The strength of an acid refers to how readily an acid will lose or donate a proton while the strength of a base is determined by its ability to deprotonate from other compounds.
Acid Strength and Strong Acids
Acids have long been recognized as a distinctive class of compounds whose aqueous solutions exhibit the following properties:
- A characteristic sour taste.
- Changes the color of litmus from blue to red.
- Reacts with certain metals to produce gaseous H2.
- Reacts with bases to form a salt and water.
Acidic solutions have a pH less than 7, with lower pH values corresponding to increasing acidity. Common examples of acids include acetic acid (in vinegar), sulfuric acid (used in car batteries), and tartaric acid (used in baking). Acids react with carbonates to create a salt, water and carbon dioxide.
There are three common definitions for acids:
- Arrhenius acid: any substances that increases the concentration of hydronium ions (H3O+) in solution.
- Brønsted-Lowry acid: any substance that can act as a proton donor.
- Lewis acid: any substance that can accept a pair of electrons.
The strength of an acid refers to how readily an acid will lose or donate a proton, oftentimes in solution. A stronger acid more readily ionizes, or dissociates, in a solution than a weaker acid. The six common strong acids are:
- hydrochloric acid (HCl)
- hydrobromic acid (HBr)
- hydroiodic acid (HI)
- sulfuric acid (H2SO4; only the first proton is considered strongly acidic)
- nitric acid (HNO3)
- perchloric acid (HClO4)
Each of these acids ionize essentially 100% in solution. By definition, a strong acid is one that completely dissociates in water; in other words, one mole of the generic strong acid, HA, will yield one mole of H+, one mole of the conjugate base, A−, with none of the unprotonated acid HA remaining in solution. By contrast, however, a weak acid, being less willing to donate its proton, will only partially dissociate in solution. At equilibrium, both the acid and the conjugate base will be present, along with a significant amount of the undissociated species, HA.
Two key factors contribute to overall strength of an acid:
- polarity of the molecule
- strength of the H-A bond
These two factors are actually related. The more polar the molecule, the more the electron density within the molecule will be drawn away from the proton. The greater the partial positive charge on the proton, the weaker the H-A bond will be, and the more readily the proton will dissociate in solution.
Acid strengths are also often discussed in terms of the stability of the conjugate base. Stronger acids have a larger Ka and a more negative pKa than weaker acids.
Base Strength and Strong Bases
There are three common definitions of bases:
- Arrhenius base: any compound that donates an hydroxide ion (OH–) in solution.
- Brønsted-Lowry base: any compound capable of accepting a proton.
- Lewis base: any compound capable of donating an electron pair.
In water, basic solutions will have a pH between 7-14.
A strong base is the converse of a strong acid; whereas an acid is considered strong if it can readily donate protons, a base is considered strong if it can readily deprotonate (i.e, remove an H+ ion) from other compounds. As with acids, we often talk of basic aqueous solutions in water, and the species being deprotonated is often water itself. The general reaction looks like:
A−(aq)+H2O(aq)→AH(aq)+OH−(aq)
Thus, deprotonated water yields hydroxide ions, which is no surprise. The concentration of hydroxide ions increases as pH increases.
Most alkali metal and some alkaline earth metal hydroxides are strong bases in solution. These include:
- sodium hydroxide (NaOH)
- potassium hydroxide (KOH)
- lithium hydroxide (LiOH)
- rubidium hydroxide (RbOH)
- cesium hydroxide (CsOH)
- calcium hydroxide (Ca(OH)2)
- barium hydroxide (Ba(OH)2)
- strontium hydroxide (Sr(OH)2)
The alkali metal hydroxides dissociate completely in solution. The alkaline earth metal hydroxides are less soluble but are still considered to be strong bases.
In the equation for acid HA is a protonated acid, H+ is the free acidic proton, and A– is the conjugate base.
Because strong acids and strong bases completely ionize in solution, the concentration of the strong acid will be the same as the concentration of the free acidic proton, and the -log[H+], or pH, can be simply calculated. The same is true of strong bases, and the pOH (-log[OH-]).
The equivalent mass (or equivalent weight) of an acid is the mass that yields one mole of hydrogen ions, or for a base it is the mass of the base that reacts with one mole of hydrogen ions.
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Key Points
• An acid is a substance that donates protons (in the Brønsted-Lowry definition) or accepts a pair of valence electrons to form a bond (in the Lewis definition).
• A base is a substance that can accept protons or donate a pair of valence electrons to form a bond.
• Bases can be thought of as the chemical opposite of acids. A reaction between an acid and base is called a neutralization reaction.
• The strength of an acid refers to its ability or tendency to lose a proton; a strong acid is one that completely dissociates in water. The molar concentration of acid then equals the molar concentration of protons, which can be used to calculate pH.
Key Terms
Valence electron: Any of the electrons in the outermost shell of an atom; capable of forming bonds with other atoms.
Ka: The acid dissociation constant is the equilibrium constant of the dissociation reaction of an acid.
pKa: A method used to indicate the strength of an acid. pKa is the negative log of the acid dissociation constant or Ka value. A lower pKa value indicates a stronger acid. That is, the lower value indicates the acid more fully dissociates in water.
equivalent mass: mass of acid that yields one mole of hydrogen ions, or mass of the base that reacts with one mole of hydrogen ions.