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Chapter 14 Acids and Bases. Some Definitions. Arrhenius Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions. Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions. Some Definitions. Brønsted – Lowry

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Chapter 14 acids and bases

Chapter 14Acids and Bases


Some definitions
Some Definitions

  • Arrhenius

    • Acid: Substance that, when dissolved in water, increases the concentration of hydrogen ions.

    • Base: Substance that, when dissolved in water, increases the concentration of hydroxide ions.


Some definitions1
Some Definitions

  • Brønsted–Lowry

    • Acid: Proton donor

    • Base: Proton acceptor


A Brønsted–Lowry acid…

…must have a removable (acidic) proton.

A Brønsted–Lowry base…

…must have a pair of nonbonding electrons.


If it can be either
If it can be either…

...it is amphiprotic.

HCO3−

HSO4−

H2O


What happens when an acid dissolves in water
What Happens When an Acid Dissolves in Water?

  • Water acts as a Brønsted–Lowry base and abstracts a proton (H+) from the acid.

  • As a result, the conjugate base of the acid and a hydronium ion are formed.


Acid dissociation ionization reactions
Acid Dissociation (Ionization) Reactions

  • Write the simple dissociation (ionization) reaction (omitting water) for each of the following acids:

    • a. Hydrochloric acid

    • b. Acetic acid

    • c. Ammonium ion

    • d. Anilinium ion (C6H5NH3)

    • e. Hydrated aluminum (III) ion [Al(H2O)6]3+


Conjugate acids and bases
Conjugate Acids and Bases:

  • From the Latin word conjugare, meaning “to join together.”

  • Reactions between acids and bases always yield their conjugate bases and acids.


Acid and base strength
Acid and Base Strength

  • Strong acids are completely dissociated in water.

    • Their conjugate bases are quite weak.

  • Weak acids only dissociate partially in water.

    • Their conjugate bases are weak bases.


Acid and base strength1
Acid and Base Strength

  • Substances with negligible acidity do not dissociate in water.

    • Their conjugate bases are exceedingly strong.


Acid and base strength2
Acid and Base Strength

In any acid-base reaction, the equilibrium will favor the reaction that moves the proton to the stronger base.

HCl(aq) + H2O(l) H3O+(aq) + Cl−(aq)

H2O is a much stronger base than Cl−, so the equilibrium lies so far to the right K is not measured (K>>1).


Acid and base strength3

C2H3O2(aq) + H2O(l)

H3O+(aq) + C2H3O2−(aq)

Acid and Base Strength

Acetate is a stronger base than H2O, so the equilibrium favors the left side (K<1).


Relative base strength
Relative Base Strength

Using the following Ka values, arrange the following species according to their strength as bases:


Autoionization of water

H2O(l) + H2O(l)

H3O+(aq) + OH−(aq)

Autoionization of Water

  • As we have seen, water is amphoteric.

  • In pure water, a few molecules act as bases and a few act as acids.

  • This is referred to as autoionization.


Ion product constant
Ion-Product Constant

  • The equilibrium expression for this process is

    Kc = [H3O+] [OH−]

  • This special equilibrium constant is referred to as the ion-product constant for water, Kw.

  • At 25°C, Kw = 1.0  10−14


Calculating h oh
Calculating [H+] & [OH-]

  • Calculate [H+] & [OH-] as required for each of the following solutions at 250C, & state whether the solution is neutral, acidic, or basic.

    • a. = 1.0 x 10-5 M OH-

    • b. a. = 1.0 x 10-7 M OH-

    • c. 10.0 M H+


pH

pH is defined as the negative base-10 logarithm of the hydronium ion concentration.

pH = −log [H3O+]


pH

  • In pure water,

    Kw = [H3O+] [OH−] = 1.0  10−14

  • Because in pure water [H3O+] = [OH−],

    [H3O+] = (1.0  10−14)1/2 = 1.0  10−7


pH

  • Therefore, in pure water,

    pH = −log (1.0  10−7) = 7.00

  • An acid has a higher [H3O+] than pure water, so its pH is <7

  • A base has a lower [H3O+] than pure water, so its pH is >7.


pH

These are the pH values for several common substances.


Other p scales
Other “p” Scales

  • The “p” in pH tells us to take the negative log of the quantity (in this case, hydrogen ions).

  • Some similar examples are

    • pOH −log [OH−]

    • pKw−log Kw


Watch this
Watch This!

Because

[H3O+] [OH−] = Kw = 1.0  10−14,

we know that

−log [H3O+] + −log [OH−] = −log Kw = 14.00

or, in other words,

pH + pOH = pKw = 14.00


Calculating ph poh
Calculating pH, pOH

pH = -log10(H3O+)

pOH = -log10(OH-)

Relationship between pH and pOH

pH + pOH = 14

Finding [H3O+], [OH-] from pH, pOH

[H3O+] = 10-pH

[OH-] = 10-pOH


Calculating ph poh1
Calculating pH & pOH

  • Calculate pH & pOH for each of the following solutions at 250C.

