Acid and Base Equilibria

Acid and Base Equilibria

Memorize – strong acids and bases

Definitions Arrhenius: acids produce H+ ions in water and

bases produce OH- ions. Bronsted-Lowry: an acid is a proton donor and a

base is a proton acceptor.HC2H3O2 + H2O C2H3O2

- + H3O+

acid base base acid

NH3 + H2O NH4+ + OH-

base acid acid base

Species that differ by a proton, like H2O and H3O+, are called conjugate acid-base pairs

The reaction of HCl and H2O. HCl is the acid because it donates a proton. Water is the base because it accepts a proton.

(a) Formic acid transfers a proton to a water molecule. HCHO2 is the acid and H2O is the base. (b) When a hydronium ion transfers a proton to the CHO2

- ion, H3O+ is the acid and formate ion is the base.

Conjugate Acid/Base pairs

Acid/BaseHC2H3O2/C2H3O2

-

NH4+/NH3

H3O+/H2O

H2O/OH-

H2O is amphoteric because it can act as either an acid or a base.

OHaqCOHaqOHaqHCO

OHaqCOaqOHaqHCO

23233

2233

)()()(

:base a As)()()(

:acidan As

An amphoteric substances can act as either an acid or baseFor example, the hydrogen carbonate ion:

The strength of an acid is a measure of its ability to transfer a proton

Acids that dissociate completely with water (like HCl and HNO3) are classified as strong

Acids that are less than completely ionized are called weak acids

Bases can be classified in a similar fashion.

Strong and Weak Acids and Bases

Acetic acid (HC2H3O2) is a weak acid It ionizes only slightly in water

The hydronium ion is a better proton donor than acetic acid (it is a stronger acid)

The acetate ion is a better proton acceptor than water (it is a stronger base)

The position of an acid-base equilibrium favors the weaker acid and base (the reactants are favored in this example)

basestronger acidstronger baseaker we acid weaker

23232232 )()( )( aqOHCaqOHOHaqOHHC

increases from left to right within the same period

For example, PH3 < H2S < HCl

increase from top to bottom within a group

For example, HCl < HBr < HI

The strengths of the binary acids

Strength of All Acids1. The POLARITY of the X-H bond The greater the polarity of the X-H bond the

greater the strength of the acid.  Polarity is measured by the difference in electronegativity between the bonded atoms.

When an acid dissociates in water, the X-H bond is broken.  The greater the polarity of the bond, the easier it is to break and produce H+ ions, and thus the stronger the acid.

Formula     Ka value      EN HF 7.2 x 10-4 1.8 H2O 1.8 x 10-16 1.2 NH3 1 x 10-33 0.8 CH4 1 x 10-49 0.4

Explain the difference in the Ka Values

2.  The CHARGE on the acid or base Compounds becomes less acidic and

more basic as the negative charge increases.

It is easier to remove a positive ion (H+) from a neutral atom or molecule than a negatively charged one.  

Bases (H+ acceptors) become stronger as their negative charge increases because they have a stronger force of attraction for pulling in extra H+ ions.

Compounds becomes less acidic and more basic as the negative charge increases.

Formula pH H3PO4 1.5 H2PO4

- 4.4 HPO4

2- 9.3 PO4

3- 12.0

Oxyacids

When the polarity, size, and charge of two compounds are all the same (e.g. oxyacids of the same element) we must find another way to measure the relative strengths of these acids.

Trends in oxoacids (acids of hydrogen, oxygen, and one other element)

Oxyacid Oxyacid - An acid in which the acid hydrogen

atoms are attached to an oxygen atom

Examples of oxyacids of the same element: H2SO4 and H2SO3

HNO3 and HNO2

3. Oxidation State As the oxidation state of an atom increases, its

tendency to draw electrons increases.   In an oxyacid, the central atom pulls electrons

away from the oxygen, consequently making the oxygen more electronegative.  

The O-H bond, therefore becomes more polar, making it easier to form ions and thus increasing the strength of the acid.

As oxidation state increases so does the acidity of the oxyacid.

Oxyacid      Ka value     Oxidation # of ClHClO 2.9 x 10-8

HClO2 1.1 x 10-2

HClO3 5.0 x 102

HClO4 1 x 103

+1+3

+5+7

When the central atom holds the same number of oxygen atoms, the trend is the same as for binary acids across a period, but the reverse for down a column.

Acid strength: HClO4 > HBrO4 > HIO4

Acid strength: HClO4 > H2SO4 > H3PO4

For a given central atom, the acid strength of an oxoacid increases with the number of oxygens held by the central atom

Acid strength: H2SO4 > H2SO3

It is a further generalization, or broadening, of the species that can be classified as either an acid or base

The definitions are based on electron pairs and are called Lewis acids and bases

There is a third definition for acid and bases

Definitions

Lewis acid - accepts a pair of electrons to form a coordinate covalent bond.

