On Tap for Today—BONDS

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On Tap for TodayBONDS

• Not the pieces of paper issued by the Gvmnt
• But because atoms graduate to molecules

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Chapter 7 Pre-Review

• This chapter has the greatest long term impact on
  success in chemistry (particularly Organic).
• Organic topics covered, bonding, nomenclature
  structures
• First we cover all types of bonds, focus on
  covalent and polar covalent
• It really pays to thoroughly understand this

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End of Public Service Anoncmnt

• Overview of Chemical Bonds (types, Lewis,
  octets)extension of whats known
• Ionic bondingstrongest of all chemical
  bondsbased almost entirely on ECs
• Covalent bonds (this is the organic)
• Polar covalent/electronegativity/etc
• Lewis structures (biggest lesson to learn)

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More on Lewis Structures

• Strategy for writing them (HINT, as much ART as
  it is SCIENCE)
• Skeletons, resonance, central atoms, of bonds,
  formal charge, all in due time
• Exceptions (there are always exceptions)
• Bond length/strength data

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Lewis structures continued

• Also covered term octet rulewhich states atoms
  tend to gain or lose electrons in order to adopt
  a noble gas configuration
• Remember, noble gases have ns2np6
• Start with Lewis symbols (dots)

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Covalent BondingLewis Structutres

• Covalent bondsatoms share electrons
• Number of dots of valence electrons
• Number of unpaired electrons bonds
• C 4
• N / B 3
• O 2
• F 1

Already paired
Unshared electron
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• Determine the number of valence electrons
• Draw a skeletal structure (takes practiceuse
  table below)
• Distribute excess electronsmake octet
• Excess electrons go to central atom,
  deficiencymultiple
• Calculate formal charge of each atom (table below
  helps)

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Number of valence electrons

• The reason Ive stressed Group/Valence electron
  trends starts to make sense now. We need to
  know
• HOW MANY ELECTRONS does a molecule have?
• Each atom brings valence electrons to the table,
  so to speak
• Anions ADD electrons (more negative charge)
• Cations subtract electrons (less es)
• I KNOW THIS SOUNDS DUMBBUT ADD THEM UP
• ALWAYS!! Its important!! If you miss a lone
  pair, you mess up geometry (later)so just GET
  USED TO THIS STEP!!

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That was step 1

• Now that you know the of bonds an atom
  typically formsstep two is drawing the bare
  bones structure (skeletal)
• This is more art than science. BUT, here are
  some tips
• H is always a terminal atom (forms only 1 bond)
  NEVER gt 1!!!
• Central atoms normally lowest electronegativity
• Terminal atoms normally higher electronegativity,
  F, ALWAYS
• Symmetrical/compact structures better

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Electronegativity Trends
I HHAATTEE memorization, but you should remember
the big four elements N, O, F, Cl.
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Formal ChargeQuick and Dirty

• Books give a formula for calculating formal
  charge, but those are pretty clunkysomething
  simpler
• Count all bonds (as 1) and lone pair electrons (1
  for each electron)
• Btwthiocyanate? Replace the O with an S

O C N valence es (for element) 6 4 5 Number
es (see counting) 6 4 6 Subtract 2 from
1 0 0 -1 Formal Charge resides on Nitrogen
Here, formal charge is on the oxygen. Are
these resonance forms?
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Formal Charge
Usually, the most plausible Lewis structure is
one with no formal charges
When formal charges are required, they should be
as small as possible and negative formal charges
should appear on the most electronegative atoms
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Resonance Delocalized Bonding
Resonance theory states that whenever a molecule
or ion can be represented by two or more
plausible Lewis structures that differ only in
the distribution of electrons, the true structure
is a composite, or hybrid, of them Resonance
structures
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Molecules that Dont Followthe Octet Rule
Molecules with an odd number of valence electrons
have at least one of them unpaired and are called
free radicals
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Molecules that Dont Followthe Octet Rule
Some molecules have incomplete octets. These are
usually compounds of Be, B, and Al, generally
have some unusual bonding characteristics, and
are often quite reactive
EOS
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Molecules that Dont Followthe Octet Rule
Some compounds have expanded valence shells,
which means that the central atom has more than
eight electrons around it
EOS
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Molecules that Dont Followthe Octet Rule
An expanded valence shell may also need to
accommodate lone-pair electrons as well as
bonding pairs
EOS
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Valence Shell Electron Pair Reupulsion

