Bonding

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Bonding
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Background information

• Happen through electrostatic forces
• Studied through physical properties
• -
• Strength of bond studied through _____________

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Covalent Bonds

• Electrons are ____________
• Bonds form when lowest possible energy is
  achieved.
• Distance between two atoms where the energy is
  the lowest determines the bond length.

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Figure 8.1 a b (a) The Interaction of Two
Hydrogen Atoms (b) Energy Profile as a Function
of the Distance Between the Nuclei of the
Hydrogen Atoms
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Polarity in Covalent Bonds

• _____________ sharing of electrons
• Atoms will form partial positive and negative
  charges.
• More electronegative ________________
• Based on electronegativity
• Ability of atom to attract e- to itself.
• Measured by comparing measured bond energy with
  expected bond energy.
• Increases from left to right and up the periodic
  table.

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Electronegativity Values
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The Relationship Between Electronegativity and
Bond Type
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Examples

• Order the following bonds according to polarity
• H-H, O-H, Cl-H, S-H, F-H
• Ans

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Dipole Moments

• A molecule with center of charge, (partial
  charges) is dipolar or has a dipole moment.
• Represent with an arrow pointing to the negative
  charge center and tail indicating charge.
• Ex H2O NH3

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Dipole Moments

• Some molecules have polar bonds but not dipole
  moments.
• Molecule is arranged so bond polarities cancel.
• Ex CO2, SO3, CH4

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Ionic Bonds

• Atoms lose/gain electrons to form bonds.
• -
• Energy interaction between a pair of ions can be
  calculated using Coulombs law
• Where
• E
• Q1/Q2
• r

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Examples

• Which has the most exothermic lattice energy
  NaCl or KCl?
• Ans

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Examples

• In solid sodium chloride the distance between the
  centers of Na and Cl- is 2.76 Å. What is the
  ionic energy per pair of ions? (1 nm 10 Å)
• Ans
• Note Can be used to calculate repulsive charges
  when two like charges are brought together.
  Calculate value will be positive.

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Size of Ions

• Review Size increases down and to the right.
• For _______________ ions (contain same of
  _______) consider of e- to of p ratio
• As of p increases, e- experience a greater
  attraction smaller ion.

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Examples

• Put these in order of increasing size
• F-, Mg2, O2-, Al3, Na
• Ans

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Trends

• Bigger charges _______________
• Smaller charges _______________
• Charges close together __________
• Charges Far apart _____________

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Binary Ionic Compounds

• The strength of the ion attraction in an ionic
  bond is indicated by its
• Lattice energy-.
• If the energy change is negative _____________
• If the energy change is positive ______________

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Figure 8.9 The Energy Changes Involved in the
Formation of Lithium Fluoride from Its Elements
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Lattice Energy Calculations

• Calculating energy can be done in a modified
  version of Coulombs law
• Where
• k
• Q1/Q2
• r

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Size of Ions

• Review Size increases down and to the right.
• For isoelectronic ions (contain same of e-)
  consider of e- to of p
• As of p increases, e- experience a greater
  attraction smaller ion.

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Examples

• Put these in order of increasing size
• F-, Mg2, O2-, Al3, Na
• Ans O2-, F-, Na, Mg2, and Al3 because all
  of the ions have the same number of e- but O2-
  has the least of p and Al3 has the least .

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Intermolecular Forces

• Forces that exist within the molecule-to hold it
  together.
• Vary in strength
• Types
• -
• -
• -

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Dipole-Dipole

• Occur in ________________ molecules
• Greater polarity
• Overall all dipole-dipole moments relatively
  weak.
• -
• -
• Most compounds held together by dipole-dipole
  forces are __________________ at room temp.

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London dispersion forces

• Occur between _______________ molecules.
• Weak attractions occur because of random e-
  motion on the atom within the molecule.
• Moving e- can lead to spontaneous polarity.
  (____________________) Acts as a weak dipole.
• More e-
• Weak attraction
• -
• -
• Generally _____________ at room temp.

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Hydrogen Bonds

• Similar to ________________
• Hydrogen end is attracted to negatively charged
  end of a molecule containing extremely
  electronegative element
• -
• Much stronger than dipole-dipole because after
  giving up e-, nucleus is unshielded.
• -

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Examples

• Arrange the following in order of increasing BP
• NH3, C2H6, SO2
• C2H6, CH4, C4H10
• Ans

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Water

• Why is water unlike other molecules by having a
  less dense solid form than liquid form?
• Ans

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Other properties

• Vapor Pressure-
• Stronger IMF
• Viscosity
• Strong IMF
• Heat of Vaporization
• Stronger IMF
• Enthalpy of Fusion
• Stronger IMF