    • a. = 1.0 x 10-3 M OH-

    • b. a. = 1.0 M H+


Calculating ph
Calculating pH

  • The pH a sample of human blood was measured to be 7.41 at 250C. Calculate pOH, [H+], & [OH-] for the sample.


How do we measure ph
How Do We Measure pH?

  • For less accurate measurements, one can use

    • Litmus paper

      • “Red” paper turns blue above ~pH = 8

      • “Blue” paper turns red below ~pH = 5

    • An indicator


How do we measure ph1
How Do We Measure pH?

For more accurate measurements, one uses a pH meter, which measures the voltage in the solution.


Strong acids
Strong Acids

  • You will recall that the seven strong acids are HCl, HBr, HI, HNO3, H2SO4, HClO3, and HClO4.

  • These are, by definition, strong electrolytes and exist totally as ions in aqueous solution.

  • For the monoprotic strong acids,

    [H3O+] = [acid].


Ph of strong acids
pH of Strong Acids

  • Calculate pH of 0.10 M HNO3.

  • Calculate pH of 1.0 x 10-10 M HCl.


Dissociation constants

Kc =

[H3O+] [A−]

[HA]

HA(aq) + H2O(l)

A−(aq) + H3O+(aq)

Dissociation Constants

  • For a generalized acid dissociation,

    the equilibrium expression would be

  • This equilibrium constant is called the acid-dissociation constant, Ka.


Dissociation constants1
Dissociation Constants

The greater the value of Ka, the stronger the acid.


Calculating k a from the ph

[H3O+] [COO−]

[HCOOH]

Ka =

Calculating Ka from the pH

  • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

  • We know that


Calculating k a from the ph1
Calculating Ka from the pH

  • The pH of a 0.10 M solution of formic acid, HCOOH, at 25°C is 2.38. Calculate Ka for formic acid at this temperature.

  • To calculate Ka, we need the equilibrium concentrations of all three things.

  • We can find [H3O+], which is the same as [HCOO−], from the pH.


Calculating k a from the ph2
Calculating Ka from the pH

pH = −log [H3O+]

2.38 = −log [H3O+]

−2.38 = log [H3O+]

10−2.38 = 10log [H3O+] = [H3O+]

4.2  10−3 = [H3O+] = [HCOO−]


Calculating k a from ph
Calculating Ka from pH

Now we can set up a table…


Calculating k a from ph1

[4.2  10−3] [4.2  10−3]

[0.10]

Ka =

Calculating Ka from pH

= 1.8  10−4


Solving weak acid equilibrium problems
Solving Weak Acid Equilibrium Problems

  • 1. List the major species in the solution.

  • 2. Choose the species that can produce H+ and write balanced equations for the reactions producing H+ .

  • 3. Using the values of the equilibrium constants for the reactions you have written, decide which equilibrium will dominate in producing H+ .

  • 4. Write the equilibrium expression for the dominant equilibrium

  • 5. ICE the problem

  • 6. Substitute the equilibrium [ ] into the equilibrium expression

  • 7. Solve for x the “easy” way; that is, by assuming [HA]0-x  [HA]0

  • 8. Use the 5% rule to verify whether the approximation is valid

  • 9. Calculate [H+] and pH


The ph of weak acids
The pH of Weak Acids

  • The hypochlorite ion (OCl-) is a strong oxidizing agent often found in household bleaches & disinfectants. It is also the active ingredient that forms when swimming pool water is treated with chlorine. In addition to its oxidizing abilities, the hypochlorite ion has a relatively high affinity for protons (it is a much stronger base than Cl-for example) & forms the weakly acidic hypochlorous acid (HOCl, Ka = 3.5 x 10-8 ). Calculate pH of 0.100 M aqueous solution of hypochlorous acid.


The ph of weak acids continued
The pH of Weak Acids continued

  • HOCl, Ka = 3.5 x 10-8 Calculate pH of 0.100 M aqueous solution of hypochlorous acid.


The ph of weak acid mixtures
The pH of Weak Acid Mixtures

  • Calculate the pH of a solution that contains 1.00 M HCN (Ka = 6.2 x 10-10) and 5.00 M HNO2 (Ka = 4.0 x 10-4) . Also calculate the concentration of cyanide ion in this solution at equilibrium.


Percent dissociation
Percent Dissociation

  • In general, the more dilute the weak acid solution, the greater the percent dissociation of the weak acid.


Calculate percent dissociation
Calculate Percent Dissociation

  • Calculate the percent dissociation of acetic acid (Ka = 1.8x 10-5) in each of the following solutions.

