Chapter 16: Acids and bases Chemistry 1062: Principles of Chemistry II Andy Aspaas, Instructor

Acids and bases Three concepts to describe behaviors of acids and bases Arrhenius concept: ionization of water molecules Bronsted-Lowry concept: donation and acceptance of protons Lewis concept: donation and acceptance of electron pairs

Arrhenius concept of acids and bases Svante Arrhenius, 1884 Arrhenius concept: Acid: increases concentration of hydronium ion, H3O+(aq), when dissolved in water Base: increases concentration of hydroxide ion, OH-(aq), when dissolved in water Hydronium and hydroxide ions are in equilibrium with water molecules, and addition of acids or bases alters this equilibrium

Arrhnenius concept Strong acid: completely ionizes in water to give a hydronium ion and an anion HClO4(aq) + H2O(l)  H3O+(aq) + ClO4-(aq) Other strong acids: H2SO4, HI, HBr, HCl, and HNO3 Strong base: completely ionizes in water to give a hydroxide ion and a cation NaOH(aq)  Na+ + OH- Group IA and IIA hydroxides are strong bases (except beryllium hydroxide) Singles out OH- as source of base character, even though other ions can give the effects of bases

Bronsted-Lowry concept 1923, Johannes Bronsted and Thomas Lowry Bronsted-Lowry concept involves proton-transfer reactions Acid: species which donates a proton Base: specis which accepts a proton H3O+(aq) + NH3(aq)  H2O(aq) + NH4+(aq)

Reversible acid-base reactions In reversible reactions, both forward and reverse reactions involve a proton transfer NH3(aq) + H2O(aq) ↔ NH4+(aq) + OH-(aq) Conjugate acid-base pair: compounds that differ only by the loss or gain of a proton NH3 (base) and NH4+ (acid) are a conjugate acid/base pair H2O (acid) and OH- (base) are the other pair

Amphiprotic species Amphiprotic species: species which can act as either an acid or a base, depending on the nature of the other reactants Water is amphiprotic: NH3(aq) + H2O(aq) ↔ NH4+(aq) + OH-(aq) Base Acid Acid Base HC2H3O2(aq) + H2O(aq) ↔ C2H3O2-(aq) + H3O+(aq) Acid Base Base Acid

Lewis concept Lewis concept: Lewis acid: species which can form a covalent bond by accepting an electron pair from another species Lewis base: species which can form a covalent bond by donating an electron pair to another species BF3 + :NH3  BF3—NH3

Relative strengths of acids and bases Strong acids ionize completely in water HCl(aq) + H2O(aq)  Cl-(aq) + H3O+(aq) Water acts as a base, accepting a proton from HCl The forward reaction is predominant, but in the reverse reaction, H3O+ would be the acid Of the two, HCl is the stronger acid, since it more readily donates its proton The arrow in a reaction containing a strong acid will point to the side which contains the weaker acid The same applies to the bases, water is a stronger base than Cl- since it more successfully attracts the proton The arrow also points towards the weaker base

Relative strengths of acids and bases But, in a weak acid like acetic acid, a small amount of its molecules are ionized HC2H3O2(aq) + H2O(aq) = C2H3O2-(aq) + H3O+(aq) The reverse direction predominates, and therefore H3O+ is a stronger acid than acetic acid, and acetate is a stronger base than water Stronger acids have weaker conjugate bases Weaker acids have stronger conjugate bases

Molecular structure and acid strength Electronegativity and bond strength determine the acidity of a proton If a proton is bonded to a more electronegative element, the bond is more polar and the proton is more easily lost Prevalent in comparisons across rows Long covalent bonds are weaker than short, so as the size of the atom H is bonded with increases, its acidity also increases Prevalent in comparisons down columns Oxoacids (H-O-Y-) increase in acidity with increasing electronegativity of Y, and with increasing number of oxygens attached to Y (both factors increase the partial negative charge of Y)

Self-ionization of water Pure water is ionized to a small extent A proton from one water molecule is transferred to another H2O(l) + H2O(l) = H3O+(aq) + OH-(aq) But since this only occurs to a small extent, the equilibrium constant for this process is small, and the concentration of water remains essentially unchanged Kw = [H3O+][OH-] = 1.0 x 10-14 at 25 °C Known as the ion-product constant The product of hydronium and hydroxide concentration for any aqueous solution is always Kw at 25 °C

Calculating ion concentrations In a strong acid solution, H3O+ will come completely from the acid itself The autoionization equilibrium is reversed due to Le Chatelier’s principle [H3O+] will equal the acid concentration In a solution of 0.1 M HCl, [H3O+] = 0.1 M Similarly for strong bases, [OH-] = base conc. Substitute into the Kw equation to find the other concentration

pH of a solution In acidic solutions, [H3O+] > 1.0 x 10-7 M In neutral solutions, [H3O+] = 1.0 x 10-7 M In basic solutions, [H3O+] > 1.0 x 10-7 M It is more convenient to give these values as pH pH = -log [H3O+] So, if [H3O+] = 1.0 x 10-3 M, pH = 3.00 Acidic solutions, pH < 7 Neutral solutions, pH = 7 Basic solutions, pH > 7

pH, pOH, and Kw pOH = -log [OH-] Kw = [H3O+][OH-] = 1.0 x 10-14 pH + pOH = 14 So, to find the pH of a basic solution, first find pOH, then subtract it from 14

Acid-base indicators Indicators change color to indicate the pH of a solution Phenolphthalein is colorless in acidic solutions, and pink in basic solutions Protonated phenolphthalein is a colorless acid. When deprotonated, it becomes a pink-colored base.