Chapter 17: Acids and Bases• Acid-base reactions involve proton

(hydrogen ion, H+) transfer

• The generalization of the Arrhenius definition of acids and bases is called the Brønsted-Lowry definitions:

• An acid is a proton donor

• A base is a proton acceptor

• This allows for gas phase acid-base reactions

• 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.

• An amphoteric substances can act as either an acid or base

• These are also called amphiprotic, and can be either molecules or ions– For example, the hydrogen carbonate ion:

OHaqCOHaqOHaqHCO

OHaqCOaqOHaqHCO

2233

2233

)()()(

:base a As

)()()(

:acidan As

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

• Acids that react 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 bases, like the oxide ion, react

completely– Weak bases, like NH3, undergo incomplete

reactions

• The strongest acid in water is the hydronium ion

• If a more powerful proton donor is added to water, it quantitatively reacts with water to produce H3O+

• Similarly, the strongest base that can be found in water is the hydroxide ion, because more powerful proton acceptors react quantitatively with water to produce OH-

• 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

basestronger acidstronger baseaker we acid weaker

23232232 )()( )( aqOHCaqOHOHaqOHHC

• This can be generalized:– Stronger acids and bases tend to react with each

other to produce their weaker conjugates– The stronger a Brønsted acid is, the weaker is

its conjugate base– The weaker a Brønsted acid is, the stronger is

its conjugate base

• These ideas can be applied to the binary acids (acids made from hydrogen and one other element)

• The strengths of the binary acids increases from left to right within the same period

• For example, HCl is stronger acid than H2S which is a stronger acid than PH3

• The strengths of the binary acids increase from top to bottom within a group

• For example, HI is a stronger acid than HBr which is a stronger acid than HCl

• Trends are also present in the oxoacids (acids of hydrogen, oxygen, and one other element)

• When the central atom holds the same number of oxygen atoms, the acid strength increases from the bottom to top within a group and from left to right within a period

• 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

• The strength of an acid can be analyzed in terms the the basicity of the anion formed during the ionization

• The basicity is the willingness of the anion to accept a proton from the hydronium ion

• Consider H2SO4 and H3PO4:

acid phosphoric acid sulfuric

||

|

||

||OH

O

OH

PHOOH

O

O

SHO

• The anions are:

• In oxoanions, the lone oxygens carry most of the negative charge, making the hydrogen sulfate ion a weaker base than the hydrogen phosphate ion

O) lone (2 O) lone (3

424

||

|

||

||

POHHSO

O

O

OH

PHOO

O

O

SHO

• In terms of the percentage of molecules that are ionized, sulfuric acid is a stronger acid than phosphoric acid

• There is a third definition for acid and bases

• 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

• A Lewis acid is any ionic or molecular species that can accept a pair of electrons in the formation of a coordinate covalent bond

• A Lewis base is any ionic or molecular species that can donate a pair of electrons in the formation of a coordinate covalent bond

• Neutralization is the formation of a coordinate covalent bond between the donor (base) and acceptor (acid)

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 (a Lewis acid) reacts with hydroxide ion (a 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

• Consider a proton transfer from the Lewis perspective

• For example, the proton transfer between the hydronium ion and ammonia:

4232 NHOHNHHOH

• The elements most likely to form acids are the nonmetals in the upper right-hand corner of the periodic table

• The elements most likely to form basic hydroxides are the IA and IIA metals along the left of the periodic table

• The elements themselves can be classified according to the ability of their oxides to form acids or bases

• In general, most metal oxides react with water to form bases, and nonmetal oxides react with water to form acids

• In Section 5.5 metal oxides were called base anhydrides and nonmetal oxides were called acid anhydrides

• When cations dissolve in water, they form species called hydrated ions

• Hydrated metal ions tend to be Brønsted acids

• For the monohydrate of the metal ion Mn+ the equilibrium can be represented as

OHMOHOHOHM nn3

)1(2 2)(

The metal ion makes the hydrogens on the water more acidic.

• The charge density of a cation is its charge divided by its volume

• The higher the charge density, the better a cation is at drawing electron density from a O-H bond and the more acidic it is

• Within a given period, the cation size increases, and the charge density decreases, from top to bottom

• As a result, the most acidic hydrated cations are found at the top of a group

• As the cation charge increases, it becomes more acidic

• When the charge (oxidation number) is small, its oxide tends to be basic

• When the cation ion charge is +3, the oxide begins to become acidic

• An amphoteric species is capable of acting as both an acid and base

• Aluminum oxide is an example of an amphoteric compound

• Many nonmetal oxides are acid anhydrides– For example:

OHaqAl(aq)H(s)OAl

OAl

OHAlO(aq)OH(s)OAl

OAl-

23

32

32

2232

32

3)(26

:base a as

22

:acidan as

acid carbonic )()(

acid sulfuric )()(

3222

4223

aqCOHOHgCO

aqSOHOHgSO

• Water undergoes self-ionization or autoionization making it a weak electrolyte

• This equilibrium is described by the ion product of water

OHOHOHOH 322

]][[]][[ 3 OHHOHOHKw

• At 25oC in pure water it has been found that

• In any aqueous solution, the product of [H+] and [OH-] equals Kw

• This provides an alternate a way to define the acidity or basicity of a solution

)25(at 100.1

thatso

100.1][][

o14

7

CK

MOHH

w

• Neutral solutions: [H3O+] = [OH-] or [H+] = [OH-]

• Acidic solutions: [H3O+] > [OH-] or [H+] > [OH-]

• Basic solutions: [H3O+] < [OH-] or [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 (see Table 17.3, page 762)

• 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 solutions– This assumption fails for very dilute solutions

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