  • a. 1.00 M HC2H3O2

  • b. 0.100 M HC2H3O2


Calculation ka from percent dissociation
Calculation Ka from Percent Dissociation

  • Lactic acid (HC3H5O3) is a waste product that accumulates in muscle tissue during exertion, leading to pain & feeling of fatigue. In a 0.100 M aqueous solution, lactic acid is 3.7% dissociated. Calculate the value of Ka for this acid.


Strong bases
Strong Bases

  • Strong bases are the soluble hydroxides, which are the alkali metal and heavier alkaline earth metal hydroxides (Ca2+, Sr2+, and Ba2+).

  • Again, these substances dissociate completely in aqueous solution.


The ph of strong bases
The pH of strong bases

  • Calculate the pH of a 5.0x 10-2 M NaOH solution.


Weak bases
Weak Bases

  • Many types of proton acceptors (bases) do not contain hydroxide ions. When dissolved in water, they increase the concentration of hydroxide ions because of their reaction with water.

  • Ex.

  • Bases such as ammonia typically have at least one unshared pair of electrons that is capable of forming a bond with a proton.


Weak bases1
Weak Bases

  • We will solve weak base problems in the same manner we solved weak acid problems (look back over the steps you were given

  • We will use Kb instead of Ka and will find [OH-] instead of [H+]

    • Remember the process for “switching” from pOH to pH &/or from [OH-] [H+]


Ph of a weak base i
pH of a Weak Base (I)

  • Calculate the pH for a 15.0 M NH3 solution (Kb = 1.8x 10-5).


Ph of a weak base ii
pH of a Weak Base (II)

  • Calculate the pH for a 1.0 M methylamine solution (Kb = 4.38x 10-4).


Polyprotic acids
Polyprotic Acids

  • Some acids furnish more than one acidic proton such as H2SO4, H3PO4.

    • Ex. H2CO3


Ph of a polyprotic acid
pH of a polyprotic acid

  • Calculate the pH of a 5.0 M H3PO4 solution and the equilibrium concentrations of the species H3PO4, H2PO4-, HPO42-, & PO43- .


Ph of sulfuric acid
pH of sulfuric acid

  • Calculate the pH of a 1.0 M H2SO4 solution.

  • Calculate the pH of a .0100 M H2SO4 solution.



Acid base properties of salts1
Acid-Base Properties of Salts

These salts simply dissociate in water:

KCl(s)  K+(aq) + Cl-(aq)


Acid base properties of salts2
Acid-Base Properties of Salts

The basic anion can accept a proton from water:

C2H3O2- + H2O  HC2H3O2 + OH- base acid acid base


Acid base properties of salts3
Acid-Base Properties of Salts

The acidic cation can act as a proton donor:

NH4+(aq)  NH3(aq) + H+(aq)

Acid Conjugate Proton

base


Acid base properties of salts4
Acid-Base Properties of Salts

  • IF Ka for the acidic ion is greater than Kb for the basic ion, the solution is acidic

  • IF Kb for the basic ion is greater than Ka for the acidic ion, the solution is basic

  • IF Kb for the basic ion is equal to Ka for the acidic ion, the solution is neutral


Acid base properties of salts5
Acid-Base Properties of Salts

Step #1:

AlCl3(s) + 6H2O  Al(H2O)63+(aq) + Cl-(aq)

Salt water Complex ion anion

Step #2:

Al(H2O)63+(aq)  Al(OH)(H2O)52+(aq) + H+(aq)

Acid Conjugate base Proton


Salts as weak bases
Salts as Weak Bases

Calculate the pH of a 0.30 M NaF solution. The Ka value for HF is 7.2 x 10-4


Salts as weak acids i
Salts as Weak Acids I

Calculate the pH of a 0.10 M NH4Cl solution. The Kb value for NH3 is 1.8 x 10-5


Salts as weak acids ii
Salts as Weak Acids II

Calculate the pH of a 0.010 M AlCl3solution. The Kavalue for Al(H2O)63+ is 1.4 x 10-5


The acid base properties of salts
The Acid/Base Properties of Salts

Predict whether an aqueous solution of each of the following salts will be acidic, basic, or neutral.

NH4C2H3O2

NH4CN

Al2(SO4)3


Effect of structure on acid base properties
Effect of Structure on Acid-Base Properties

Hypochlorous

acid

Chlorous

acid

Chloric

acid

Perchloric

acid

Increasing Acidity


The lewis acid base model
The Lewis Acid-Base Model

  • Lewis Acid: An electron pair acceptor

    • Has an empty orbital to accept a pair of electrons

  • Lewis Base: An electron pair donor

    • Has a lone pair of electrons


Lewis acids bases
Lewis Acids & Bases

  • For each reaction, identify the Lewis acid & base.

  • Ni2+(aq) + 6NH3(aq) Ni(NH3)62+(aq)

  • H+(aq) + H2O(l) H3O+(aq)


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