Acids and Bases

Acids and Bases




An understanding of acids and bases forms an essential part of a total grasp of chemistry. It encompasses many of the chemical concepts you learn early in a first semester. Organic chemistry and biochemistry cannot be understood without a mastery of acids and bases.







Acids and bases are composed of molecules which form ions in solution. Historically, three definitions were needed to explain acid-base behavior as chemists became aware of ever more broadening classes of compounds which underwent acid-base reactions,



Because of this, the trend makes each kind of acid and base more inclusive as better approximations revealed these kinds of reactions.




Arrhenius Acids and Bases




Arhhenius acid base
Figure 1: Arrhenius acid and bases classified on what they do to an aqueous solution




When an Arrhenius acid dissolves in water, it ionizes into an H+1 and negative anion.  For example, when HCl bubbles through water, it produces an excess of H+1 cations, Equation (1).



\LARGE{HCl \longrightarrow H^{+1} + Cl^{-1}}                                        (1)



HCl dissociates into a positive cation, H+1, and negative anion, Cl-1. Notice the overall charge of the solution remains neutral. Bear that in mind when you think about acid and base dissociation. The overall charge conserves its neutral charge.


Sometimes to emphasize an interaction with water, the same acidic behavior is depicted as in in Equation (2).




\LARGE{HCl + H_2O \longrightarrow H_3O^{+1} + Cl^{-1}}               (2)


This shows H+1 transfers from its bond with Cl to form a new bond between O and H. This results in a positive H3O+1 polyatomic ion and a Cl-1 anion. Students can become confused about two ways to represent the same thing. However for the moment, regard these two ways of representing an aqueous acid as identical.






In a similar vein, when a compound dissolves in water which produces an excess of hydroxide anions, HO-1, it is a base, (3). Chemical equation (3) shows a simple dissociation which shows NaOH as aqueous ions in water. Equation (4) includes water as a reactant in the same way as (2), in the case of an acid.



\LARGE{NaOH \longrightarrow Na^{+1} + Cl^{-1}}                        (3)


\LARGE{NaOH + H_2O \longrightarrow Na^{+1} + Cl^{-1}}+ H_20        (4)



In both representations of a base, excess hydroxide ions make a solution basic.



Bronsted-Lowry Acids and Bases



Compounds however react as acids and bases even when not in a solution of water. They can can show acid-base behavior without necessarily producing an excess of H+1 in the case of an acid, or excess OH-1 in the case of a base. In fact, although the definition of an acid or base may appear similar, this definition focuses on the behavior of a compound within the context of a reaction




A quick mention on vocabulary: the term “proton” is used to describe the H+1 ion. When you take a hydrogen atom and remove an electron, you have a proton. So the naked H+1 ion is considered synonymous with a “proton”. 



Bronsted Acid Base
Figure 2: Bronsted acid-base definition depends only on where and how a proton is transferred



In this way a compound is not an acid or base, but given the role of an acid or base in a specific reaction with a specific reaction partner. The same substance can act as an acid or base, depending on the conditions. To assign what role a compound takes, the acid is the species that donates an H+1. A base accepts an H+1



Nothing in this definition dictates what the external condition of a solution might be, or limits the definition to a compound’s behavior to what happens in aqueous solution. It also accounts for the fact that H+1 as a separate entity does not actually exist as such in a solution. It is passed from one molecule to another; from donor to acceptor. 







A familiar example of an acid-base reaction is the reaction of sulfuric acid in water, (5):



\LARGE{H_2SO_4 + H_2O \longrightarrow HSO_4^{-1} + H_3O^{+1}}     (5)



Sulfuric acid donates a proton while water accepts a proton. In this case, sulfuric acid is the acid and water is the base. In aqueous solution, you meet no surprises. Sulfuric acid acts just like an Arrhenius acid. However water does not act like an Arrhenius base. Water does not produce excess hydroxide ions.



The identity of a Bronsted-Lowry acid might not be as obvious when you consider the reaction of sodium bicarbonate with sodium hydroxide:




\LARGE{NaHCO_3^{-1} + 2  NaOH \longrightarrow Na_2CO_3 + H_2O}    (6)


Although NaOH is an obvious base, notice no excess hydroxide ions are produced as a result of the reaction. More over,  you may believe if a substance bears a negative charge it must be a base. However the only reliable way to assign which reactant is an acid and which is a base is to examine the reaction and identify which reactant donates donates a proton.







In a strict manner like Bronsted aids, bases follow the strict definition that a Bronsted-Lowry base accepts a proton. This makes it possible to include reactants like triethylamine, (CH3CH2)3N, which reacts with acids to form ammonium salts, (CH3CH2)3NH+1 :


\LARGE{HCl + (CH_3CH_2)_3N \longrightarrow (CH_3CH_2)_3NH^{+1} + Cl^{-1}}    (6)



The Bronsted-Lowry acid-base definition has many important applications to general chemistry and organic chemistry.

Lewis Acids and Bases



The final most common classification for acids and bases focuses on the movement of electrons rather than the movement of protons.



This particular definition becomes necessary to include the acid-base behavior of reactants like boron fluoride BF3, which is an acid or trimethylamine N(CH3)3  which is a base.


Electrons move from elements which are electron deficient to elements which have an electron excess, Figure 3.


This kind of movement of electrons should not be confused with an oxidation-reduction reaction. In this case, the charge of the element does not change.




Lewis acid base
Figure 3: Lewis acid base definition based on the direction of electron transfer






Any element which can accommodate more electrons can act as a Lewis acid. The previous example of BF3, like B(OH)3, or FeCl3, can act as a Lewis acid, Figure 4.






Lewis acid base
Figure 4: examples of Lewis acids




Lewis bases have elements with a surplus of electrons in the form of lone pairs of electrons, or multiple bonds which act as a source of labile electrons.



Lewis base structures
Figure 5: A few examples of Lewis bases