vsper, molecular orbitals, and organic molecules
TRANSCRIPT
LS50 2015
Lecture 10: VSPER, molecular orbitals, and organic molecules
Cassandra Extavour Dpts. OEB & MCB
Recap: what happened in Lecture 09
You learned how to: • Describe and explain some basic properJes of an element
based on periodic trends • Figure out the number of valence electrons for an atom • Understand electronegaJvity; use this to predict formaJon of
compounds in general terms • State the octet rule
– name and define the types of bonds that atoms can engage in to saJsfy it – recognize the three major types of violaJons of this rule
• Draw Lewis dot structures of atoms, ionic compounds, and covalent compounds
Learning goals for today
By the end of this lecture, you should be able to: • Use VSPER to predict molecular geometries • Understand and apply basic chemical nomenclature to name chemical compounds • Use molecular orbitals to explain observed bond angles and magneJc properJes of
compounds • Draw organic structures based on chemical formula • Understand different schemaJc representaJons 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
• SomeJmes you realize that there is more than one possible correct Lewis dot structure – Resonance: the same atoms are linked to each other but in a different
bonding paVern
– Structural isomers: different atoms are linked to each other
• e.g. O3, CO32-‐
Resonance – one example
O O O
O O O
O O OBoth structures are correct!
Ozone O3
• The bonding is an “average” of these two structures: electrons are “delocalized” between the atoms
• So bond strength is (2+1)/2 = 1.5 rather than an integer
Resonance – another example
What is the bond strength?
Carbonate* CO32-‐
* this is an example of a polyatomic ion – we’ll talk about those in a couple of slides
Compound Nomenclature • Ionic compounds
– First part of the name = name of the metal element – Second part of the name – name of the non-‐metal element with “ide” suffix – For transiJon metals, you have to indicate the charge on the metal ion since they can have
more than one charge: use roman numerals in parentheses afer the name of the ion to do this
– Examples: what are the names of these compounds? • Al2O3 • FeCl2 • FeCl3
• Covalent compounds – Nonmetal atoms can combine in >1 set of atomic raJos – To prevent ambiguity, use greek prefixes to indicate the number of atoms of each element in
the compound • mono, di, tria, tetra, penta, hexa, hepta, octa • if there is no prefix, mono is assumed
– Examples: what are the names of these compounds? • NO • NO2 • N2S3 • PS5
Compound Nomenclature Polyatomic ions: • Two or more nonmetal ions covalently bonded with an overall charge • a metal ion or other atom could bond ionically with it • 3 rules for naming these: 1. if you add H, add “hydrogen” to the beginning of the name (you have to know
the name of the covalent ion first: see Table in Lecture 10 road map) – e.g. CO3
2-‐ carbonate, HCO3-‐ hydrogen carbonate
2. if you remove an O, change the end of the name of the ion from “ate” to “ite” – e.g. NO3
-‐ nitrate, NO2-‐ nitrite (note the charge is unchanged)
– if you lose a further O call it “hypo_ite” – 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 – e.g. SO4
2-‐ sulfate, SeO42-‐ selenate (selenium), TeO4
2-‐ tellurate (tellurium) – there are a number of polyatomic ions that you can learn the names of to help you
apply the above rules – see Table in Lecture 10 road map
Valence Shell Electron Pair Repulsion (VSEPR)
• Used to give you informaJon about the shape of covalent compounds (the geometry around a central atom)
• Basic premise: electrons repel each other, and arrange themselves in space so as to minimize the repulsion
• A region of space around a central atom that has at least one electron pair is a domain = a concentraJon of electron charge density in space
• The number of domains determines the geometry in roughly predictable ways • The number of domains, not the number of pairs (bond order) is what is important • For resonance structures, you can use any one of the resonance forms to idenJfy
the number of domains and predict the gross shape
• VSEPR caveats: – Ligands that are poorly electronegaJve may not have the expected shape – It mostly ignores the repulsive effects of ligands – It is not good at predicJng shapes of transiJonal metal complexes – Doesn’t give us informaJon about magneJc properJes of molecule – à that’s what molecular orbitals are for (coming up)
VSEPR Examples: 2 domains Basic Geometries for 2, 3, and 4 Electron Pairs
BeCl2 Cl Be Cl
Two pairs minimize repulsions if oriented 180o from oneanother, giving a linear geometry.
BCl3 B
Cl
Cl Cl
Three pairs minimize repulsions if oriented 120o from oneanother, giving a trigonal planar geometry.
CH4 C
H
HH
H
Four pairs minimize repulsions if oriented 109.5o from oneanother, giving a tetrahedral geometry.
• Two electron pairs (two domains) minimize repulsion if they are oriented 180º away from each other
• à Linear geometry
Double and Triple Bond Domains
L Shape is determined by the number of electron domains(regions), not simply the number of electron pairs.
