The student is expected to define acids and bases and distinguish between Arrhenius and Brønsted-Lowry definitions and predict products in acid base reactions, precipitation reactions, and oxidation-reduction reactions AND understand and differentiate among acid-base reactions, precipitation reactions, and oxidation-reduction reactions AND define pH and use the hydrogen or hydroxide ions concentrations to calculate the pH of a solution AND distinguish between degrees of dissociation for strong and weak acids and bases.
According to the Arrhenius definition of acids and bases in aqueous solutions, an acid is a substance that forms more hydrogen ions (H+) than hydroxide ions (OH-), while a base forms more hydroxide ions (OH-) than hydrogen ions (H+). According to the Brønsted-Lowry definition of acids and bases, an acid is a substance that can donate a proton (H+) and a base is a substance that can accept a proton.
During an acid-base reaction, the hydrogen ions (H+) from the acid will combine with the hydroxide ions (OH-) from the base. There must be a proton transfer for the reaction to occur. Strong acids will more likely transfer their protons to a strong base. The favored reaction direction will have a weak acid and base as products.
Acid base reactions involve the transfer of protons. In these neutralization reactions, the acid and the base react to form water and salt. A precipitation reaction occurs when two ionic compounds exchange cations in a solution and an insoluble precipitate is formed. Both precipitation and acid-base reactions are types of double replacement reactions. Oxidation-reduction (redox) reactions involve the transfer of electrons, which changes the oxidation numbers of the elements involved in the chemical change.
The pH scale is used to determine the strength of an acid and equal to the negative logarithm (base 10) of the hydrogen ion concentration [H+]. The concentrations of hydrogen and hydroxide ions multiply together equal the ion-product constant for water (1 x 10-14 ). In this way, the hydroxide ion concentration can also be used to calculate the pH of a solution.
When an acid or base dissolves in an aqueous solution, the ions may or may not dissociate completely. Complete dissociation of ions occurs with strong acids and strong bases, while partial dissociation of ions occurs with weak acids and weak bases. The conjugate base of a strong acid is a weak base. The conjugate acid of a strong a base is a weak acid.
Arrhenius Definition
When talking about acids and bases, there are several different definitions. The first scientist to recognize and describe acids and bases was the Swedish chemist Svante Arrhenius. In the late 1800s, Arrhenius defined acids and bases in aqueous solutions as follows:
Arrhenius Acid: An ionic compound that ionizes in aqueous solution and produces hydrogen ions (H+).
Arrhenius Base: An ionic compound that ionizes in aqueous solution and forms hydroxide ions (OH-).
According to Arrhenius’s definition, only substances that release hydrogen ions may be classified as acids. Likewise, only substances that release hydroxide ions may be classified as bases. However, later scientists found many substances that do not fit these definitions, yet still have properties of either acids or bases.
Brønsted-Lowry Definition
In 1923, a Danish chemist named Johannes Brønsted and an English chemist named Thomas Lowry independently proposed new definitions to describe acids and bases. In the Brønsted-Lowry definition, an acid is a compound that donates a proton (hydrogen ion H+) and a base is a compound that accepts a proton. The ion or molecule remaining after the acid has lost a proton is known as that acid’s conjugate base, and the species created when the base accepts the proton is known as the conjugate acid. This is expressed in the following reaction:
acid + base ↔ conjugate base + conjugate acid
According to the Brønsted-Lowry definition of acids and bases:
Brønsted-Lowry Acid: An acid is a substance that donates a proton (H+).
Brønsted-Lowry Base: A base is a substance that accepts a proton.
Brønsted-Lowry acids and bases behave in pairs such that there is always a proton donor and a proton acceptor. This is an important point as it highlights the differences between the Arrhenius and Brønsted-Lowry definitions of acids and bases. Water behaves as a Brønsted-Lowry base involving hydrochloric acid, and as a Brønsted-Lowry acid involving ammonia. However, water would not be classified as either an acid or a base according to the Arrhenius definition.
Types of Reactions
Both precipitation and acid-base reactions are double-replacement reactions.
