Chemistry: Acids and Bases Flashcards
Litmus Paper
Turns red in acidic solution and blue in basic solution.
Acids Properties
- Have a sour taste. - Aqueous solutions can conduct electricity. - React with bases to form water and a “salt.” - Nonoxidizing acids react with metals to produce hydrogen gas. - Cause color changes in plant dyes, turn litmus paper red.
Bases Properties
- Have a bitter taste. - Aqueous solutions can conduct electricity. - React with acids to form water and a “salt.” - Feel slippery to the touch. - Cause color changes in plant dyes, turn litmus paper blue.
Arrhenius Definition
Swante Arrhenius formed this toward end of 19th century. He defined an acid as a species that produces H+ (a proton) in an aqueous solution and a base as a species that produces OH- (a hydroxide ion) in an aqueous solution. These definitions, though useful, fail to describe acidic and base behavior in nonaqueous media. An example of an Arrhenius acid, base, and acid-base reaction, respectively, are: HCl(aq) –> H+(aq) + Cl-(aq) NaOH(aq) –> Na+(aq) + OH-(aq) HCl(aq) + NaOH(aq) –> NaCl(aq) + H2O(l)
Bronsted-Lowry Definition
More general definition proposed by Johannes Bronsted and Thomas Lowry in 1923. An acid is a species that donates protons, while a base is a species that accepts protons. Advantage to this definition is it’s not limited to aqueous solutions. Acids and bases always occur in pairs, called conjugate acid-base pairs. The two members of a conjugate pair are related by the transfer of a proton. For example, H3O+ is the conjugate acid of the base H2O and NO2- is the conjugate base of HNO2: H3O+(aq) H2O(aq) + H+(aq) HNO2(aq) NO2-(aq) + H+(aq)
Lewis Definition
At same time as Bronsted and Lowry, Gilbert Lewis also proposed a definition. He defined an acid as an electron-pair acceptor and a base as an electron-pair donor. Lewis’s are the most inclusive definitions. Just as every Arrhenius acid is a Bronted-Lowry acid, every Bronsted-Lowry acid is also a Lewis acid. However, the Lewis definition encompasses some species not included within the Bronsted-Lowry definition. For example, BCl3 and AlCl3 can each accept an electron pair and are therefore Lewis acids, despite their inability to donate protons.
Nomenclature of Arrhenius Acids
The name of an acid is related to the name of the parent anion (the anion that combines with H+ to form the acid). Acids formed from anions whose names end in -ide have the predfix hydro- and the ending -ic. F Flouride = HF Hydrofluoric acid Br Bromide = HB Hydrobromic acid Acids formed from oxyanions are called oxyacids. If the anion ends in -ite (less oxygen), then the acid will end with -ous acid. If the anion ends in -ate (more oxygen), then the acid will end with -ic acid. Prefixes in the names of anions are retained. ClO- Hypochlorite –> HClO Hypochlorous Acid ClO2- Chlorite –> HClO2 Chlorous Acid CLO3- Chlorate –> HClO3 Chloric Acid ClO4- Perchlorate –> HClO4 Perchloric Acid NO2- Nitrite –> HNO2 Nitrous Acid NO3- Nitrate –> HNO3 Nitric Acid
Hydrogen Ion Equilibria (pH and pOH)
Hydrogen ion concentration, [H+], is generally measured as pH, where: pH = -log[H+] = log(1/[H+]) Likewise, hydroxide ion concentration, [OH-], is measured as pOH where: pOH = -log[OH-] = log(1/[OH-]) In any aqueous solution, the H2O solvent dissociates slightly: H2O(l) H+(aq) + OH-(aq) This dissociation is an equilibrium reaction and is therefore described by a contsant, Kw, the water dissociation constant. Kw = [H+][OH-] = 10^-14 Rewriting this equation is log form gives: pH + pOH = 14 In pure H2O, [H+] is equal to [OH-] because for every mole of H2O that dissociates, one mole of H+ and one mole of OH- are formed. A solution with equal concentration of H+ and OH- is neutral and has a pH of 7 (-log 10^-7 = 7). A pH below 7 indicates a relative excess of H+ ions and, therefore, an acidic solution; a pH above 7 indicates a relative excess of OH- ions and, therefore, a basic solution.
