Chapter 8: Acids & Bases

Chapter 8: Acids & Bases • Acids are all around us. • Foods like lemons, limes, and oranges contain acidic compounds. • When wine is oxidized, it becomes acetic acid. • Our muscles produce lactic acid. • Our stomach uses hydrochloric acid to break down food.

Acids • Arrhenius definition – an acid is a substance that produces H+ ion when added to water. HCl(aq) H+(aq) + Cl-(aq) • Acids taste sour, are electrolytes, and neutralize bases.

Bases • Arrhenius definition – a base is a substance that produces OH- when added to water. NaOH(aq) Na+(aq) + OH-(aq) • Bases taste bitter, have a slippery, soapy feel, and neutralize acids.

Bronsted-Lowery Theory • Expanded definition of acids & bases. • Acid = a proton (H+) donor • Base = a proton acceptor • Important note about the use of H+ in equations. • For weak acids and bases, the reactions are reversible, equilibria.

Bronsted-Lowery Theory • HC2H3O2 + H2O  H3O+ + C2H3O2- • NH3 + H2O  NH4+ + OH-

Conjugates • Write the conjugate base of an acid. • Remove an H+ 1. HBr – H+ = Br− 2. H2S – H+ = HS− • Write the conjugate acid of a base. • Add an H+ 1. NO2− + H+ = HNO2 2. NH3 + H+ =NH4+

Learning Check 1. The conjugate base of HCO3− is a. CO32− b. HCO3− c. H2CO3 2. The conjugate acid of HCO3− is a. CO32− b. HCO3−c. H2CO3 3. The conjugate base of H2O is a. OH−b. H2O c. H3O + 4. The conjugate acid of H2O is a. OH−b. H2O c. H3O+

Strong Acids • Strong acids completely ionize (dissociate) in water. • Strong acids are also strong electrolytes. • There are six: HClO4, H2SO4, HI, HBr, HCl, and HNO3.

Weak Acids • Weak acids only partially dissociate to produce ions in solution. • Weak acids are weak electrolytes. • Too many to list, but some common ones are: H3PO4, HC2H3O2, HF, and HC6H7O6.

Comparison

Strong Bases • Strong bases completely ionize in water and are also strong electrolytes. • Only group 1A and 2A hydroxides are strong bases. • All other metal hydroxides are insoluble in water.

Weak Bases • Weak bases only partially react with water (accepting a proton) and are, thus, weak electrolytes. • Ammonia, NH3, is the most common weak base. • Lone pair on N group will accept a proton. • Other organic weak bases.

Ionization of Water • Water can act as both an acid and a base. • Any substance that does this is called amphoteric. • Pure water – two water molecules will occasionally react with each other where one is an acid and one is a base.

Ionization of Water • H2O + H2O  H3O+ + OH- • [ X ] = symbol for molarity. • In pure water, [H3O+] = [OH-]. • Kw = [H3O+] x [OH-]; where Kw = 1 E-14. • Thus, [H3O+] = [OH-] = ____________

Acidic, Basic, and Neutral Solutions

Calculating [H3O+] and [OH-] • If we know [H3O+], then [OH-] = • If we know [OH-], then [H3O+] = • Ex) [H3O+] = 2.5 E-5 • Ex) [OH-] = 4.8 E-3

pH Scale • One convenient method for measuring the acidity or basicity of a solution is to use the pH scale. • pH is a logarithmic (log) scale and equal to: pH = -log[H3O+]. • pOH = -log[OH-]. • pH + pOH = 14.

pH Scale • A word about significant figures. • 2.4 x 10-3M pH = 2.62 • Red numbers are the significant digits. • Blue numbers are exact numbers.

Basic Calculators Enter concentration Press “log” key Change the sign Record answer to proper s.f.’s TI-83 or TI-89 Press negative sign Press “log” key (select in catalog) Enter concentration, close parenthesis Enter key and round to proper s.f.’s Guide to Using Your Calculator

pH to Concentration • To convert a pH back to a concentration, you will use the “antilog” key = 10x. • [H3O+] = 10-pH • [OH-] = 10-pOH • Can also use the universal power (^) key.

Comparison of Values

Measuring pH • Can be done with a meter, pH paper, or an indicator.

Fill in the Chart

Reactions of Acids • An acid will react with most metals. • Mg(s) + 2 HCl(aq) MgCl2(aq) + H2(g)

Reactions of Acids • Acids react with any carbonate (CO3-2) and bicarbonate (HCO3-) to generate CO2.

Environmental Note: Acid Rain • Normally, rain is slightly acidic – pH of 5.5 to 6.2 – due to dissolved CO2. • Burning fossil fuels, which contain small amounts of Sulfur, produces SO3. • This is converted to H2SO4. • H2SO4 + CaCO3 CaSO4 + H2O + CO2

Neutralization • Acids neutralize bases and vice versa. • Acid + Base  Water + Salt • HCl + NaOH  H2O + NaCl • Balancing more complex acid-base neutralization. • Each H+ needs one OH-. • Each H+ and OH- makes one H2O.

Solution Stoichiometry • Acid-base titration – can use a known acid or base solution to determine an unknown counterpart. • Endpoint – when all of the unknown acid or base has reacted. • Indicator – a substance that changes color when it changes pH.

Solution Stoichiometry • Requires precise glassware to deliver the known solution = buret. • Stopcock allows for delivery drop-by-drop. • Volumes can be read to nearest 0.05mL.

Buffers • When a small amount of acid or base is added to pure water, the pH swings drastically. • Some solutions, though, will resist these wild swings in pH and are called buffer solutions. • Made from a ___________ and a salt containing the ______________.

Buffers • 1.0L of pure water + 0.0100moles (0.365g) of HCl. • [H3O+] = 0.010 mol / 1.0L = 0.010M • pH = • 1.0L of pure water + 0.0100moles (0.400g) of NaOH. • [OH-] = 0.010 mol / 1.0L = 0.010M • pH =

Buffers • Buffer of HF and NaF (note: Na+ is a spectator ion). • Ideal buffer would contain a 50 / 50 mixture of each. • 1.0L of 0.10 moles HF (2.0g) and 0.10 moles of NaF (4.2g) will have a pH of 3.17.

Buffers • HF + H2O  H3O+ + F- 50% 50% (from NaF) • Addition of strong acid • reacts with F- ion to generate more HF • Addition of strong base • reacts with HF to produce water plus more F-

Buffers • Buffer plus 0.010 moles of HCl. • pH = 3.08 (from 3.17). • Buffer plus 0.010 moles of NaOH. • pH = 3.26 (from 3.17).

Buffers in Blood • Normal blood pH is 7.35 to 7.45. • Outside this range, cells cannot function properly. • Two buffer systems are present to maintain this pH. • H2CO3 / HCO3- • H2PO4- / HPO4-2

Chapter 8: Acids & Bases