CH339K

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CH339K

• Lecture 2

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Bonding

• Covalent
• Ionic
• Dipole Interactions
• Van der Waals Forces
• Hydrogen Bonds

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Covalent Bonds

• Electrons form new orbitals around multiple
  atomic nuclei
• Bond energy results from electrostatic force
  between redefined electron cloud and nuclei
• Strong typically 150 400 kJ/mol

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Ionic Interactions

• Energy from non-directional force between ions
• Biomolecules frequently have large numbers of
  charged groups
• Charge-charge interactions stabilize intra- and
  intermolecular structures
• Coulombs Law
• Energy drops off as function of distance between
  charges

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Dipoles

• Fixed dipoles
• Molecules with asymmetric charge distributions
  form dipoles
• Induced Dipoles
• One dipole can induce a charge in an adjacent
  molecule

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van der Waals Interactions

• Technically, all induced dipole interactions are
  van der Waals interactions
• Biochemists usually mean induced dipole-induced
  dipole (London Dispersion) forces
• Any atom will have an uneven distribution of
  charge at any given instant

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Van der Waals (cont.)

• That temporary dipole will induce a dipole in
  adjacent atoms
• This results in a net attractive force between
  atoms
• Force is weak - .5 to 2 kJ/mol
• Net biochemical effect molecules that FIT
  together STICK together.

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Van der Waals (cont.)

• If you live in Central Texas, you see van der
  Waals forces in action every summer night

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Hydrogen Bonds

• Hydrogen Bonds form between
• A hydrogen covalently bound to an electronegative
  atom
• Another electronegative atom

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Hydrogen Bonds (cont.)

• The group to which the hydrogen is covalently
  bound is the donor.
• The other group is the acceptor.
• Donors
• -OH, -NH2, -SH (lesser donor)
• Acceptors
• -N, O, -O

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Hydrogen Bonds (cont.)

• Hydrogen bonds are not just electrostatic
  partially covalent
• Therefore, they are directional
• Intermdiate strength 5 10 kJ/mol

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Water Structure
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Water Forms Clusters in Solution
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Hydrophobic Effect
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Water

• Water
• Has a high specific heat
• Has a high heat of vaporization
• Is an excellent solvent for polar materials
• Is a powerful dielectric
• Readily forms hydrogen bonds
• Has a strong surface tension
• Is less dense when it freezes (i.e. ice floats)

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Acids and Bases

• Definitions
• Arrhenius
• Bronsted-Lowry
• Lewis

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Conjugate Pairs

• Every acid has its conjugate base
• Every base has its conjugate acid

Conjugate Acid Conjugate Base
H3C - COOH H3C-COO-
NH4 NH3
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Acids and bases pH

• Water ionizes

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Typical pH Values
Substance pH
Stomach acid 1.5 - 2.5
Coca-cola 2.5
Human saliva 6.5
Human blood 7.5
Human urine 5 - 8
Oven cleaner 14
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Acids and Bases

• Water thus acts as both a weak acid and a weak
  base
• (A Strong acid is one that dissociates completely
  in water a weak acid is one that doesnt.)
• All biochemically significant acids and bases are
  weak (except for HCl stomach acid)

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Acids and Bases

• Just like water, a weak acid has an ion product,
  the Ka
• For the weak acid HA
• Therefore

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Acids and Bases

• Kas for weak acids range over several orders of
  magnitude
• They are generally small
• More convenient to define
• pKa -log Ka
• Just like pH -logH

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Typcal Kas and pKas
Acid Ka pKa
Acetic 1.8 x 10-5 4.74
Formic 1.7 x 10-4 3.77
Benzoic 6.5 x 10-5 4.19
Carbonic 4.3 x 10-7 6.37
Imidazole 2.8 x 10-7 6.55
Phenol 1.3 x 10-10 9.89
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pH for Strong Acids

• Since a strong acid dissociates completely
• pH -log(Acid)
• For a 0.1 M (100 mM) solution of HCl,
• pH -log(0.1) 1

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pH for Weak Acids

• Whats the pH of a 100 mM solution of Acetic Acid?

H 0.00134 M
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Shortcut

• The quadratic solution is a pain, but we can
  approxmate

H 0.00134 M
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Titrating a Strong Acid

• 10 ml of an HCL sln
• Titrate with 0.5 M NaOH
• OH- H gt H2O
• Takes 8.5 ml NaOH to bring solution to neutrality

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Titrating a Weak Acid

• Titrating .1 M Hac
• Initial pH is 2.88 instead of 1
• Little change until large amounts of NaOH have
  been added
• Buffering effect
• Caused by equilibrium that exists between a weak
  acid and conjugate base.

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Henderson-Hasselbach Equation
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Predicting pH

• Lets make 1 liter of a solution that is 0.1 M in
  acetic acid ( pKa 4.74 ) and 0.3 M in sodium
  acetate.

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Buffering Effect

• Addition of significant amounts of acid or base
  changes the ratio of conjugate base to conjugate
  acid
• pH changes as the log of that ratio
• Result is resistance to pH change in a buffered
  solution