P-Scale Values
A useful skill for various problems involving acids and bases, as well as their corresponding buffer solutions, is the ability to quickly convert pH, pOH, pKa, and pKb into nonlogarithmic form and vice versa. When the original value is a power of 10, the operation is relatively simple; changing the sign on the exponent gives the corresponding p-scale value directly. Example: If [H+] = 0.001, or 10^-3, then pH = 3. If Kb = 1,0 x 10^-7, then pKb = 7. More difficulty arises when the original value is not an exact power of 10; exact calculation would be excessively onerous, but a simple method or approximation exists. If the nonlogarithmic value is written in proper scientific notation, it will look like n x 10^-m, where n is a number between 1 and 10. The log of this product can be written as log(n x 10^-m) = -m + log n, and the negative log is thus m - log n. Now, since n is a number between 1 and 10, its logarithm will be a fraction between 0 and 1, thus, m - log n will be between m - 1 and m. Further, the larger n is, the larger the fraction log n will be, and therefore the closer to m -1 our answer will be.
Relative Strengths of Acids and Bases
The strength of an acid or base will depend largely upon its ability to ionize. The strength of an acid, for example, can be measured by the fraction of the molecules of that acid undergoing ionization. Subsequently, the acid strength can be expressed by the following equation. % Ionization = (ionized acid concentration at equilibrium/initial concentration of acid) x 100% When an acid or base is strong, its conjugate base and acid will be weak. The stronger the acid, the weaker the conjugate base. Furthermore, within a series of weak acids, the stronger the acid, the weaker its conjugate base for all acids and bases included.
Strong Acids and Bases
Strong acids and bases are those that completely dissociate into their component ions in aqueous solution. For example, when NaOH is added to water, it dissociates completely. Hence, in a 1-M solution of NaOH, complete dissociation gives 1 mole of OH- ions per liter of solution. pH = 14 - (-log[OH-]) = 14 + log[1] = 14 Virtually, no undissociated NaOH remains.
Note that the [OH-] contributed by the dissociation of H2O is considered to be negligible in this case.
Strong Acids & Bases: H2O Contribution
The contribution of OH- and H+ ions from the dissociation of H2O can be neglected only if the concentration of the acid or base is greater than 10^-7 M.
For example, the pH of a 1 x 10^-8 M HCl solution might appear to be 8 because -log(1x10^-8) = 8. However, a pH of 8 is in the basic pH range, and an HCl solution is not basic. This discrepancy arises from the fact that at low HCl concentrations, H+ from the dissociation of water does contribute significantly to the total [H+]. The [H+] from the dissociation of water is less than 1x10^-7 M due to the common ion effect.
The total concentration of H+ can be calculated from Kw = (x+1x10^-8)(x) = 1.0x10^-14, where x = [H+] = [OH-] (both from the dissociation of water molecules). Solving for x gives x = 9.5x10^-8 M, so [H+]total = (9.5x10-8 + 1x10-8) = 1.05x10-7 M, and pH -log(1.05x10-7) = 6.98, slightly less than 7, as should be expected for a very dilute, yet acidic solutions.
Strong Acids & Bases: Common Ones
Strong acids commonly encountered in the lab include HClO4 (perchloric acid), HNO3 (nitric acid), H2SO4 (sulfuric acid), and HCl (hydrochloric acid). Commonly encountered strong bases include NaOH (sodium hydroxide), KOH (potassium hydroxide), and other soluble hydroxides of Group IA and IIA metals.
Strong Acids & Bases: Calculation of pH and pOH
Calculation of the pH and pOH of strong acids and bases assumes complete dissociation of the acid or base in solution: [H+] = normality of strong acid and [OH-] = normality of strong base.
Weak Acids and Bases
Those that only partially dissociate in aq solutions. A weak monoprotic acid (HA) in aqueous solution will achieve the following equilibrium after dissociation (H3O+ is equivalent to H+ in aqueous solution): HA(aq) + H2O(l) H3O+(aq) + A-(aq)