Types of chemical bonds

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Types of chemical bonds

• Bond Force that holds groups of two or more
  atoms together and makes the atoms function as a
  unit.
• Example H-O-H
• Bond Energy Energy required to break a bond.
• Ionic Bond Attractions between oppositely
  charged ions.
• Example Na Cl-

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Types of chemical bonds

• Ionic Compound A compound resulting from a
  positive ion (usually a metal) combining with a
  negative ion (usually a non-metal).
• Example M X- ? MX
• Covalent Bond Electrons are shared by nuclei.
• Example H-H
• Polar Covalent Bond Unequal sharing of electrons
  by nuclei.
• Example H-F
• Hydrogen fluoride is an example of a molecule
  that has bond polarity.

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Lewis structures

• Lewis Structure Representation of a molecule
  that shows how the valence electrons are arranged
  among the atoms in the molecule.
• Bonding involves the valence electrons of atoms.
• Example Na?
  H-H

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Lewis structures of elements

• Dots around elemental symbol
• Symbolize valence electrons
• Thus, one must know valence electron configuration

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Lewis Structures of molecules

• Single Bond Two atoms sharing one electron pair.
• Example H2
• Double Bond Two atoms sharing two pairs of
  electrons.
• Example O2
• Triple Bond Two atoms sharing three pairs of
  electrons.
• Example N2
• Resonance Structures More than one Lewis
  Structure can be drawn for a molecule.
• Example O3

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Rules for Lewis structures of molecules

• Write out valence electrons for each atom
• Connect lone electrons because lone electrons are
  destabilizing
• Become two shared electrons
• Called a bond
• Check to see if octet rule is satisfied
• Recall electron configuration resembling noble
  gas
• In other words, there must be 8 electrons (bonded
  or non-bonded) around atom
• Non-bonded electron-pair
• Called lone pair

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Lets do some examples on the board

• H2
• Duet rule
• F2
• Octet rule
• O2
• N2

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Lewis structures

• Example
• Write the Lewis Structure for the following
  molecules
• H2O
• CCl4
• Where does the carbon go why?
• PH3
• H2Se
• C2H6

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Lewis structures continued

• CO2
• C2H4
• C2H2
• SiO2

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Polyatomic ions

• If positive charge on ion
• Take away electron from central species
• If negative charge on ion
• Add electron to central species
• Example
• H3O

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Your turn

• NH4
• ClO-
• OH-

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Resonance structures

• When structures can be written in more than one
  way
• O3
• Actual molecule is in-between
• Resonance hybrid
• Another example
• HCO3-
• What would its resonance hybrid look like?

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Practice

• NO2-
• NO3-

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Formal Charge

• Charge calculated on atom based on Lewis
  structure
• Yields best Lewis structure of competitors
• FC VE - LE ½(BE)
• Rules
• Sum of all FCs must equal to charge on species,
  if any
• Smaller or zero FCs on atoms better than large
  FCs
• Negative FC should be on most electronegative
  species

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Examples

• HBr
• FC on H 1-0 ½ (2) 0
• FC on Br 7 6 ½ (2) 0
• Net sum of FCs charge on ion 0
• OH-
• FC on O 6 6 ½(2) -1
• FC on H 1 0 ½(2) 0
• Net sum of FCs charge on ion -1

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Practice

• H2O2
• H3O

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Aberrant compounds

• Odd-electron species
• NO
• NO2

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Aberrant compounds

• Incomplete octet
• BH3

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Aberrant compounds

• Expanded octet
• Some central atoms can exceed an octet
• Third period and higher elements can do this
• E.g., Al, Si, P, S, Cl, As, Br, Xe, etc.
• d-orbitals can accommodate extra electrons

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Examples

• AsI5
• XeF2

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Practice

• SCl6
• XeF4

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Aberrant compounds

• Write this out
• SO42-
• Can we reduce the formal charges?
• If so, how?
• We can also find the average FC
• Lets take a look

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Aberrant compounds

• Its OK to expand the octet for those atoms that
  can take it in order to lower FCs

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Practice

• SO32-
• PO33-
• SO2
• SO3
• H2SeO4

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Electroneutrality principle

• Electrons distributed so that charges on atoms
  are closest to zero
• If - charge present, should be on most
  electronegative atom
• (so, charge should be on least
  electronegative atom)
• Good for deciding which resonance structure is
  best
• Example OCN-

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Electronegativity

• Electronegativity The relative ability of an
  atom in a molecule to attract shared electrons to
  itself.
• Example Fluorine has the highest
  electronegativity.
• Similar electronegativities between elements give
  non-polar covalent bonds (0.0-0.4)
• Different electronegativities between elements
  give polar covalent bonds (0.5-1.9)
• If the difference between the electronegativities
  of two elements is about 2.0 or greater, the bond
  is ionic

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Electronegativity

• Example
• For each of the following pairs of bonds, choose
  the bond that will be more polar.
• Al-P vs. Al-N
• C-O vs. C-S

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Dipole moment

• Dipole Moment
• A molecule that has a center of positive charge
  and a center of negative charge
• Will line up on electric field
• In Debye units
• 1 D 3.34 x 10-30 C ? m

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Examples

• F2
• CO2
• H2O
• NH3
• BF3
• CCl4

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Molecular polarity

• Net-dipole moment leads to molecular polarity
• Thus the following two that have net-dipole
  moments are polar
• H2O
• NH3

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Molecular structure

• Molecular Structure or geometric structure
  refers to the three-dimensional arrangement of
  the atoms in a molecule.
• Bond Angle The angle formed between two bonds in
  a molecule.

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Molecular structureVSEPR

• The VSEPR Model The valence shell electron pair
  repulsion model is useful for predicting the
  molecular structures of molecules formed from
  nonmetals.
• The structure around a given atom is determined
  by minimizing repulsions between electron pairs.
• The bonding and nonbonding electron pairs (lone
  pairs) around a given atom are positioned as far
  apart as possible.

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Molecular StructureVSEPR

• Steps for Predicting Molecular Structure Using
  the VSEPR Model
• 1. Draw the Lewis structure for the molecule.
• 2. Count the electron pairs and arrange them in
  the way that minimizes repulsion (that is, put
  the lone pairs as far apart as possible).
• 3. Determine the positions of the atoms from the
  way the electron pairs are shared.
• 4. Determine the name of the molecular structure
  from the positions of the atoms.

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Example

• Br2
• CO2
• CF4
• PF3

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Your turn

• NH4
• XeF4
• AsI5
• SF3
• I3 -