vsepr
TRANSCRIPT
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VSEPR
Valence Shell Electron Pair Repulsions
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Covalent Bond: A Model
• Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit
• They result from the tendency of a system to seek its lowest possible energy
• Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms
Note: The next three slides will repeat at the end. This is preliminary intro info that may make more sense at the end.
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Example: Methane
• 1652 kJ of energy are required to break a mole of methane into separate C and H atoms
• OR 1652 kJ of energy are released when one mole of methane is formed from one mole of C atoms and four moles of H atoms
• Methane is therefore a stable molecule relative to its stable atoms
• Since there are four H atoms arranged around the central C, it is natural to envision four individual attractions between C and H (bonds)
• Each bond has an associated bond energy, found by dividing the total energy by four (1652/4 = 413 kJ)
• The positive Bond Energy value indicates the energy required to break the bond between C and H atoms
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Bonding Model• Models originate from our observations of
the properties of nature• Atoms can form stable groups by sharing
electrons, shared electrons give a lower energy state because simultaneously attracted to two nuclei
• Remember: Models are human inventions that allow us to explain and predict. A model is a useful way of thinking; they include simplifications and assumptions. A model does not equal reality.
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Covalent Bonds
• Electron pair(s) – shared between two atoms– attracted to both nuclei
• Location of a single shared pair– Directly between two nuclei– Maximizes attractions with shortest distance between
two positive nuclei– Minimizes repulsions with negative electrons between
positive nuclei that would repel one another
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Multiple covalent bonds around the same atom determine the shape
• Negative e- pairs with same charge repel each other
• Repulsions push the pairs as far apart as possible
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Single bonds
• Sigma bond ө• Overlap of orbitals allow electron pair to be
shared between the two atoms
• Electron pair shared directly between two nuclei
• Only one pair may be shared in this space - just as only one pair of electrons may occupy a single atomic orbital
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Double and Triple Bonds
• Pi bonds π• Since the space between the nuclei is
occupied, e- pair is shared above and below the plane or front and back
• Overlap of p-orbital lobes allow for this sharing above and below OR front and back
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Lewis Structures
• Drawn to show the bonds between the atoms in the structure
• Only shows whether single, double or triple bonds
• Does not show the shape
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Lewis Structure
• Represents the arrangement of valence electrons among atoms in the molecule
• Rules based upon observations of thousands of molecules, which show that in most stable compounds the atoms achieve noble gas configurations
• Duet Rule – hydrogen stable with only a pair of e-
• Octet Rule – other atoms stable with 4 pairs of e-
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Rules for Drawing Lewis Structures
1) Count the number of valence electrons2) Draw the skeleton structure- the central is
generally listed first in formula3) Distribute electrons to give each atom a stable
octet4) Reconcile # e-
a) Do you have enough electrons? You may need to use double or triple bonds.
b) Do have too many electrons? You may need to explain the octet, but only if empty d-orbital available
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Determine # Valence e- from column #
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Electron Clouds repel each other, thus structure around an atom is determined
principally by minimizing repulsions
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2 electron pairs (2 EP) around central atom
• Two clouds pushed as far apart as possible– Greatest angle possible 180º– LINEAR shape
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3 electron pairs (3 EP) around central atom
• Three clouds pushed as far apart as possible– Greatest angle possible 120º– TRIGONAL PLANAR shape
(3) (flat)
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4 electron pairs (4 EP) around central atom
• Four clouds pushed as far apart as possible– Greatest angle no longer possible
in two dimensions– Requires three-dimensional– TETRAHEDRAL shape
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Orbital Hybridization #1
• Atomic orbitals such as s and p are not well suited for overlapping and allowing two atoms to share a pair of electrons
• Remember: best location of shared pair is directly between two atoms
• e- pair spends little time in best location– With overlap of two s-orbital– With overlap of two p-orbitals
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Orbital Hybridization #2
• Hybrid orbitals (cross of atomic orbitals)• Remember: The pink flower hybrid cross
of the red and white flower• Hybrid orbitals
– Shape more suitable for bonding• One large lobe and one very small lobe• Large lobe oriented towards other nucleus
– Angles more suitable for bonding• Angles predicted from VSEPR
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Overlap of two s-orbitals
NOT A GOOD LOCATION-
far from one nucleus
Note: shared in this overlap the e- pair would spend most of the time in an unfavorable location
GOOD SPOT between both nuclei
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Overlap of two p-orbitals
One atom & its p-orbital The other atom & its p-orbital
represents the nucleus
BAD location far from other nucleus GOOD SPOT
between both nuclei
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Hybrid Orbitals yield more favorable shape for overlap