Lewis base – donate a pair of electrons

N:H

H

H

+ B

Cl

ClCl

N B

H

H

H

Cl

Cl

Cl

NH3 (a Lewis base) forms a coordinate covalent bond with BF3 (a Lewis acid) during neutralization. NH3BF3 is called an addition compound because it was made by joining two smaller molecules.

Carbon dioxide (Lewis acid) reacts with hydroxide ion (Lewis base) in solution to form the bicarbonate ion. The electrons in the coordinate covalent bond come from the oxygen atom in the hydroxide ion.

Lewis acids:Molecules or ions with incomplete valence

shells (for example BF3 or H+)Molecules or ions with complete valence

shells, but with multiple bonds that can be shifted to make room for more electrons (for example CO2)

Molecules or ions that have central atoms capable of holding additional electrons (usually, atoms of elements in Period 3 and below, for example SO2)

Lewis bases:Molecules or ions that have unshared pairs

of electrons and that have complete shells (for example O2- or NH3)

All Brønsted acids and bases are Lewis acids and bases, just like all Arrhenius acids and bases are Brønsted acids and bases

Neutral solutions: [H+] = [OH-] Acidic solutions: [H+] > [OH-] Basic solutions: [H+] < [OH-] To make the comparison of small values

of [H+] easier, the pH was defined:

In terms of the pH: Neutral solutions: pH = 7.00 Acidic solutions: pH < 7.00 Basic solutions: pH > 7.00

pH10 ][or ]log[ pH -HH

The pH of some common solutions. [H+] decreases, while [OH-] increases, from top to bottom.

The pH of a solution can be measured with a pH meter or estimated using a visual acid-base indicator

An acid-base indicator is a species that changes color based on the pH

Calculating the pH of a strong acid or base is “easy” because they are 100% dissociated in aqueous

For example, the pH of 0.10 M HCl is 1.00 and the pH of 0.10 M NaOH is 13.00

In the last example it was assumed that the total concentration of [H+] was due to the strong acid (HCl) and [OH-] was due to the strong base (NaOH)

This assumption is valid because the autoionization of water is suppressed in strongly acidic or strongly basic solutionsThis assumption fails for very dilute solutions

of acids or bases (less than 10-6 M)

Equilibrium Constant Expression

The equilibrium or ionization constants for weak acids and bases and water are given the labels of Ka, Kb and Kw.

1) A weak acid:HA + H2O H3O+ + A-

Ka = [H+][A-]

[HA]

Equilibrium Constant Expressions

2. A weak base:B + H2O BH+ + OH-

Kb = [BH+][OH-]

[B]

3. Water dissociation:H20 H+ + OH-

Kw = [H+][OH-]

(1 x 10-7)(1x 10-7)= 1 x 10-14

Equilibrium Constant Relationships The product of the Ka and Kb for an acid and its conjugate

base is the Kw

KaKb = Kw = 1 x 10-14

The greater the Ka or Kb the greater the dissociation of the acid or base

• Ka and Kb values are usually very small since they are weak acids and bases.

• pKa and pKb are used to show the equilibrium concentrations.

w

ba

b

a

KOHOHAOHHA

HAAOHKK

AOHHAKOHHAOHA

HAAOHKAOHOHHA

]][[ ][

]][[][

]][[

isproduct the][

]][[

:base conjugate for the][

]][[

:acid weak For the

3

3

2

332

Ionization Constant Relationships pKb = -log Kb pKa = -log Ka pKw = -log Kw = -log(1 x 10-14)

pKw = 14So, pKa + pKb = 14And, pH + pOH = 14 If you know ka for a weak acid you can always

find kb for its conjugate base and visa-versa.

Example: Morphine is very effective at relieving intense pain and is a weak base. What is the Kb, pKb, and percentage ionization of morphine if a 0.010 M solution has a pH of 10.10?

At equilibrium, [OH-] = x = 10-pOH

x x x -0.010 Ex x x - C

x-0.010x 0 0 0.010 I

][]][[ )()(

2

2

B

OHBHKOHaqBHOHaqB b

SOLUTION: Use pOH = 14.00 – pH, substituting:

%3.1

%100010.0

ionization %

and 5.80,p so 106.1

)103.1010.0()103.1(

010.0

then ,103.1

10][

6

4

242

4

)10.1000.14(

xK

xxK

M

OH

b

b

Ionization Constant Calculations1. A monoprotic acid solution has a concentration

of 0.100 M and the pH is 2.44 @ 25oC. Calculate the Ka and pKa.

Ka = 1.36 x 10-4, pKa = 3.86

2. Calculate the pH of a 0.010 M solution of HCl.pH = 2.0

3. Hydrazine, N2H4 has a concentration of 0.025 M. Calculate the pH and % ionization. The Kb = 1.7 x 10-5

pH = 10.81, % ionization = 2.6 %