Download VSPER, Molecular Orbitals, and Organic Molecules

LS50 2015
INTRO TO CHEMISTRY WEEK: SEPTEMBER 14 – 18, 2015
Learning goals
By the end of this week, you should understand the following:
• The key findings and interpretations of some critical experiments that led to our current
understanding of atomic structure and electron configurations (Lecture 08)
• How electron configurations are related to chemical bonds and interactions (Lecture 09)
• How chemical bonds are related to assembly and structure of biological molecules (Lectures 0910)
• How molecular assembly and structure are related to molecular function (Lectures 11-12)
• How knowing molecular function facilitates prediction of the evolution of biological molecules
(Lecture 12)
Lecture 10 – VSPER, Molecular orbitals, Organic molecules and chirality
Learning goals
By the end of this lecture, you should be able to:
• Understand and apply basic chemical nomenclature to name chemical compounds
• Use VSPER to predict molecular geometries
• Use molecular orbitals to explain observed bond angles and magnetic properties of compounds
• Draw organic structures based on chemical formula
• Understand different schematic representations of organic compounds
• Determine polarity of bonds as well as molecules
• Define dipole moments and hydrogen bonds
• Name the six most abundant elements found in biological molecules
•
Resonance (one more concept related to Lewis dot diagrams from Lecture 09)
• Sometimes you realize that there is more than one possible correct Lewis dot structure
• There are two types of these:
o Structural isomers: different atoms are linked to each other
§ In this case, the “real” (most likely) structure is the one with the lowest formal
charge
§ The formal charge is calculated for each atom in the molecule by subtracting the
atomic number from the number of electrons present in the dot structure being
evaluated (counting a shared electron pair as one electron)
§ e.g. Cl2O
o Resonance: the same atoms are linked to each other but in a different bonding pattern
§ In this case, the “real” structure is an average of the resonance structures, which
can result in fractional bond orders
2o e.g. O3, CO3
•
Chemical nomenclature
• Ionic compounds
o First part of the name = name of the metal element
o Second part of the name – name of the non-metal element with “ide” suffix
o For transition metals, you have to indicate the charge on the metal ion since they can
have more than one charge: use roman numerals in parentheses after the name of the
ion to do this
• Polyatomic Ions
o This is two or more nonmetal ions covalently bonded with an overall charge: a metal ion
or other atom could bond ionically with this
o 3 rules for naming these:
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•
1. if you add H, add “hydrogen” to the beginning of the name (you have to know the
name of the covalent ion first)
2a. e.g. CO3 carbonate, HCO3 hydrogen carbonate
2. if you remove an O, change the end of the name of the ion from “ate” to “ite”
a. e.g. NO3 nitrate, NO2 nitrite (note the charge is unchanged)
b. if you lose a further O call it “hypo_ite”
c. if you add an extra O, call it “per_ate”
3. if you replace the central atom in the ion with another atom from the same group,
just replace the corresponding part of the name
222a. e.g. SO4 sulfate, SeO4 selenite (selenium), TeO4 (tellurium)
o there are a number of polyatomic ions that you can learn the names of to help you apply
the above rules (see Table 1)
Covalent compounds
o Nonmetal atoms can combine in >1 set of atomic ratios
o To prevent ambiguity, use greek prefixes to indicate the number of atoms of each
element in the compound
o (if there is no prefix, mono is assumed)
Chemical formulas
o Hill system: has three rules:
1. Write C, then H, then everything else in alphabetical order
2. If ionic compound, write positive ion then negative ion
3. If oxide, acid or hydroxide,
a. Acid: start name with H
b. Hydroxide: end with OH
c. Oxide: end with multiples of O
•
Valence shell electron pair repulsion (VSEPR)
o Also called electron domain theory
o Used to give you information about the shape of covalent compounds (the geometry
around a central atom)
o The basic premise is that electrons repel each other, and will arrange themselves in
space so as to minimize the repulsion
o A region of space around a central atom that has at least one electron pair is a domain =
a concentration of electron charge density in space
o The number of domains determines the geometry in roughly predictable ways
§ e.g. central atom with only bonded pairs: BeCl2, BCl3, PCl5, SF6, H2CO, CO2,
HCN
§ e.g. central atom has some non-bonded pairs: SnCl2, NH3, H2O
o VSEPR caveats:
§ Ligands that are poorly electronegative may not have the expected shape
§ It mostly ignores the repulsive effects of ligands
§ It is not good at predicting shapes of transitional metal complexes
•
Molecular orbitals (MOs)
o These result from the overlap of atomic orbitals (AO), but are shared by the entire
molecule (a group of bonded atoms)
o Provides a way to think about molecular structure by considering electrons as moving
under the influence of all nuclei of the molecule
o We need this to explain the magnetic properties of molecules
o The basic way to build these is to know
§ the AOs containing the valence electrons involved in the bond
§ the total number of electrons involved in the bond
o once you know that information
§ combine the AOs that interact at each bond (the nature of the MO formed will
depend on the specific combination of AOs – see below)
§ determine the relative energy states of the new MOs that have been created
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o
o
o
o
o
o
§ distribute all of the electrons involved in the bond among the MOs
the total number of MOs formed should equal the total number of AOs involved
when AOs combine to make an MO, there are two types of MO that can be formed:
i.
