NATURE OF ACIDS & BASES

General Properties:

• Listed below are some common acids and their common uses:

NATURE OF ACIDS & BASES

• The structural formulas for several common acids are shown below. Note that all acids have a hydrogen bonded to an oxygen (referred to as the acidic hydrogen) that ionizes in aqueous solutions.
  
  
  

• Acetic acid is an organic acid that contains the carboxylic acid group, that is shown below:
  

Carboxylic acid group

• Some common bases are listed below. Organic bases (called alkaloids) are also bitter in taste and are often poisonous.
  

DEFINITION OF ACIDS & BASES

• There are three definitions of acids and bases used in chemistry. Each definition is useful in a given instance, and are discussed next.

Arrhenius Definition:

• Arrhenius definition of acids and bases in based on their behavior in water (aqueous solutions), and was developed by Swedish chemist Svante Arrhenius.
• Based on this definition:
• Acid: A substance that produces ions in aqueous solution.
• Base: A substance that produces ions in aqueous solution.
• Therefore, according to the Arrhenius definition, HCl is an acid because it produces ions in solution:

• The ions produced are very reactive and bond to water molecules in aqueous solution, forming :
  

  

• Similarly, according to the Arrhenius definition, NaOH is a base because it produces in solution.

• Based on Arrhenius definition, acids and bases combine to form water, neutralizing each other in the process:
  

  

DEFINITION OF ACIDS AND BASES

Brønsted-Lowry Definition:

• The Brønsted-Lowry definition of acids and bases focuses on the transfer of ions in an acid-base reaction. Since is a proton, this definition focuses on the idea of proton donor and proton acceptor:
• Acid: proton ( ion) donor
• Base: proton ( ion) acceptor
• According to this definition, HCl is an acid because, in solution, it donates a proton to water:

• According to this definition, is a base because, in solution, it accepts a proton from water:

• According to the Brønsted-Lowry definition, acids (proton donors) and bases (proton acceptors) always occur together in an acid-base reaction. For example:

• According to the Brønsted-Lowry definition, some substances-such as water-can act as acids or bases. These substances are called amphoteric.

Examples:

DEFINITION OF ACIDS AND BASES

• In a reversible acid-base reaction, both forward and reverse reactions involve transfer
  

  

• The pair of acid and base from each side of these reactions are referred to as acid-base conjugate pairs. A conjugate acid is any base to which a proton has been added, while a conjugate base is any acid from which a proton has been removed.
  

  

  
  

  

Examples:

1. Identify the Brønsted-Lowry acid-base conjugate pairs in each of the following equations:
   a)
   b) (aq) (l) (aq) (aq)
2. Which pair is not a conjugate acid-base pair?
   a)
   b)
   c)

LEWIS ACIDS AND BASES

• A third definition of acids and bases is the Lewis model, which further broadens the range of substances that can be considered acids and bases.
• This model, developed by Gilbert Lewis, defined acids and bases in terms of electron-pair transfer.
  

  

• Therefore, based on the Lewis model:
• Lewis Acid: electron pair acceptor
• Lewis Base: electron pair donor
  

  

• The Lewis model does not substantially expand the substances that can be considered a base, since a proton acceptor must have an electron pair to bond the proton.
• However, it does significantly expand the substances that can be considered an acid. According to this model, substances need not even contain hydrogen to be an acid. For example:

LEWIS ACIDS AND BASES

• Since molecules with incomplete octets have empty orbitals, they can serve as Lewis acids. For example, both and have incomplete octets:
  
  
  
• As a result, both these molecules can act as Lewis acids, as shown below:
  

  

  
  

  

• Some molecules that do not initially contain empty orbitals can rearrange their electrons to act as Lewis acids. For example, consider the reaction between water and carbon dioxide:
  

  

LEWIS ACIDS AND BASES

• Some cations, since they are positively charged and have lost electrons, have empty orbitals that allow them to also act as Lewis acids. For example, consider the hydration of ions shown below:
  

  

• In General:

Examples:

1. Classify each species below as a Lewis acid or a Lewis base:
   a)
   b)
   c)
   d)

ACIDS & BASES: LEWIS VS. BRONSTED-LOWRY

• The Bronsted-Lowry and Lewis definitions of acids and bases are complementary and are often used to explore the common ways in which acids and bases are involved in many reactions.
• The general reaction schemes below show the difference between the two definitions.

