Basic Concepts of Bonding

1
Basic Concepts of Bonding

• Chapter 8 and Chapter 9
• Brown and LeMay

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3 Types

• Ionic
• Metal nonmetal
• Electrostatic force
• Transfer of electrons
• Covalent
• Two nonmetals
• Metallic
• Positive nuclei in sea of electrons

3
General Stuff

• Valence electrons involved
• Shown by Lewis dot diagram
• Octet Rule 8 is magic number why?

4
Ionic Bonding and Energy

• Exothermic Process
• Remove electron endothermic
• Gain electron exothermic
• Example ionization energy for Na is more
  positive than electron affinity for chlorine is
  negative, so this appears to be endothermic
• Something else energy is released when ions are
  attracted to each other

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Lattice Energy

• How much energy?
• Measured by the amount of energy it takes to
  break up the crystal, called lattice energy
• Definition energy needed to separate a mole of
  a solid ionic compound into its gaseous elements
• Most greatest charge and smallest radius
• Compare AgCl, CuO, CrN

6
Electron Configuration of ions

• Try oxgen, magnesium, vanadium

7
Polyatomic Ions

• Review

8
Covalent Bonding

• Much more common!
• Electrons repel, nucleus attracts electrons
• Attraction must overcome repulsion
• Shared electrons act as glue to hold nuclei
  together
• Show H2, Cl2 (animation)
• Cant do it with single bonds, try double, then
  triple
• Distance between nuclei decreases as more are
  shared

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Nonpolar Covalent Bond

• Electrons are shared equally
• H2, Cl2 etc
• CH4 (how do you know?)

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Electronegativity

• Based upon the ability of an atom in a molecule
  to attract electrons to itself.
• Scale based upon lots of data
• 0 4
• Fluorine is the greatest, cesium is the least
• Predict by knowing trends

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Trends
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Polar Bond

• Covalent bond in which one nucleus attracts the
  electrons to itself more than the other.
• Electronegativity difference usually between .4
  and 2.0
• If the difference is around 2.0 or greater, the
  bond is considered ionic

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Comparison
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Polar Molecule

• Molecule containing asymmetric polar bonds
• A dipole is established whenever two equal
  charges are separated by a distance
• Dipole moment (µ) is the magnitude of the dipole

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Nomenclature

• Try BaCl2 and NO2
• Compounds made with metals of high oxidation
  numbers may not behave as ionic compounds do.
  Often named as molecular.
• Example TiO2

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Drawing Lewis Structures

• 1. Sum all valence electrons
• 2. Central atom is usually least electronegative
• Arrange with all single bonds (count for 2)
• Complete octets with remaining
• If extra expand octet of central atom
• If not enough double or triple up
• Try NO ion, PO43- ion

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Formal Charge

• How can you decide from several possible
  structures?
• Formal charge is a quick estimate of
  reasonableness (has no chemical reality)
• Equal to ( valence electrons on an atom
  assigned electrons)
• Definition of assigned all unshared half of
  those bonded
• Sums must be 0 formal charge for a compound,
  charge of ion for an ion.

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Practice with CO2

• Try 1 single, 1 triple each way and 2 doubles.
• Evaluate!
• Best 0 for as many atoms as possible,
  negative value for most electronegative atom.
• How is NCO- put together?

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Resonance

• Ozone, O3 do Lewis structure
• Which is correct?
• Does it matter?
• If not, this is resonance
• In reality, bond length is a blend

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Benzene

• C6H6
• Benzene
• Basis of lots of chemistry today

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Exceptions to the Octet Rule

• Odd number, less, or more
• Odd NO
• Less boron (6) and beryllium (4)
• BF3 and BeF2
• Memorize!
• Too many called expanded octet

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Expanded Octet

• Largest class
• Period 3 and beyond because it involves d
  orbitals
• More common for larger central atoms like I
• Try ICl4-

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Bond Strength

• Determined by energy needed to break bonds
• Bond enthalpy ?H for breaking a bond in a mole
  of the gaseous substance
• Show equation
• Designated D(Cl-Cl)
• Average in more complex molecules (C-H)
• All positive, greater value means stronger bond

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Estimating ?H of a reaction

• Use when other values (? Hf) are not available
• Sum of bonds broken sum of bonds formed
• Just an estimation

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Strength of Multiple Bonds

• As number of bonds between two atoms increases,
  strength is greater and distance decreases

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Chapter 9 - Geometry

• VSEPR Model Valence-shell electron pair
  repulsion theory
• Electron domains repel each other and therefore
  arrange each other in space so that they repel
  the least.
• Electron domain single bond, multiple bond, or
  unshared pair of electrons

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Handout

• Domains 2 6
• Learn shapes
• Electron domains for Unshared electrons
• nonbonding electrons have more repulsive force
  and therefore reduce bonding angles slightly
• Also true for doubles and triples compared to
  single bonds

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Expanded Octets

• Trigonal Pyramidal axial and equatorial domains
  are different (show)
• Unshared are the equatorial domains
• Octahedral doesnt matter for first, second is
  opposite

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Larger

• Describe geometry around a specific atom
• Double and triple bonds change geometry that is
  why Lewis structure is not always enough!
• Try acetic acid

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Polarity

• The kind of bonds and arrangement determine if a
  molecule is polar.
• You must have an asymmetric polar bond to make a
  molecule polar
• Try H2O and CCl4

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Isomers

• Molecules with the same number and kind of atom
  but different arrangements and therefore
  different properties.
• Three kinds
• Structural (C-H backbone)
• Functional (alcohol vs. ether)
• Geometric (stereo isomers or mirror images)

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Hybridization

• This is the process of changing the orbitals as
  atoms approach each other
• Example - BeF2 linear (show)
• Page 336 summary
• Reported as sp, sp2, sp3, sp3d, sp3d2

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Sigma and Pi bonds

• Describes orientation of orbitals
• Single or first of double and triple are sigma
  bonds (s)
• Other bonds in a double or triple bond are pi
  (p)
• Lets try formaldehyde, CH2O
• structure, resonance, isomer, hybridization
  around C, sigma and pi bonds of C