• Atomic orbitals are not shaped to maximize attractions nor minimize repulsions
• Hybrid orbital shape – One large lobe oriented towards other atom– Notice the difference in this shape compared
to p-orbital shape
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Angles and Shape
• Atomic orbitals are not shaped to maximize attractions nor minimize repulsions
• BUT the angles are also not favorablep-orbitals are oriented at 90º to
each otherOther angles are required 180º,
120º or 109.5º
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Orbital Hybridization #3
• Each e-pair requires a hybrid orbital• If two hybrid orbitals required than two atomic
orbitals must be hybridized, an s and a p orbital forming two sp orbitals at 180º
sp hybrids
2 EP 4 EP3 EP
sp2 hybrids sp3 hybrids
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Electron-Pair Geometryvs
Molecular Geometry
• Electron-pair geometry– Where are the electron pairs– Includes
• bonding pairs (BP) – shared between 2 atoms• nonbonding pairs (NBP) – lone pair
• Molecular geometry– Where are the atoms– Includes only the bonding pairs
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Examples of 3 EP• 3 BP + 0 NBP = 3 EP
– 3 EP = EP geom is trigonal planar– All locations occupy by an atom, – so molecular geometry is also trigonal planar
• 2 BP + 1 NBP = 3 EP– 3 EP = EP geom is trigonal planar– Only two bonding pairs– One of the locations is only lone pair of e-– so molecular geometry is bent
O
O O
N
O O
O
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Carbonate Ion (CO32-)
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Nitrate Ion (NO3-)
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Nitrite Ion (NO2-)
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Examples of 4 EP• 4 BP + 0 NBP = 4 EP
– Both EP geom and molecular geom– tetrahedral
• 3 BP + 1 NBP = 4 EP– 4 EP so EP geom is tetrahedral– Molecular geom is TRIGONAL PYRAMIDAL – No atom at top location
• 2 BP + 2 NBP = 4 EP– 4 EP so EP geom is tetrahedral– Molecular geom is BENT – no atoms at two locations
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4 BP + 0 NBP = 4 EP
TETRAHEDRAL
Cl Cl
Cl
Cl
S
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3 BP + 1 NBP = 4 EP
TRIGONAL PYRAMIDAL
NH H
H
●●
lone pair of e-
NBP
H H
H
N
107
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2 BP + 2 NBP = 4 EP
BENT
O H
H
●●
lone pair of e-
NBP H
H
O
●●
lone pair of e-
NBP
104.5
H
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Hydronium Ion (H3O+)
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Ammonia Molecule (NH3)
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Summary of 4 EP
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Exceptions to Octet Rule
• Reduced Octet – H only forms one bond- only one pair of e-– Be tends to only form two bonds
• only two pair of e-
– B tends to only form three bonds• only three pair of e-
• Expanded Octet– Empty d-orbitals can be used to
accommodate extra e-– Elements in the third row and lower can expand– Up to 6 pairs of e- are possible
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Lewis Structures in Which the Central Atom Exceeds an Octet
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5 EP
Trigonal bipyramidal
Orbital hybridizationRequires 5 hybrid orbitals
So, 5 atomic orbitals required
sp3d
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Trigonal planar shape
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5 EP = trigonal pyramidal
molecular geometry5 BP + 0 NBP = 5 EP 4 BP + 1 NBP = 5 EP
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5 EP = trigonal pyramidal
molecular geometry 3 BP + 2 NBP = 5 EP 2 BP + 3 NBP = 5 EP
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6 EP
Octahedral
Orbital hybridizationRequires 5 hybrid orbitalsSo, 5 atomic orbitals requiredsp3d
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6 EP = octahedral
6 BP
+0 NBP
6 EP
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6 EP = octahedral
5 BP+1 NBP 6 EP
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6 EP = octahedral
4 BP+2 NBP 6 EP
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Summary: Molecular Geometry of Expanded Octets
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Summary of EP Geometry
2 EP 3 EP 4 EP 5 EP 6 EP
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Summary of EP Geometry
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Predict the geometry, angles and orbital hybridization
Predict the geometry, angles and orbital hybridization
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Covalent Bond: A Model
• Chemical bonds can be viewed as forces that cause a group of atoms to act as a unit
• They result from the tendency of a system to seek its lowest possible energy
• Bonds occur when collections of atoms are more stable (lower in energy) than the separate atoms
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Example: Methane
• 1652 kJ of energy are required to break a mole of methane into separate C and H atoms
• OR 1652 kJ of energy are released when one mole of methane is formed from one mole of C atoms and four moles of H atoms
• Methane is therefore a stable molecule relative to its stable atoms
• Since there are four H atoms arranged around the central C, it is natural to envision four individual attractions between C and H (bonds)
• An average bond energy associated with each bond is found by dividing the total energy by four (1652/4 = 413 kJ)
• The positive Bond Energy value indicates the energy required to break the bond between C and H atoms
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Bonding Model• Models originate from our observations of
the properties of nature• Atoms can form stable groups by sharing
electrons, shared electrons give a lower energy state because simultaneously attracted to two nuclei
• Remember: Models are human inventions that allow us to explain and predict. A model is a useful way of thinking; they include simplifications and assumptions. A model does not equal reality.
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Bond Energy and Enthalpy
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