Acids and Bases

Electrolytes Substance that conducts electricity when dissolved in water Ionizes in water Examples: aqueous ionic solutions, acids, bases

Properties of Acids Acids have a pH less than 7.0 Acids will burn your skin Dilute acids have a sour taste Strong acids are good conductors of electricity Acids react with bases to form water and a salt (neutralization reaction) Acids react with certain metals to produce hydrogen gas Metals above H2 will react with acids to produce H2(g)

Properties of Bases Bases have a pH greater than 7.0 Bases have a slippery or soapy feeling Dilute bases have a bitter taste Strong bases are good conductors of electricity Bases react with acids to form water and a salt (neutralization reaction)

Common Acids and Bases

pH Scale Scale ranges from 0-14 Acids = pH less than 7 Bases = pH greater than 7 Neutral = pH = 7

pH of common substances

Indicators Common indicators are listed in Table M Examples: What color will thymol blue be in a solution with a pH of 6.5? What color will methyl orange turn if a solution has a pH of 7.0? What color will phenolphthalein turn in an acid?

Arrhenius Theory Definition of acids and bases Acids – substances whose water solution contain hydrogen ions (H+), ionize to produce H+ Bases – substances whose water solutions contain hydroxide ions (OH-), ionize to produce OH- Properties of acids and bases are due to an excess of H+ or OH- ions

Hydrogen/Hydronium H+ cannot exist unbonded in a system; instead it is bonded with a water molecule to make the hydronium ion (H3O+) H+ and (H3O+) are both used to indicate the presence of an acid

Bronsted-Lowry Theory Acids – donate/lose H+ Bases – gain H+ Example: NH3 + H2O  NH4+ + OH- H2O is the NH3 is the

Strength of Acids/Bases Strength is proportional to the degree to which it ionizes in solution Greater dissociation (more ions), stronger

Ionization Constant for Water Kw For pure water, at 25oC Kw = 1.0 x 10-14 = [H+][OH-] Kw is a constant therefore [H+][OH-] = 1.0 x 10-14 Examples: 1. Find the [OH-] concentration if [H+] = 10-8 2. Find the [H+] concentration if [OH-] = 10-3 3. Find the [OH-] concentration if [H+] = 10-7 * Exponents add up to -14

Hydrogen Ion Concentration (pH) used for convenience pH = -log [H+] Examples: [H+] = 1.0 x 10-7, pH = 2. [H+] = 1.0 x 10-3.5, pH = Acids: pH is lower than 7, [H+] is greater than [OH-] Bases: pH is greater than 7, [H+] is less than [OH-]

Hydroxide Ion Concentration (pOH) pOH = -log [OH-] Since [H+][OH-] = 1.0 x 10-14 pH + pOH = 14 Examples: Find the pOH of a solution with a pH = 8. Is this solution acidic or basic? Find the pH of a solution with a pOH = 12. Is this solution acidic or basic?

Concentration and pH A change of 1 in pH is a tenfold increase in acid or base strength A pH of 4 is 10 times more acidic than a pH of 5 A pH of 12 is 100 times more basic than a pH of 10

Mono/Di/Triprotic Monoprotic acids Diprotic aicds Triprotic acids produce a single hydrogen ion Examples: HCl, HBr Diprotic aicds produce two hydrogen ions Examples: H2SO4, H2S Triprotic acids produce three hydrogen ions Examples: H3PO4

Neutralization Reactions An acid and a base react together to form water and an ionic salt Examples: HCl + NaOH  H2O + NaCl H2SO4 + Ca(OH)2  2H2O + CaSO4 HNO3 + Ca(OH)2 

Acid-Base Titrations Process of adding a measured volume of an acid (or a base) of known molarity to a base (or an acid) of unknown molarity until neutralization occurs Standard Solution – acid or base of known molarity End Point – point of neutralization Unknown molarity is calculated using the titration formula

Titration MAVA = MBVB equation MA = Molarity of H+ MB = Molarity of OH- Moles of acid = moles of base (moles = molarity x liters)

Titration Examples What volume of 1.0M sulfuric acid can be neutralized by 50.0mL of 3.0M sodium hydroxide? 50.0mL of a 0.250M KOH are needed to neutralize 20.0mL of a HCl solution of unknown concentration. What is the concentration of the HCl?