ACIDS and BASES

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.

Relative Strengths of Acids and Bases

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 STRONG acids and bases Strong electrolytes Ionize (separate) 100% in water Strong Acids: HCl, HNO3, H2SO4, HClO4, HI, HBr Strong Bases: NaOH, KOH, Ca(OH)2 Most Group 1 and 2 metal hydroxides are strong bases

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 Pouvoir hydrogene: “Hydrogen Power” Uses [H3O+] or [H+] 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

pH Scale

pOH “Hydroxide Power” Uses [OH-] pOH = -log [OH-] 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+]eqm, 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 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:

Calculations pH = -log [H+] pOH = -log [OH-] pOH + PH = 14 [H+] x [OH-]= 1 x 10-14 So, given one value, you should be able to determine the other 3!

Practice pH pOH Calculate pH, pOH, and [OH-] [H+] = 1.00 x 10-3M [OH-] = 1.00 x 10-9M [OH-] = 5.72 x 10-5M Calculate pH, pOH, and [OH-] [H+] = 4.25 x 10-9M