Chapter 16: Acids and bases

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

2. 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

3. 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

4. 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

5. 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)

6. 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

7. 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 BaseAcid

8. 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

9. 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

10. 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

11. 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)

12. 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 Kwat 25 °C

13. 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 Kwequation to find the other concentration

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

15. 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

16. 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.