Acids and Bases

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Chapter 10

• Acids and Bases

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Sec 10.1 Arrhenius Acid-Base

• There are two main definition sets for acids and
  bases
• Arrhenius definitions (this section)
• Bronsted-Lowry definition (sec 10.2)
• Arrhenius acid produces H3O ions in water
• Arrhenius base produces OH- ions in water

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Sec 10.1 Arrhenius Acid-Base

• H goes to H3O , we think of them as being
  interchangeable
• H is really a proton (why?)
• Acids ionize, that is they dissolve in solution
  to give positive and negative ions

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Sec 10.1 Arrhenius Acid-Base

• Bases produce OH- ions in water either by
• 1) Being a metal hydroxide that dissociates
• NaOH, Mg(OH)2, etc
• 2) Reacting with water molecules in such a way to
  create OH-
• NH3 H2O NH4 OH-
• NH3 is a base in this reaction...what is H2O ?

5
Sec 10.2 Bronsted-Lowry

• Bronsted-Lowry is a more complete definition of
  acids and bases
• B-L Acid proton donor
• B-L Base proton acceptor
• Acid base reaction Transfer of a proton from
  an acid to a base
• Note In an acid-base reaction there is ALWAYS an
  acid and a base

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Sec 10.2 Bronsted-Lowry

• Bronsted-Lowry came up with the idea of a
  conjugate acid-base pair. In an acid-base
  reaction there are actually two conjugate pairs.
  Lets identify them
• Examples
• HNO3 H2O NO3- H3O
• CH3COOH NH3 CH3COO- NH4

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Sec 10.2 Bronsted-Lowry

• General Points and Rules
• 1. Acids can be positively, neutral or negatively
  charged
• 2. Bases can be negatively or neutral charged but
  not positively
• The acid in the pair always has one more H atom
  and one fewer negative (-) charge

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Sec 10.3 Mono, Dri, Triprotic

• Acids can be monoprotic, diprotic, or triprotic
• These terms refer to the acidic hydrogens.
• Not every hydrogen for every molecule is acidic
• The acidic ones are normally written at the start
  of the molecule
• Examples H2SO4, H2CO3, H3PO4, HCH3OO

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Sec 10.3 Mono, Dri, Triprotic

• Acids can be monoprotic, diprotic, or triprotic
• Examples H2SO4, H2CO3, H3PO4
• Some molecules can function both as an acid as a
  base, (water, HCO3-)
• This is called amphiprotic or amphoteric

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Sec 10.4 Strengths of Acids

• The relationship between an acid and its
  conjugate base is an inverse one
• Ie, if the acid is strong, then the conjugate
  base will be weak
• The weaker the acid, the stronger its conjugate
  base will be
• Think of this in terms of desire for the proton.
  If the acid wants to donate the proton in the
  first place, how likely is it to come back?

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Sec 10.4 Strengths of Acids
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Sec 10.4 Strengths of Acids

• Both acids and bases have scales of strength
• Strength is really a measurement of the degree of
  dissociation
• The more an acid dissociates, the more H3O
  particles are formed, and vice versa for a base
• We think of strong acids and strong bases as
  those that completely dissociate

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Sec 10.4 Strengths of Acids

• Table 10.2 Page 237 for Acids
• Strong Acids HCl, HBr, HI, HNO3, H2SO4, HClO4
• Table 10.3 Page 238 for bases
• Strong Bases LiOH, NaOH, KOH, Ba(OH)2
• Weak acids and bases nearly anything (that
  qualifies as acid or base) not listed above

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Sec 10.4 Strengths of Acids

• Finally, remember that strength and concentration
  are two different things
• We could have a strong acid at low concentration
• Or a weak acid at high concentration

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Sec 10.5 Ionization Constants

• General form of acid dissociation
• HA H2O A- H3O
• So Keq for this expression would be
• Keq A-H3O / HAH2O
• But water is a liquid and reduces to 1
• We write Ka as the acid ionization constant
• Ka A-H3O / HA

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Sec 10.5 Ionization Constants

• The pH or pKa are expressed in the logarithmic
  scale
• pH -log H3O and
• pKa -log Ka
• The larger the value of Ka, the stronger the acid
  strength. Why is this?
• Ka and pKa are inverse

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Sec 10.5 Ionization Constants

