Acid–Base Concepts

 Acids and bases were first characterized by taste and other simple properties: acids are sour and bases are bitter. There are three descriptions (concepts):

Arrhenius:
- acids increase the concentration of H+  ions (protons), and bases increase the concentration of OH-  (hydroxide) ions in solution. In water, H+ ions are hydrated, and are often denoted as the H3O+ (hydronium) ion.
Shortcomings: cannot be used for nonaqueous phases or where structural differences exist. For example, the following acid–base reaction cannot be described using the Arrhenius definitions: 
HCl(g)+ NH3(g) --> NH4Cl(s)
Brønsted–Lowry:
- acid is a proton donor and a base is a proton acceptor.

Example:

           acid                           base

A conjugate acid–base pair consists of two species in an acid–base reaction, one acid and one base, that differ by the loss or gain of a proton.
An amphiprotic species is a species that can act as either an acid or a base (it can lose or gain a proton), depending on the other reactant. For example, HCO3- acts as an acid in the presence of OH- but as a base in the presence of HF. Anions with ionizable hydrogens, such as HCO3-, and certain solvents, such as water, are amphiprotic.

Identifying Acid and Base Species In the following equations, label each species as an acid or a base. Show the conjugate acid–base pairs.
  (a) HCO3-(aq) + HF(aq) --> H2CO3(aq) + F-(aq)
  (b) HCO
3-(aq) + OH-(aq) --> CO32-(aq) + H2O(l)

Problem Strategy Since a Brønsted–Lowry acid is a proton donor, and the base is a proton acceptor. Examine each equation to find the proton donor on each side. Then label the acids and bases.
In the reaction a) H2CO3 and HCO3- are a conjugate acid­base pair, as are HF and F-.
In the reaction b) HCO3- and CO32- are a conjugate acid­base pair, as are H2O and OH-

Shortcomings: cannot be used for acid–base reaction without a proton. For example: 
Na2O(s) + SO3(g) --> Na2SO4(s)

Lewis:
- acid is an electron pair acceptor, and a base is an electron pair donor. 

The formation of complex ions can also be looked at as Lewis acid­base reactions. Usually when it is not empasized, the acid-base description is viewed in terms of Brønsted–Lowry. Lewis acids and bases, on the other hand, are named as Lewis acid or Lewis base.

Identifying Lewis Acid and Base Species In the following reactions, identify the Lewis acid and the Lewis base.
  (a) Ag+ + 2NH3 -->Ag(NH3)2+
  (b) B(OH)
3 + H2O --> B(OH)4- + H+

Problem Strategy Write the equations using Lewis electron-dot formulas. Then identify the electron-pair acceptor, or Lewis acid, and the electron-pair donor, or Lewis base.

Acid–Base Strength

Strong Bronsted-Lowry acids readily ionize releasing H+ ions. As defined for electrolytes, strong acids and bases ionize nearly completely in water. A combination of molecular structure and individual bond polarity determines the strength of an acid or base. An acid–base reaction normally goes in the direction of the weaker acid.

HCl(aq) + H2O(l) --> Cl-(aq) + H3O+(aq)
stronger     stronger                 weaker       weaker  
acid           base                       base           acid  

    A strong acid reacts with water to form a strong acid and a weak base. The Acid–Base pairs on opposite sides of the equation are called conjugate Acid–Base pairs. Recognizing strong and weak acids helps predict if an Acid–Base reaction will occur. It is important to understand that the terms stronger and weaker are used here only in a comparative sense. The H3O+ ion is a relatively strong acid.

Relative Strengths of Acids and Bases

                         Acid                       Base                        
Strongest acids



























Weakest acids

HClO4  
H2SO4  
HI  
HBr  
HCl  
HNO3  
H3O 
HSO4 
H2SO3  
H3PO4  
HNO2  
HF  
HC2H3O2  
Al(H2O)63+  
H2CO3  
H2S  
HClO  
HBrO  
NH4 
HCN  
HCO3 
H2O2  
HS 
H2 
ClO4 
HSO4 
I 
Br 
Cl 
NO3 
H2O  
SO42-  
HSO3 
H2PO4 
NO2-  
F 
C2H3O2 
Al(H2O)5OH2+  
HCO3 
HS  
ClO 
BrO 
NH3  
CN 
CO32-  
HO2 
S2-  
OH 
Weakest bases



























Strongest bases

Example: For the following reaction, decide which species (reactants or products) are favored at the completion of the reaction.
SO42-(aq) + HCN(aq) --> HSO4-(aq) + CN-(aq)
Use the table to compare the relative strengths of acids and bases. If you compare the relative strengths of the two acids HCN and HSO4-, you see that HCN is weaker. Or, comparing the bases SO42- and CN-, you see that SO42- is weaker. Hence, the reaction would normally go from right to left.

