Acids and Bases

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Acids and Bases

• Chapter 19

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Review

• Electrolyte
• A substance that conducts an electrical current
  when melted or in solution
• Ionic compounds

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Acid-Base Theories

• Different definitions of acids and bases
• Arrhenius
• Bronsted-Lowry
• Lewis

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Arrhenius

• Acid
• Compounds that ionize to produce hydrogen ions
  (H) in aqueous solutions
• Examples HCl, HBr, H2SO4, CH3COOH
• Note CH3COOH is an organic acid

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Arrhenius

• Base
• Compounds that ionize to produce hydroxide ions
  (OH-) in aqueous solutions
• Examples KOH, NaOH, LiOH
• Note CH3OH is not a base, its an organic
  alcohol

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Bronsted-Lowry

• Acid
• Hydrogen ion donor
• Examples HCl, HBr, H3O
• Base
• Hydrogen ion acceptor
• Examples H2O, NH3

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Lewis

• Acid
• Accepts a pair of electrons
• Example H
• Base
• Donates a pair of electrons
• Example OH-

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Acid-Base Theory
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Properties of Acids

• Taste Sour
• Will change color of acid base indicator
• Can be strong or weak electrolytes in an aqueous
  solution

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Properties of Bases

• Taste Bitter
• Feel Slippery
• Will change color of acid base indicator
• Can be strong or weak electrolytes in an aqueous
  solution

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Ionization

• Electrolytes will dissociate into ions when
  dissolved in water
• Strong Electrolytes will completely dissociate
• Weak Electrolytes will only partially dissociate

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Ionization of Water

• Water can be split into 2 ions
• H and OH-
• Ionization of Water
• H2O ? H OH-
• H2O H2O ? H3O OH-

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Strong Acids

• Completely dissociate when in solution
• HCl(s) ? H(aq) Cl-(aq)
• HNO3(s) ? H(aq) NO3-(aq)

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Polyprotic Acids

• Acids that have more than one H
• ExamplesH2SO4, H3PO4
• Can release more than one H into solution
• H2SO4(s) ? 2H(aq) SO42-(aq)

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Strong Bases

• Completely dissociate when in solution
• NaOH(s) ? Na(aq) OH-(aq)
• KOH(s) ? K(aq) OH-(aq)

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Neutral Solutions

• For neutral solutions
• H OH-
• For all aqueous solutions
• H OH- 1.0 x 10-14

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Measuring Acidity (Alkalinity)

• Traditionally we measure H
• pH -log H
• Neutral solution H 1.0 x 10-7
• pH 7
• pOH -log OH-
• pH pOH 14

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Acidity

• Acidic Solutions pH lt 7.0
• H gt 1.0 x 10-7
• Basic Solutions pH gt 7.0
• H lt 1.0 x 10-7

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Changes in pH

• pH increases by 1 for every decrease in H by a
  magnitude of 10

H pH
1.010-7 7
1.010-8 8
1.010-9 9
1.010-10 10
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Measuring pH

• Litmus paper
• Red in acid
• Blue in base
• pH paper
• pH Meter
• Acid Base Indicators (Table M)

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Table M
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Neutralization

• Acid Base ? Water Salt
• Double Replacement Reaction
• HA BOH ? HOH BA

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Neutralization

• Examples
• HCl NaOH ? H2O NaCl
• HNO3 LiOH ? H2O LiNO3
• H2SO4 2KOH ? 2 H2O K2SO4

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Titration

• Process in which a volume of solution known
  concentration is used to determine the
  concentration of another solution
• Usually shown by a color change of an indicator
  (end point)

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Titration Example

• 2 Liters of an unknown conc. of NaOH is titrated
  with 1 Liter of 6M HCl, what is the concentration
  of the base?
• MAVA MBVB
• (6M) (1L) (X) (2L)
• X 3M NaOH

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Polyprotic Acids

• Acids that have more than one H
• ExamplesH2SO4, H3PO4
• H2SO4 1M
• H 2M
• H3PO4 1M
• H 3M

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Another Titration Example

• 500 milliliters of an unknown conc. of NaOH is
  titrated with 1 Liter of 1M H2SO4, what is the
  concentration of the base?
• MAVA MBVB
• (2M) (1L) (X) (0.5L)
• X 4M NaOH

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Hydrolysis

• In salt hydrolysis, cations or anions of
  dissociated salts remove hydrogen ions from or
  donate hydrogen ions to water

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Hydrolysis

• Anions of weak acids remove hydrogen ions leaving
  solution basic

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Hydrolysis

• Cations of weak bases donate hydrogen ions
  leaving solutions acidic

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Buffer

• Solution in which the pH remains relatively
  constant when small amounts of acid or base are
  added
• Solution of a weak acid and one of its salts, or
  a solution of a weak base and one of its salts
• H2CO3/HCO3-