Summary
Highlights
For oxyacids, more oxygen atoms result in a stronger acid (e.g., H2SO4 > H2SO3). A more electronegative central atom in an oxyacid also leads to a stronger acid (e.g., HClO3 > HBrO3 > HIO3). For binary acids (without oxygen), a weaker H-A bond (easier to break) results in a stronger acid. For instance, HF is a weak acid because fluorine's high electronegativity prevents it from releasing hydrogen easily.
This lecture will review the basics of acids and bases, including identifying them, understanding Bronsted-Lowry definitions, writing ionization equations, and using the ion product of water (Kw) for pH calculations. It also covers amphoteric behavior, where a substance like water can act as both an acid and a base. The content is primarily a review, so it will be covered quickly.
Acids are substances that donate hydrogen ions (protons, H+) to water, forming hydronium ions (H3O+). In aqueous solutions, H+ and H3O+ are considered equivalent. Acid characteristics include turning blue litmus paper red, having a sour taste, and being corrosive to skin. Six common acids are sulfuric, nitric, phosphoric, hydrochloric, acetic, and carbonic acids, which are frequently used in chemistry and everyday products.
Strong acids ionize completely (100% dissociation) in solution, indicated by a single-headed arrow in their ionization equation. Nitric acid is an example, dissociating into H+ and NO3-. There are six strong acids that are crucial to know. Weak acids do not hydrolyze completely, meaning they only partially dissociate and exist in equilibrium, represented by a double-headed arrow. Acetic acid is an example of a weak acid.
Bases either donate hydroxide ions (OH-) to a solution or accept protons/hydronium ions. Strong bases like sodium hydroxide dissociate to release hydroxide ions into the solution. Common bases include sodium hydroxide (used in soaps), magnesium hydroxide (in antacids), and ammonia. Bases taste bitter, feel slippery or slimy on the skin, and turn red litmus paper blue.
A neutral substance either doesn't ionize or has an equal number of H+ and OH- ions, like water. pH measures the concentration of hydronium (H+) ions in an aqueous solution. A high concentration of H+ ions results in a low pH, and vice versa. pH is essential for biological systems, such as maintaining blood pH for enzyme function.
A strong acid can exist at a low concentration, and vice versa. Strength refers to the degree of ionization, while concentration refers to the amount of substance. The pH scale uses logarithms: pH = -log[H+] and pOH = -log[OH-]. The relationship between pH and pOH is pH + pOH = 14. Each pH unit represents a tenfold change in H+ concentration. The Kw for water is 1 x 10^-14, meaning [H+] and [OH-] are both 1 x 10^-7 M.
Cations are stronger acids than neutral molecules, which are stronger than anions (e.g., H3O+ > H2O > OH-). Conversely, anions are stronger bases than neutral molecules, which are stronger than cations. A stronger acid will have a weaker conjugate base, and vice versa.
Acid strength is directly related to pH. Strong acids have 100% dissociation, while weak acids only partially dissociate. The pH scale generally ranges from 0 to 14, but highly concentrated strong acids or bases can have negative pH values or values above 14. Group 1 and 2 hydroxides are strong bases. All strong acids are monoprotic except H2SO4, which is diprotic, increasing H+ concentration significantly. The Ka (acid dissociation constant) for strong acids is essentially infinite. For weak acids, Ka relates the concentrations of products and reactants at equilibrium (Ka = [A-][H3O+]/[HA]), excluding water. A larger Ka indicates a stronger acid. The negative log of Ka is pKa, which will be used along with ICE tables for calculations.