

• [H3O+] = 3.2 x 10-2 M [OH-] = 3.2 x 10-13 M

pH Calculations & the Strength of Acids and Bases

• For strong acids and bases, the hydronium and hydroxide ion concentrations can be directly calculated.

• For weak acids and bases, hydronium and hydroxide ion concentrations cannot be directly calculated because not all of the molecules are ionized.

Homework

• Complete worksheet

Titrations

• Technique used to measure the amount of acid or base present in a solution.

• Involves an acid/base neutralization reaction

• Equivalence Point: point at which the two solutions used in the titration are present in chemically equivalent amounts.

• Endpoint: The point in a titration at which an indicator changes color.

Acid-Base Indicator

•. . . marks the end point of a titration by changing color.

•The equivalence point is not necessarily the same as the end point.

Most common acid-base indicators are weak acids, HIn

One color with H+ ( HIn )

Another color without H+ ( In-)

HIn H+ + In-

Red Blue

Figure 15.6The Acid and Base Forms of the Indicator

Phenolphthalein

Figure 15.8

The Useful pH Ranges for Several Common Indicators

Titration (pH) Curve

• A plot of pH of the solution being analyzed as a function of the amount

of titrant added.

• Equivalence point (stoichiometric point): Enough titrant has been added to react exactly with the solution being analyzed.

The pH Curve for the Titration of 100.0 mL of 0.10 M HCI with 0.10 M NaOH

The pH Curve for the Titration of 50 mL of 0.1 M

HC2H3O2 with 0.1 M

NaOH

Titration Technique

• Go to page 615 in textbook

Titration

• Concentration of one solution is known precisely : Standard solution

• Concentration of other solution calculated from the chemically equivalent volumes.

Acid-Base Properties of Salts

Salts = Ionic compounds

Salts can behave as ACIDS or BASES.

1. Salts that produce neutral solutions.

Composed of cations from strong bases and anions from strong acids.

Example: NaCl. NaNO3, KCl

2. Salts that produce basic solutions.

Composed of cations with neutral properties and anions which are the conjugate

base of a weak acid.

Example: NaCH3COOMajor species:

Na+ is neutralCH3COO- is conjugate base of weak acidH2O is weakly amphoteric

CH3COO1- + H2O CH3COOH + OH1-

CH3COO1- in water produces OH1- ions Basic solution

3. Salts that produce acidic solutions.

Composed of cations which are the conjugate acid of a weak base andanions with neutral properties.

Example: NH4ClMajor species: Cl-, H2O, & NH4

+

NH41+ (aq) NH3 (aq) + H1+ (aq)

Section 15.2 A Buffered Solution

•. . . resists change in its pH when either H+ or OH are added.

•1.0 L of 0.50 M H3CCOOH

– + 0.50 M H3CCOONa

• pH = 4.74•Adding 0.010 mol solid NaOH raises the pH of the solution to 4.76, a very minor change.

• One of the most practical applications of buffers is our blood.

• Most important one is the HCO3

- and H2CO3.

BUFFERS

• Weak acid and its SaltHF & NaFOr

• Weak base and its Salt

NH3 & NH4Cl

Key Points on Buffered

Solutions

• 1. They are weak acids or bases containing a common ion.

.

Buffered Solution Characteristics

Buffers contain relatively large amounts of weak acid and corresponding base.

Added OH reacts to completion with the weak acid.

OH- + HA A- + H2O The net result is the hydroxide is

replaced by A-.

Buffered Solution Characteristics

Buffers contain relatively large amounts of weak acid and corresponding base.

Added H+ reacts to completion with the weak base.

H+ + A- HA Free H+ converted to HA

Homework