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
• Common indicators are listed in Table M
Examples:
1. What color will thymol blue be in a solution with a pH of 6.5?
2. What color will methyl orange turn if a solution has a pH of 7.0?
3. What color will phenolphthalein turn in an acid?
Indicators
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:
1. [H+] = 1.0 x 10-7, pH = 2. 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: 1. Find the pOH of a solution with a pH = 8. Is
this solution acidic or basic?
2. 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 – 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:
1. HCl + NaOH H2O + NaCl
2. H2SO4 + Ca(OH)2 2H2O + CaSO4
3. 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
1. What volume of 1.0M sulfuric acid can be neutralized by 50.0mL of 3.0M sodium hydroxide?
2. 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?