Acid Base Equilibrium
Pure water undergoes self-ionization into hydronium and hydroxyl ions as shown below.
H2O + H2O ⇆ H3O+(aq) + OH–(aq)
Using the equation above:
Kc = ([H3O+][OH–])/[H2O]2
Because H2O is a liquid, it can be excluded from the formula, and end up with
Kc = [H3O+][OH–]
This formula can be simplified to
Kc = [H+][OH–]
In this case, Kc is defined as the autoionization constant of water, with the symbol Kw. At 25°C, Kw has the value of 1.0 x 10–14.
When;
[H+] > [OH–] the solution is acidic. When [H+] = [OH–] the solution is neutral, and if [H+] is less than [OH–], the solution is basic.
Neutral solutions, where [H+] = [OH–] mean that the Kw = √1.0 x 10–14 = 1.0 x 10–7.
Pure water has equal concentrations of H+ and OH–, therefore it is neutral. Note that, as temperature rises Kw increases.
Percent ionization is similar to the concept of percent yield;
pH is defined as the negative log of [H+] i.e. pH = -log[H+]
The pH scale condenses possible values of [H+] to a 14 point scale where pH =7 actually indicates [H+] = 1.0 x 10–7
Percent ionization = (ionized acid concentration at equilibrium/Initial concentration of acid ) * 100
Percent Ionization = ([H3O+]eq/[HA]initial)*100.
A strong acid is explained as an acid that reacts quantitatively with water to form hydronium ions. It has a percent ionization > 99%.
A weak acid is explained as an acid that reacts partially with water to form hydronium ions (most less than 50%).
The Arrhenius Theory defined acids as substances that ionize to produce hydrogen ions plus an anion, and bases as substances that dissociate in water to produce hydroxide ions plus a cation. These definitions apply in several situations but have failed to predict the acidic or basic properties of some compounds.
To address this weakness, Arrhenius modified the theory to define acids as substances that react with water to produce hydronium ions. A hydronium ion is represented as H3O+. A base was then defined as a substance that reacts with water to produce hydroxide ions. Limitations still persisted with these modified definitions for example, they did not explain reactions that do not occur in aqueous solutions.
Bronsted–Lowry revised this definition and excluded the concept of a substance being an acid or a base. Instead, a substance is referred to act as an acid, or act as a base in the context of a specific reaction. According to the Bronsted–Lowry concept, an acid is a proton donor, and a base is a proton acceptor in the specific reaction.
Consider two examples:
1. HCl(aq) + H2O(l) → H3O+(aq) + Cl–(aq)
In this example, HCl acts as the acid, donating a proton (H+) to make hydronium ions. Water is the base, accepting the proton.
2. NH3(aq) + H2O(l) → OH - (aq) + NH4 + (aq)
In this second example, NH3(aq) acts as the base, accepting a proton, while H2O(l) acts as the acid, donating a proton.
Notice that water has the ability to either accept or donate a proton, making it amphiprotic. The bicarbonate ion HCO3- is also amphiprotic.
The Bronsted–Lowry concepts results in an acid–base dynamic equilibrium, where the forward and reverse reactions occur at the same time and at the same rate. A pair of substances that differ only by a proton is called a conjugate acid–base pair. The example below shows two conjugate pairs:
CH3COOH(aq) + H2O CH3COO-(aq) + H3O+(aq)
In this example above, Acetic acid has a strong attraction for its own proton and doesn’t readily donate it, this makes it a weak acid. It’s conjugate base, the acetate ion, is a stronger base than water. It has a greater attraction for protons. The weaker an acid is, the stronger it’s conjugate base. The stronger an acid, the weaker it’s conjugate base.
In a system that contains several different possible acid-base reactions, the only significant reaction is a proton transfer from the strongest acid present to the strongest base present.
Steps For Predicting the Predominant Acid–Base Reaction
Acid and base strengths
Consider the dissociation of hydrofluoric acid:
HF (aq) + H2O (l) ⇆ H3O+(aq) + F- (aq)
The acid ionization constant (Ka) indicates the extent to which an acid will react with water. It is a ratio of the dissociated form of the acid to the undissociated form.
The H2O is ignored because it is liquid. Strong acids have high Ka values while weak acids have lower Ka values. Ka less than 0.001 are weak acids, Ka between 0.001 and 1 are moderate acids, and Ka greater than 1 are strong acids.
Dissociation refers to the formation of ions when a chemical comes apart. Should be differentiated from Ionization which refers to the formation of ions when two chemicals react. Even though we write HCl ⇆ H+ + Cl–, this is just the short form. In reality HCl reacts with H2O, thus it is an ionization not a dissociation. on the other hand, NaCl can also dissociate in water. This is not an ionization, since water is only required to stabilize ions, it is not needed as a reactant.
A base ionization constant (Kb) indicates the extent to which a base will react with water. For example:
C6H5O73-(aq) + H2O(l) ⇆ HC6H5O72-(aq) + OH-(aq)
Amphoteric substances can act as acids or as bases. If Ka > Kb, then it acts as an acid. If Ka is less than Kb, then it acts as a base.
A monoprotic acid has one ionizable H atom per molecule such as hydrochloric acid (HCl).
A polyprotic acid has more than one ionizable H atom per molecule such as sulfuric acid, H2SO4. The protons are removed in steps not all at once. It is always easier to remove the first proton in a polyprotic acid than the second. i.e. Ka1 > Ka2 > Ka3.
The pH curves of polyprotic and polybasic species can have more than 1 equivalence point or end point because the acid/base will react in successive steps until all protons have been transferred.
Terminology:
End point is a point of a titration at which a sharp change in a measurable and characteristic property occurs, for example, a color change in the indicator.
An equivalence point the measured quantity of titrant recorded when the endpoint occurs, the point when equal amounts of acid and base have reacted.
Buffering refers to the resistance to changes in pH by a solution, so only a slight change in pH occurs. Buffering occurs because initially, any acid added immediately reacts with the excess hydroxide and is converted to water.
The steep slope is where complete neutralization occurs
When a strong acid is titrated with a strong base.
When a strong base is titrated with a strong acid.
When a strong acid is titrated with a weak base.
When a strong base is titrated with a weak acid.
The titration curve of a polyprotic acid.