Scientists Resonance
have observed that certain compounds with uneven bonding (double bond on one side, single bond on the other for example) have bond lengths that differ from the calculated value
Ozone: The O=O bond length should be 121 pm The O-O bond length should be 132 pm Actual length of both bonds is between these 2 values!
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Resonance
Linus Pauling proposed the theory of resonance to solve this problem Resonance structures are structures that represent the bonding in a molecule or ion when a single Lewis structure fails to accurately describe the actual electronic structure The `real' structure is a hybrid of the resonance structures The predominant resonance structure is on with the formal charges of all atoms as close to zero as possible
(Nobel Laureate for Chemistry AND Peace)
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Resonance
Let's look at the Carbonate ion
Or the structure of Benzene (Classic Resonance structure)
Note Double Headed Arrows!
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Exceptions to the Octet Rule
(Compounds with an atom that has fewer than 8 valence electrons)
Boron (Group 3A metalloid) Has 3 valence electrons Will form 3 covalent bonds and "have" six valence electrons
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Coordinate Covalent Bond (aka: Dative Bond)
Many boron compounds are HIGHLY reactive The 4th valence pair for Boron can be provided by another compound
Note the bonding arrow from Nitrogen to Boron This implies that there is a dative bond and Nitrogen is supplying a bonding pair to Boron
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Compounds with Atoms Possessing more than 8 valence electrons
Elements in the 3rd or higher period may sometimes form compounds with more than 8 valence electrons
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Second period elements will not form compounds that give them more than 8 valence electrons
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Why do these elements do that?
Elements in the 3rd period have empty orbitals in the d subshell that will allow them to accommodate up to 12 electrons in the valence shell. Be aware that they do exist and know why they behave as they do. Silicon, Phosphorous, Chlorine are good examples
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Free Radicals
What are they? We've heard them described in commercials, scientific literature and songs. Compounds with an unpaired electron are called `free radicals' VERY REACTIVE
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The Shapes of Molecules
Molecular shapes can be predicted by noting how many bonds and electron pairs surround individual atoms and applying what is called the valence-shell electron-pair repulsion (VSEPR) model. The basic idea of the VSEPR model is that the negatively charged clouds of electrons in bonds and lone pairs repel each other, and keep as far apart as possible.
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Molecular Geometries
Trigonal Bipyramidal and octahedral geometries are only expected when the central atom is in period 3 or higher. Why?
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Linear Geometry
Usually trinary compounds with double bonds or Boron compounds Central atom in a trinary compound w/o double bonds MUST NOT have an octet in order to see this geometry
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Trigonal Planar Geometry
Central atom MUST NOT have an octet of electrons
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Tetrahedral Geometry
Most molecules with quaternary compositions have this geometry
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Trigonal Planar Geometry
Most molecules with quaternary compositions have this geometry
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Trigonal Bipyramidal and Octahedral Geometries
Central atom does not follow the octet rule
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Central Atoms with Lone Pairs AND Bonding Pairs
Now we have to introduce another term: Electron Pair Geometry The Electron Pair Geometry is the geometry taken up by ALL of the valence electron pairs around a central atom The Molecular Geometry is the geometry of the central atom and the atoms surrounding it
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Electron Pair Geometry and Molecular Geometry
The electron pair geometry in ammonia is TETRAHEDRAL so we would expect the H-N-H bond angle to be 109.5, but it is actually less! Lone pairs occupy a larger colume than bonding pairs
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This causes the Lone Pair to squeeze the Bonding Pairs together a bit
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Electron Pair Geometry and Molecular Geometry
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Central Atoms with MORE than 4 Valence Electron Pairs
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There are three step to applying the VSEPR model: Step 1: Draw a Lewis structure of the molecule, and identify the atom whose geometry is of interest. Step 2: Count the number of electron charge clouds surrounding the atom of interest. Step 3: Predict molecular shape by assuming that the charge clouds orient in space so that they are as far from away one another as possible.
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Linear molecules have bond angles of 180. Planar triangular molecules have bond angles of 120. Tetrahedral molecules have bond angles of 109.5.
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Polar Covalent Bonds and Electronegativity
Electrons in a covalent bond occupy the region between the bonded atoms. If the atoms are identical, as in H2 and Cl2, electrons are attracted equally to both atoms and are shared equally. If the atoms are not identical, however, as in HCl, the bonding electrons may be attracted more strongly by one atom than by the other and thus shared unequally. Such bonds are known as polar covalent bonds.
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When charges separate in a neutral molecule, the molecule has a dipole moment and is said to be polar.
The dipole moment is measured in SI units called the Debye (D) 1 D = 3.34x10-30 Cm
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In HCl, electrons spend more time near the chlorine than the hydrogen. Although the molecule is overall neutral, the chlorine is more negative than the hydrogen, resulting in partial charges on the atoms. Partial charges are represented by a - on the more negative atom and + on the more positive atom. The ability of an atom to attract electrons is called the atom's electronegativity. Fluorine, the most electronegative element, assigned a value of 4, and less electronegative atoms assigned lower values.
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Elements at the top right of the periodic table are most electronegative, those at the lower left are least electronegative. Noble gases are not assigned values.
More electronegative atoms surround themselves with negative charge and more electropositive atoms tend to `give' electrons Prentice Hall 2007 26 Chapter Five up in a bond
As a rule of thumb, electronegativity differences of less than 0.5 result in nonpolar covalent bonds, differences up to 1.9 indicate increasingly polar covalent bonds, and differences of 2 or more indicate ionic bonds. There is no sharp dividing line between covalent and ionic bonds; most bonds fall somewhere in-between.
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As a rule of thumb, electronegativity differences of less than 0.5 result in nonpolar covalent bonds, differences up to 1.9 indicate increasingly polar covalent bonds, and differences of 2 or more indicate ionic bonds. There is no sharp dividing line between covalent and ionic bonds; most bonds fall somewhere in-between.
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5.9 Polar Molecules
Entire molecules can be polar if electrons are attracted more strongly to one part of the molecule than to another. Molecules polarity is due to the sum of all individual bond polarities and lone-pair contribution in the molecule. Polarity has a dramatic effect on the physical properties of molecules, particularly on melting points, boiling points, and solubility...