SO42-(aq) + HCN(aq) <-- HSO4-(aq) + CN-(aq)
  weaker             weaker                      stronger         stronger  
base                 acid                             acid                base  

Conclusion: The reactants are favored.

Self-Ionization of Water

Although it is considered a nonelectrolyte, water ionizes slightly according to the reaction: 
H2O(l) + H2O(l) --> H3O+(aq) + OH-(aq)

     The extent of ionization can be expressed as an equilibrium constant: 

Kc =

[H3O+][OH-]


[H2O]2

The equation can be rewritten as Kw, the ion product constant for water:
  [H2O]2 Kc = [H3O+][OH-] = Kw = 1.0 X 10-14 at 25°C
In pure water both H+ and OH- ions are present in equal concentration. From Kw = 1 x 10-14, it follows that [H+]= [OH-]= 1.0 x 10-7 M in pure water. If the concentration of H+ increases with addition of a Brønsted acid, the concentration of OH- must decrease to maintain Kw and vice versa.

The measure of acidity of a solution is often given in pH units. pH is calculated from the concentration of H+ as

pH = -log [H+]

where [H+] is in molar. A neutral solution will have a pH of 7, acidic solutions will have a pH below 7, and basic solutions will have a pH above 7.

     We can summarize [H+] and [OH-] for aqueous solutions: 
      Basic: [H+] < 1.0 X 10-7     [OH-] > 1.0 X 10-7 pH > 7
      Neutral: [H
+] = [OH-] = 1.0 X 10-7 pH ~ 7
      Acidic: [H
+] > 1.0 X 10-7     [OH-] < 1.0 X 10-7  pH < 7

We can also define pOH as
pOH = -log [OH-]and pKw as pKw = -log Kw = -log 1.00 x 10-14 = 14.00 at 25 oC. pH and pOH are related to Kw by:   pH + pOH = 14.00   
  The pH of a solution is typically measured using
pH meters or indicator solutions. 

Molecular Structure and Acid Strength

The strength of an acid depends on how easily the proton, H+, is lost or removed from an H––X bond in the acid species. Following factors are important in determining relative acid strengths: the strength of H––X bond, the strength of H––O bond in H3O+, and the extent to which the conjugate base, X-, of the acid is hydrated in water.

One of the determining factors is the polarity of the bond to which the H atom is attached. The H atom should have a positive partial charge:

d+    d-
H___X

The more polarized the bond is in this direction, the more easily the proton is removed and the greater the acid strength. The second factor determining acid strength is the strength of the bond that is, how tightly the proton is held. This, in turn, depends on the size of atom X. The larger atom X, the weaker is the bond and the greater the acid strength.

1) In going down a column of elements of the periodic table, the size of atom X increases, the H___X bond strength decreases, and the strength of the binary acid increases. You can predict the following order of acid strength in a group:

HF < HCl < HBr < HI

2) Going across a row of elements of the periodic table, the electronegativity increases, the H___X bond polarity increases, and the acid strength increases. For example, the binary acids of the last two elements of the second period are H2O and HF. The acid strengths are

H2O < HF

3) Oxoacids - electronegativity. An oxoacid has the structure

H––O––Y––

The acidic H atom is always attached to an O atom, which, in turn, is attached to an atom Y. Other groups, such as O atoms or O––H groups, may also be attached to Y. Bond polarity dominates in determining relative strengths of the oxoacids. This, in turn, depends on the electronegativity of atom Y. If the electronegativity of atom Y is large, the H___O bond is relatively polar and the acid strength large. For a series of oxoacids of the same structure, differing only in the atom Y, the acid strength increases with the electronegativity of Y:

HIO < HBrO < HClO

The oxoacids of chlorine provide another example where oxidation state (and thus electronegativity) of Cl increases with each additional O atom. As a result, the H atom becomes more acidic. The acid strengths increases in the following order:

HClO < HClO2 < HClO3 < HClO4

4) Polyprotic acids. For example, H2SO4 ionizes by losing a proton to give HSO4-, which, in turn, ionizes to give SO42-. HSO4- can lose a proton, so it is acidic. However, because of the negative charge of the ion, which tends to attract protons, its acid strength is reduced from that of the uncharged species. That is, the acid strengths are in the order

HSO4- < H2SO4

This shows that the acid strength of a polyprotic acid and its anions decreases with increasing negative charge