Properties of ACIDS • Sour to taste • React with some metals to form Hydrogen gas • Turn Litmus RED • Phenolphthalein stays colorless • Electrolytes (conduct) • Form H+ (H3O+ when attached to water molecules)
Properties of BASES • Bitter to taste • Slippery to touch • Turn Litmus BLUE • Phenolphthalein turns MAGENTA • Electrolytes • Many form OH- in water
Definitions of Acids and Bases • Arrhenius (traditional) • Acids: produce H3O+ in water • Bases: produce OH- in water • Examples • HCl(g) → H+(aq) + Cl-(aq) • NaOH (s) → OH-(aq) + Na+(aq) • Most common definition
Definitions of Acids and Bases • Bronsted-Lowry • Acids: H+ donor (proton donor) • Bases: H+ acceptor (proton acceptor) • Examples • HCl→ Cl- (donates H+) • NH3 → NH4+ (accepts H+)
Bronsted-Lowry cont. • When acids and bases donate or accept hydrogen ions, conjugate acids and bases are formed. • Conjugate Acid: particle formed when a Base gains a H+ • Conjugate Base: particle formed when an Acid donates a H+
Bronsted-Lowry cont. • Conjugate acid-base pair: two substances related by the loss or gain of a single H+ • Always paired with an acid and base • Examples- Label acid/base and conjugate acid/base • NH3 + H2O→ NH4+ + OH- • HCl+ H2O→ H3O+ + Cl- • Hint: Acid/Base is always Reactant; Conjugate Acid/Base is always Product
Terms • Amphoteric substances can behave as acids or bases (water) • Monoprotic acids donate only one H+ • Ex. HCl, HNO3 • Polyprotic acids donate more than one H+ • H2SO4, H2CO3
Relative Strengths of Acids and Bases • The stronger the acid, the weaker the conjugate base. • H+ is the strongest acid that can exist in equilibrium in aqueous solution. • OH- is the strongest base that can exist in equilibrium in aqueous solution.
Conjugate Acids and Bases • Any acid or base that is stronger than H+ or OH- simply reacts stoichiometrically to produce H+ and OH-. • The conjugate base of a strong acid (e.g. Cl-) has negligible acid-base properties. • Similarly, the conjugate acid of a strong base has negligible acid-base properties.
Acid-Base Equilibria • In pure water, the following equilibrium is established: • This is referred to as the autoionization of water.
Types of Acids and Bases • Weak acids and bases • Weak electrolytes • Partially ionize in water (much less than 100%) • Establish equilibria • Weak Acids • H3PO4, HC2H3O2, H2CO3 • Weak Bases • NH3, low [OH-]
pH • Pouvoirhydrogene: “Hydrogen Power” • Measure of the acidity of a solution • Uses [H3O+] or [H+] • Concentrations usually expressed as powers of 10 • pH = -log [H+] • pH scale 0-7 acid, 7 neutral, 7-14 base • Can be lower than 0 or higher than 14
pOH • “Hydroxide Power” • Measures alkalinity of a solution • Uses [OH-] • pOH = -log [OH-] • pOH scale 0-7 base, 7 neutral, 7-14 acid • Can be lower than 0 or higher than 14 • pH + pOH = 14 • From the autoionization of water
Calculating Ka from pH • Weak acids are simply equilibrium calculations. • The pH gives the equilibrium concentration of H+. • Using Ka, the concentration of H+ (and hence the pH) can be calculated. • Write the balanced chemical equation clearly showing the equilibrium. • Write the equilibrium expression. Find the value for Ka. • Write down the initial and equilibrium concentrations for everything except pure water. We usually assume that the change in concentration of H+ is x. • Substitute into the equilibrium constant expression and solve. Remember to turn x into pH if necessary
Percent ionization is another method to assess acid strength. • For the reaction
Percent ionization relates the equilibrium H+ concentration, [H+]eq, to the initial HA concentration, [HA]0. • The higher percent ionization, the stronger the acid. • Percent ionization of a weak acid decreases as the molarity of the solution increases.
Polyprotic acids • Polyprotic acids have more than one ionizable proton. • The protons are removed in steps not all at once: • It is always easier to remove the first proton in a polyprotic acid than the second. • Therefore, Ka1 > Ka2 > Ka3 etc.
Relationship between Ka and Kb • For a conjugate acid-base pair • Therefore, the larger the Ka, the smaller the Kb. That is, the stronger the acid, the weaker the conjugate base. • Taking negative logarithms:
Strong Acids and Bases • Strong electrolytes • Ionize (separate) 100% in water • Strong Acids: HCl, HNO3, H2SO4, HClO3, HClO4, HI, HBr • Usually only source of H+, so pH can be calculated from the molarity of the acid (unless < 10-6) • Strong Bases: NaOH, KOH, Ca(OH)2 • Most Group 1 and 2 metal hydroxides are strong bases • Ionic metal oxides, hydrides, and nitrides
Salt Solutions • Most salts are strong electrolytes (ionize in solution) • Acid-Base properties result from their ions. • Many ions react with water to form H+ and OH- (hydrolysis) • Anions from weak acids are basic. • Anions from strong acids are neutral. • All cations (except alkali/alkaline earth metals) are weak acids
Salt Solutions • SA + SB = Neutral Salt • SA + WB = Acidic Salt • WA + SB = Basic Salt • WA + WB = Acidic or Basic Salt • Based on relative strength of Ka and Kb • Ka > Kb = acidic • Ka < Kb = basic
Acid-Base Behavior and Chemical Structure • For compound H-X, • If H is partially positive, then it is an acid • If H is partially negative, then it is a base • Bond Strength and Polarity affects acid/base strength • Bond strength used to determine strength in a group; Bond polarity used to determine strength in a period • Acid strength tends to increase down a group; Base strength tends to decrease down a group • Acid strength tends to increase L to R; Base strength tends to decrease L to R
Oxyacids • Acids that contain OH groups bound to the central atom ( Y – O – H) • Strength depends on Y and the atoms attached to Y • Increasing electronegativity of Y = increasing acidity • Increasing the number of O atoms attached to Y increases polarity, which increases strength • Ex. HClO < HClO2 < HClO3 < HClO4
Carboxylic Acids • Contain –COOH • Additional oxygen atom on the carboxyl group increase the polarity of the O-H bond and stabilizes the conjugate base
Lewis Acids and Bases • Lewis Acid – electron pair acceptor • Lewis Base – electron pair donor • Do not need to contain protons – most general definition of acids/bases • Many Lewis acids have an incomplete octet (BF3) • Transition metal ions are usually Lewis acids • Lewis acids must have an empty orbital • Compounds with multiple bonds can be Lewis acids
Amino Acids • Amphoteric • Contains carboxyl group (acid) and ammine group (basic) • Proton of the carboxyl group is transferred to the basic nitrogen of the ammine • Results in a zwitterion or dipolar ion