• Table 10.4 Page 238 Ka Values

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Sec 10.6 Salts

• A salt is an ionic compound that dissolves in
  water to give negative and positive ions
• A salt can release ions which can be acidic or
  basic, such as NaOH
• Furthermore, salts form as the result of an
  acid-base neutralization reaction(Sec 10.7)

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Sec 10.7 Acid-Base Neutralization

• When acids and bases are mixed, the H and OH-
  ions combine to form H2O, a neutral substance.
• This reaction is called neutralization
• The products of a neutralization are water and a
  salt
• For di or triprotic acids, more molecules of base
  are needed to neutralize each molecule of acid

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Sec 10.8 Self Ionization of Water

• Remember that water is an amphoteric substance,
  it can act either as an acid or a base
• Therefore two molecules of water can react to
  self ionize, producing H3O and OH-

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Sec 10.8 Self Ionization of Water

• For pure water at room temperature,
• Kw has a value of 1.0x10-14

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Sec 10.8 Self Ionization of Water

• Because in pure water H3O and OH- are equal, the
  concentrations H3O and OH- are equal to the
  square root of Kw, or 1.0x10-7
• Kw H3OOH- and
• H3O OH- so therefore
• H3O Square root of Kw

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Sec 10.8 Self Ionization of Water

• However, it is true that for any solution of
  water Kw H3OOH-
• This means that if the H3O increases in
  concentration, then OH- must decrease
• If H3O is higher than 1.0x10-7 the solution is
  acidic
• If H3O is lower than 1.0x10-7 the solution is
  basic

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Sec 10.8 Self Ionization of Water

• Figure 10.9 Page 244

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Sec 10.9 The pH and pOH scales

• We convert H3O concentrations into the pH
  scale to make things easier, the pH scale will
  always be a positive number between 0 and 14
• pH -log H3O
• If the concentration increases, the pH will
  decrease
• Therefore a solution with a pH less than 7 is
  acidic. pH greater than 7 is basic

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Sec 10.9 The pH and pOH scales

• Table 10.5 Outlines the significance of pH
  numbers in relation to the solution

Page 244
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Sec 10.9 The pH and pOH scales

• pOH is handled in exactly the same way as pH.
• pOH -log OH-
• Furthermore, pH pOH 14
• Because pKw -log Kw -log (1.0x10-14)
• Example
• If the H3O of a solution is 3.0x10-2 what are
  the pH and pOH ?

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Sec 10.9 The pH and pOH scales

• The pH of Common Materials

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Sec 10.12 Buffers

• Buffer is a solution that is able to change only
  minimally when acid or base is added to it
• Buffers are critical for physiological systems
• Commonly, buffers are made up of a weak acid and
  its conjugate base

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Sec 10.12 Buffers

• If you have a buffer and add acid, the acid
  reacts with the conjugate base present
• If you have the same buffer and add base, the
  base reacts with the weak acid
• In either case, the buffer resists change in pH
  by neutralizing the added substance

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Sec 10.12 Buffers

• Any buffer system has some limit to the amount of
  acid or base you can add
• This limit is called buffer capacity
• Once you have reached the buffer capacity, the
  solution will no longer resist changes to pH
  because the buffer molecules are used up
• The buffer systems found in human blood are
  carbonate and phosphate

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Sec 10.12 Buffers

• Blood is a critical buffer solution essential for
  human life.
• As seen below and explained further on page 257,
  the range of safe pH is extremely narrow

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Sec 10.12 Buffers

• Some Common Buffer Salts
• Table 10.8 Page 252

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Sec 10.13 Henderson-Hasselbalch

• The Henderson-Hasselbalch equation can be used to
  make a buffer solution at a specific pH

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Sec 10.15 Acid-Base Titration

• The purpose of a titration is to determine the
  concentration of an unknown solution.
• The end point of an acid-base titration is when
  there are equal moles of acid and base present
• We can use M1V1 M2V2 to calculate the
  concentration of the unknown
• Often, an indicator is used to signal the end
  point

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Sec 10.15 Acid-Base Titration

• Indicator, is a substance that changes color over
  a certain pH range

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Sec 10.15 Acid-Base Titration

• The setup for a titration
• Figure 10.16 page 258

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Sec 10.15 Acid-Base Titration
The indicator shows that the end point has been
reached. Figure 10.17 Page 259
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Problems

• Assigned problems pages 261-265