• Fancy way of saying electrons repel each other.
• VSEPR predicts shape based on maximizing distance
• Number of Electron Groups dictates the shape a
  molecule adopts
• 2 Electron Groupslinear
• 3 Electron GroupsTrigonal planar
• 4 EGs (See Amy or Ernie)
• 5 EGs trigonal bipyramidal
• 6 OH (six bonds? OCTA hedral? What gives?

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Electron Groups defined

• Helps to define electron group
• EGs are ANY group of electrons, bonding, lone
  etc.
• Single bond
• Double bond
• Triple bond
• Lone pair of electrons
• Radical (one e)
• Remember, it maximizes the repulsion between ALL
  electrons
• EGGs ARE important though

• HOWEVER, we more often care about the shape of
  the MOLECULE bonding pairs only (lone/radicals
  be damned)
• Molecular Group Geometry describes the
  orientation of JUST the bonding pairs
• More accurately statedthe bonding GROUPS
• Maddeningly enough, some molecules have LONE
  pairs, so what to do?

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Shapes of simple structures

• Two and Three EGs
• Little variation in the shape
• Two is always linear
• Three has two options
• Bent, or trigonal planar

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4-CoordinateOrganic

• Not really going to drag this out much, you ALL
  should know this pretty well
• Methane, ammonia and water, the unholy trinity?
• 40 tetrahedral
• 31 trigonal pyramidal!
• 22 bent (I hate angular)
• Note bond angles!!

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For the uber lazythe summary panel
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Promotionfollowed by hybridztn

• Remember that promotion is energy intensive,
  takes energyBUTnot as much energy saved when
  you can form additional bonds
• Making bonds RELEASES energy, breaking takes E

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Consistent with what weve learned?

• Yesremember that orbitals are mathematical
  equationsprobability of finding an electron
• When you combine several orbitals, the eq changes

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This helps explain methane

• Also important to remember that the number of
  atomic orbitals INTO a hybrid scheme MUST equal
  the number of hybrid orbitals OUT of that scheme
• Previous example means were dealing with an sp3
  orbital (one s and three p orbitals).

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Not all hybrids are bonding orbitals

• In each casemethane, ammonia, and water are all
  sp3 hybridized, but lone pairs of electrons
  occupy some of these orbitals

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What about other schemes

• Yep, can do those too. What about an sp2 scheme
  (one s and two p orbitals)?
• 3 in, 3 out, all at 120 (just like Electron
  group geometry)
• One orbital isnt

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sp Hybridization in Be
with two unused p orbitals.
Two AOs combine to form
two hybrid AOs
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Lets relate Electron Cloud/hybrid geo!

• Note that the EGG and the hybrid orbital geometry
  are the same (which is why its often necessary
  to assign EGG in the first place).

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Hybrid Orbitals andMultiple Covalent Bonds

• Covalent bonds formed by the end-to-end overlap
  of orbitals are called sigma (s) bonds.
• All single bonds are sigma bonds.
• A bond formed by parallel, or side-by-side,
  orbital overlap is called a pi (p) bond.
• A double bond is made up of one sigma bond and
  one pi bond.
• A triple bond is made up of one sigma bond and
  two pi bonds.

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VB Theory for Ethylene, C2H4
The hybridization and bonding scheme is described
by listing each bond and its overlap.
Sigma bonds, and pi bonds
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• Formic acid, HCOOH, is the simplest carboxylic
  acid.
• Predict a plausible molecular geometry for this
  molecule.
• Propose a hybridization scheme for the central
  atoms that is consistent with that geometry.
• Sketch a bonding scheme for the molecule.