Acids and bases Chapter 19

i. Arrhenius definition Acids increase the concentration of hydrogen ions in aqueous solution Acids have more H+ ions than OH- ions Formulas for acid begin with H Bases increase concentration of hydroxide ions (OH-) in aqueous solutions Bases have more hydroxide ions than hydrogen ions Formulas for bases end with –OH-

ii. Properties Acids Bases Taste sour Taste bitter, feel slippery Reacts with metals to form H2(g) Litmus – blue Corrosive Phenolphthalein - pink Litmus – red Phenolphthalein – colorless

iii. Bronsted-lowry acids/bases Acids are hydrogen ion donors Bases accept hydrogen ions Conjugate acid – species formed after the base has taken a H+ Conjugate base – species formed after the acid has lost a H+ A base is a proton (hydrogen ion) acceptor, so the conjugate acid is the acid in the reverse reaction/process Example: Base = NO3-, Conjugate Acid = HNO3 Example = HCl, Conjugate Base = Cl- Polyprotic acids – can donate more than one H+ Example: H3PO4 (in reality, only one H is given up)

iv. Strong vs. Weak acids/bases Strong acids: completely ionize in water, strong electrolyte Strong Bases: same definition as strong acids Examples: HNO3, HCl, HI, HBr, H2SO4 Group 1 and 2 hydroxides (NaOH, KOH, Sr(OH)2, etc) Weak acids: do not ionize completely, some whole molecules remain, weak electrolytes Weak bases: same as weak acids NH3 Examples: HC2H3O2, HF

v. Neutralization reactions Acid + Base  Salt + water Salt – ionic compounds (except oxides) Salts can be Neutral – when both the acid and base are strong Acidic – when a strong acid reacts with a weak base Basic – when a strong base reacts with a weak acid

pH – measures the concentration of H+ ions in solution vi. Ph and poh pH – measures the concentration of H+ ions in solution Calculate: pH = -log[H+] pOH – measures the concentration of OH- ions in solution Calculate: pOH = -log[OH-]

c) Self-ionization of water vi. Ph and poh c) Self-ionization of water Amphoterism – water can act as either an acid or a base H2O + H2O  H3O+ + OH- H2O + NH3 H2O + HCl d) In one liter of pure water at 25oC i. the concentration of H+ = 1x10-7 ii. the concentration of OH- = 1x10-7

vi. pH and poh e) Formulas useful for pH and pOH problems (these are in your reference tables) i. pH = -log[H+] ii. pOH = -log[OH-] iii. pH + pOH = 14 iv. [H+] = 10-pH vi.[OH-] = 10-pOH vii. Kw = [H+][OH-] = 1x10-14

vi. Ph and poh f) Below 7 = acidic – more H+ ions exactly 7 = neutral – equal H+ and OH- Above 7 = basic – more OH- ions

standard solution unknown solution VII. Titration Concept: analytical method used to determine the concentration of an unknown substance. Often an acid-base reaction. Equivalence point: point at which moles of H+ = moles of OH-; can use this fact (and the volumes involved) to calculate the molarity of the unknown (similar to dilution formula) The equivalence point can be determined by measuring the pH at regular intervals and graphing the data or Estimate the equivalence point by using an indicator

c) End Point: the point at which an indicator changes color vii. titration c) End Point: the point at which an indicator changes color This will be close to (but not exactly) the equivalence point. For example – phenolphthalein changes from colorless to pink at a pH of about 8. If you are doing a strong acid/strong base titration, the equivalence point is at a pH of 7. The difference is only a mL or so, so it gives a very close approximation of the equivalence point

vii. titration

VII. titration d) Titration Calculation Formula MaVana = MbVbnb M = molarity V = volume n = number of moles of OH- (for the base) or H+ (for the acid) -use the subscript a refers to the acid, b refers to the base

vii. titration d. Titration Calculation i. example: 42.5mL of 1.3M KOH are required to neutralize 50.0mL of H2SO4. Find the molarity of the sulfuric acid