Bronsted-Lowry

• Note that in the Bronsted-Lowry definition, the acid exchanges proton with the base, and as a result forms charged species as product. In the Lewis definition, however, the product formed is an adduct of the acid and base, and is charge neutral.
• The specific examples below show this difference

Bronsted-Lowry

Lewis

ACID STRENGTH

• The strength of an electrolyte depends on the extent of its dissociation into its component ions in solution. A strong electrolyte completely dissociates into ions in solution, whereas a weak electrolyte partially dissociates into ions.
• The strength of an acid is similarly defined. A strong acid completely ionizes in solution, while a weak acid only partially ionizes. Therefore, the strength of the acid depends on the equilibrium shown below:

• For example, hydrochloric acid (HCl) is classified as a strong acid since its solution contains virtually all ions:

• Listed below are six important strong acids. Note that sulfuric acid ( ) is diprotic, while the others are monoprotic.

ACID STRENGTH

• In contrast, HF is classified as a weak acid because it does not completely dissociate in solution:

• The degree to which an acid is strong or weak depends on the attraction between the anion of the acid (the conjugate base) and the hydrogen ion, relative to the attraction of these ions to water.
• When the attraction between and is weak, the acid is strong. On the other hand, if the attraction between and is strong, the acid is weak.
• For example, in HCl , the has a relatively weak attraction for - the reverse reaction does not occur to any significant extent. In HF , on the other hand, the has a greater attraction for - the reverse reaction occurs to a significant degree.
  

  

• In general, the stronger the acid, the weaker the conjugate base, and the weaker the acid, the stronger the conjugate base.
  

  

• Listed below are some common weak acids. Note that some are diprotic and one is triprotic.

ACID IONIZATION CONSTANT ( )

• The relative strength of a weak acid can be quantified with the acid ionization constant ( ).
• is the equilibrium constant for the ionization reaction of a weak acid. For example,

• Listed below are the acid ionization constants for some monoprotic weak acids, listed in order of decreasing strength:

Acid	Formula	Structural Formula	Ionization Reaction
Chlorous acid
Nitrous acid
Hydrofluoric acid	HF	H-F
Formic acid
Benzoic acid
Acetic acid
Hypochlorous acid	HCIO
Hydrocyanic acid	HCN
Phenol

3. The hydrogen oxalate ion ( ) is amphoteric. Write a balanced equation showing how it reacts as an acid towards water and another equation showing how it reacts as a base towards water.
4. For each acid shown below, write the formula of its conjugate base:
   
5. For each base shown, write the formula of its conjugate acid:
   

AUTOIONIZATION OF WATER

• As discussed earlier, water can act both as an acid or base. It acts as an acid when it reacts with HCl , and it acts as a base when it reacts with :
  

  

• Therefore, water is amphoteric: it can act either as an acid or base. Even when pure, water can act as an acid and a base with itself, a process called autoionization:

Acid (proton donor)

• The autoionization reaction and its equilibrium constant can be written as:

• This equilibrium constant is called ion product constant for water ( ), and has a value of at .
• In pure water, . Therefore

ACIDIC & BASIC SOLUTIONS

• An acidic solution contains an acid that creates additional ions, causing to increase. Since the value remains constant in any solution, the [ ] in these solutions can be calculated as shown:

• For example, if , then the can be calculated as:

• A basic solution contains a base that creates additional ions, causing to increase. Since the value remains constant in any solution, the in these solutions can be calculated as shown:

• For example, if , then the can be calculated as:

SUMMARY

• In a neutral solution:
• In an acidic solution:

• In a basic solution:

Examples:

1. What is the concentration of the [ ] and [ ] in a 0.25 M solution of ?

2. What is the concentration of the and in a 0.040 M solution of ?
3. Identify each of the following solutions are acidic, basic or neutral:

pH SCALE

• The pH scale is a compact way to specify the acidity of a solution. pH is defined as the negative log of hydronium ion concentration:

• Therefore, a solution with (acidic) has pH of:

• Note that the significant digits in a logarithm appear after the decimal. Therefore, the 2 significant figures in the [ ] appear as 2 decimal places in the pH :
• In general,
  

  

• For a neutral solution:

• Listed on the right is the pH of some common substances. As discussed earlier, many foods are acidic and have low pH values. Few foods are basic.
• Note that:
• since pH is a negative scale, a lower pH value indicates greater concentration,
  and
• since pH is a logarithmic scale, a change of 1 pH corresponds to a 10 -fold change in concentration.

pH SCALE

• The relationship of pH scale and the are summarized below:

The pH Scale
	1	2	3	4	5	6	7	8	9	10	11	12	13	14
Acidic pH Basic
[ ]