bonding orbital
• electrons here have a higher probability of being between nuclei than
elsewhere
• tends to hold nuclei together
• called constructive interference: has a lower energy than the states of
the isolated atoms
i.
anti-bonding orbital (indicated with a superscript asterisk)
• electrons tend to spend more of their time not between the nuclei
• tends to weaken the bond
• called destructive interference: has a higher energy than the states of the
isolated atoms
you can also have a non-boding orbital, which has no positive or negative interaction
between its orbitals
these do not contribute to nor detract from the bond strength
depending on the angle at which AOs combine relative to the axis of the internuclear
bond axis, they are called σ or π MOs
it’s important to know which one is which, because they have different structural
consequences: σ bonds permit rotation around the internuclear axis, but π bonds do not
Elements of Life
• Most biological molecules are made of just six atoms: C, H, N, O, P and S
Drawing organic structures
• Many different 2D schematics are used to show what chemical structures look like
• They each have strengths and weaknesses with respect to what they are able to show and what
the need the viewer to assume in order to understand their meaning
• For example, Lewis dot diagrams are very clear at showing which electrons are principally
associated with which atoms
• However, as we now understand with MO theory, it is closer to reality to think of electrons being
associated “globally” with an entire molecule, rather than “belonging” to just one or another atom
(or being equally shared – see dipole moments in Lecture 11 – between multiple atoms)
• One trend in drawing organic structures is to simplify them more and more by using shorthand or
eliminating entirely representations of carbons, hydrogens, or lone atoms
Polarity in covalent bonds
• When an electronegativity difference exists between atoms in a covalent bond, the electron
density of the shared electron pair may not be evenly distributed between the two atoms
• In this case, the bond is polar, and is said to have a dipole moment µ
• µ is a vector quantity dependent on the charge magnitude and bond length/polarity direction
Hydrogen bonds
• hydrogen bonding is an example of an intermolecular force that results from dipole-dipole
interactions
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Table 1. Eight common classes of polyatomic ions, and related polyatomic ions
1. Carbonate CO32Hydrogencarbonate HCO32. Nitrate NO3Nitrite NO23. Hydroxide OH4. Acetate C2H3O25. Ammonium NH4+
6. Sulfate SO42Hydrogensulfate HSO4Sulfite SO32Hydrogensulfite HSO3Selenate SeO42Hydrogenselenate HSeO4Selenite SeO32Hydrogenselenite HSeO37. Phosphate PO43Hydrogenphosphate HPO42Dihydrogenphosphate H2PO4Phosphite PO33Hydrogenphosphite HPO32Dihydrogenphosphite H2PO3Arsenate AsO43Hydrogenarsenate HAsO42Dihydrogenarsenate H2AsO4Arsenite AsO33Hydrogenarsenite HAsO32Dihydrogenarsenite H2AsO38. Chlorate ClO3Chlorite ClO2Hypochlorite ClO
Perchlorate ClO4Bromate BrO3Bromite BrO2Hypobromite BrO
Perbromate BrO4Iodate IO3Iodite IO2Hypoiodite IO
Periodate IO4-
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Hybridisation – summary
Table 2. Geometries of hybrid molecular orbitals
Hybrid Atomic orbitals
isation that are mixed
Geometry
General
formula
Examples
sp
linear
AB2
BeH2
trigonal planar
AB3
BF3, CO32C2H4
sp2
s+p
s + px + py
sp3
s + px + py + pz
tetrahedral
sp3d
s + px + py + pz + dz2
Trigonal Bipyramidal AB5
s + px + py + pz + dx2-y2
square pyramidal
sp3d2
s + px + py + pz + dz2 + dx2-y2 octahedral
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AB4
AB6
SO42-, CH4,
NH3, H2O,
PCl5, SF4
SF6
[Ni(CN)4]2[PtCl4]2-
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