Acid-Base Reaction: This is a type of double-replacement reaction that occurs when equal amounts of an acid are added to a base so that the acid and the base neutralize each other. An acid-base reaction is also known as a neutralization reaction in which water and salt are produced, thus neutralizing the pH of the solution. During an acid-base reaction, the hydrogen ions (H+) from the acid will combine with the hydroxide ions (OH-) from the base. There must be a proton transfer for a reaction to occur. Note, strong acids will more likely transfer their protons to a strong base. The favored reaction direction will result in a weak acid and a weak base as products.
Precipitation Reaction: This is a also a type of double-replacement reaction. These reactions occur in aqueous solutions of ionic compounds, where one of the reactants formed is a solid that precipitates out of solution. In a precipitation reaction, the rules on solubility are followed to determine which reactants produce a solid when mixed. The formations of solids are caused by a stronger interaction between solutes than between solute and solvent.
Oxidation-Reduction (Redox) Reactions: These types of reactions involve a change in the oxidation states of each species involved in a chemical reaction. There are several different types of reactions that may also be considered redox reactions. These types of reactions include the transfer of oxygen atoms and/or electrons. The most current definition of oxidation and reduction says that the term oxidation refers to the loss of electrons or to the gain of oxygen, where the term reduction refers to the gain of electrons or to the loss of an oxygen. A decrease in the oxidation state means an electron is gained (reduction), and an increase in the oxidation state means an electron is lost (oxidation).
The pH Scale Determines Strength of Acid or Base
The pH of a solution or substance is the number that is used to describe the hydrogen ion concentration of that solution or substance. The pH scale is a range that goes from 0 to 14. If a substance is 0 on the pH scale, it is highly acidic, meaning that it has a high hydrogen ion concentration. If a substance is 14 on the pH scale, it is highly basic, meaning that it has a very low hydrogen ion concentration, but a very high hydroxide ion concentration. If a substance is 7 on the pH scale, it is neutral, meaning the hydrogen ion and hydroxide ion concentrations are equal. The pH scale is used to determine the strength of an acid or base and is equal to the negative logarithm (base 10) of the hydrogen ion concentration [H+], as follows:
pH = -logarithm (hydrogen ion concentration)
Thus, we have:
pH = -log[H+] for acidic solutions
To avoid confusion when dealing with small numbers, dissociation constants and ion concentrations are not stated in exponential form, but as the negative logarithms of the actual values. The letter phas been chosen to mean “negative logarithm of.” Thus, pH means the negative log of the hydrogen ion (H+) concentration, and pOH means the negative log of the hydroxide ion (OH-)concentration.
Pure water, or any neutral solution, has a pH of 7, which may be represented as:
pH = -log(1 x 10-7) OR
[H+] = 1 x 10-7 M
A strong acid, such as HCl, with a pH of 0, may be represented as:
[H+] = 1 x 100 M
[OH-] = 1 x 10-14 M
A strong base, such as NaOH, with a pH of 14, may be represented as:
[H+] = 1 x 10-14 M
[OH-] = 1 x 100 M
The ranges on the pH scale are as follows: below 7 for acids, above 7 for bases, and 7 for neutral solutions. The product of the concentrations of hydrogen and hydroxide ions equals the ion-product constant for water (1 x 10-14):
Ionization constant of water = (hydrogen ion concentration)(hydroxide ion concentration)
Kw = [H+][OH-]
Note, the hydroxide ion concentration can also be used to calculate the pH of a solution by first calculating the hydrogen ion concentration.
Comparing Strong and Weak Acids and Bases
When an acid or base dissolves in an aqueous solution, the ions may or may not dissociate completely. In strong acids and bases, there is a complete dissociation (ionization) of ions in aqueous solution. In weak acids and bases, there is only a partial dissociation (ionization) of ions in aqueous solution.
As mentioned earlier, acids and bases have conjugates when they dissociate in aqueous solution. Strong acids completely ionize in solution, producing large concentrations of hydrogen ions, with the pH values approaching 0. The conjugate base of a strong acid is a weak base. Strong bases completely ionize in solution, producing large concentrations of hydroxide ions, with the pH values approaching 14. The conjugate acid of a strong base is a weak acid.
On the other hand, a weak acid or a weak base will not completely ionize, producing fewer hydrogen and hydroxide ions than a strong acid or base. As expected, their pH values are closer to neutral. The closer an acid’s or base’s pH is to 7, the weaker the acid or base.