CH H
O
4 pairs in 3 regionsY trigonal planar
O C O
4 pairs in 2 regionsY linear
H C N4 pairs in 2 regions
Y linear
• Four electron pairs in two domains • à Linear geometry
Double and Triple Bond Domains
L Shape is determined by the number of electron domains(regions), not simply the number of electron pairs.
CH H
O
4 pairs in 3 regionsY trigonal planar
O C O
4 pairs in 2 regionsY linear
H C N4 pairs in 2 regions
Y linear
• Four electron pairs in two domains • à Linear geometry
Check: what are the names of these compounds?
CO2
CHN
BeCl2
VSEPR Examples: 3 and 4 domains
Basic Geometries for 2, 3, and 4 Electron Pairs
BeCl2 Cl Be Cl
Two pairs minimize repulsions if oriented 180o from oneanother, giving a linear geometry.
BCl3 B
Cl
Cl Cl
Three pairs minimize repulsions if oriented 120o from oneanother, giving a trigonal planar geometry.
CH4 C
H
HH
H
Four pairs minimize repulsions if oriented 109.5o from oneanother, giving a tetrahedral geometry.
Check: what are the names of these compounds?
VSEPR geometry representaJons
What are those green knobs?
VSEPR Examples: central atom with some non-‐bonded pairs
• Molecules with non-‐bonding pairs on a central atom sJll have shapes based on arrangement of electron domains
• However they may have bond angles altered (compared to what is shown on the previous slide) by some other factors: – repulsions between lone pairs and bond pairs: lone pairs are
“stronger” (have larger domains) than bond pairs – More electronegaJve ligands are “weaker” (have smaller domains)
than less electronegaJve ones – Double and triple bonds are “stronger” (have larger domains) than
single bonds
• e.g. H2O
Note: this is why you needed to know how to draw Lewis dot structures! It was good for something aHer all!
VSEPR Examples: central atom with some non-‐bonded pairs
Molecular orbitals (MO): fun for the whole family (of atoms in a molecule)
• Atoms that combine into molecules share electrons to some extent à the electrons “belong” to the molecule now, not the individual atoms
• For this reason, Lewis dot structures alone can’t explain the magneJc properJes or shapes of molecules with covalent bonds à MO can help us determine molecular structure (nice!)
• Central idea: the electron distribuIon for an atom in a molecule is different from the electron distribuIon it has in isolaIon
• Step 1: Combine atomic orbitals (AO). – remember s, p, d etc.? Now you’ll be glad you learned them! – The MO outcome depends on which AOs you are combining – Since orbitals are wave funcJons, you have to consider the outcome of combining
AO in phase and out of phase (the MO has its own Ψ) – The number of MO has to equal the total number of AO
• Step 2: Determine the relaJve energy states of the MO. • Step 3: Fill the MO with all of the valence electrons brought into the
structure by its consJtuent atoms. – follow the Auqau and Pauli principles just like for AO
Since AO are wave funcJons (Ψ), consider the outcome of combining AO in phase and out of phase to create the Ψ of the MO
Combining two s orbitals: H2
Sigma (σ) bonds • Cylindrically symmetrical about the plane of the internuclear axis • permit rotaJon • used in single bonds and only one of the electron pairs in bonds with order >1 • Can be created by mixing two s AO, two p AO, or combining a hybrid AO with s or p AO
1. Combine AO to make MO (# AO has to = # MO)
2. Determine relaJve energy states 3. Fill them with electrons (Auqau, Pauli)
1s σ*1s AnJ-‐bonding 1s
σ1s bonding
MO theory can help understand why some molecules don’t exist: He2
Bond order • Recall: this is the # bonds between a pair of atoms • Calculate this from MO as follows: (# bonding electrons -‐ # anJbonding electrons)/2 • For He2 we see that bond order = 0! • This means they won’t bond: not enough energeJc advantage for this to happen • In other words, they are not more stable together than they are apart
He-He
Combining s and p orbitals: CH4 (methane)
HybridizaHon of AO • Combine the AO within an atom to get MO for that atom within a parJcular molecule • The number and types of bonds shown in the Lewis diagram dictate what classes of MO
will form • Different hybrid MO have different specific geometries (e.g. sp3: tetrahedron) • Hybrid MO only form σ bonds
sp3 Hybrid orbitals
Combining s and p orbitals: C2H4 (ethylene)
Pi (π) bonds • Above or below the plane of the internuclear axis • Do not permit rotaJon • Used in bonds with order >1 • Can be created by mixing two p AO
Elements of life
“I cannot, so to say, hold my chemical water and must tell you that I can make urea without thereby needing to have kidneys, or anyhow, an animal, be it human or dog.”
– Friedrich Wöhler 1828
• 90% of our bodies are made of C, H, O, N – about 70% H2O
• Across all life, most molecules contain principally 6 elements – H C, N, O, P, S
51
=
CH2
CH3 CH3
C
H
R C H
H
C
C
C
C
C
C H
H
H
H H
H
H
H
H
H H
H=
where it is now recognised that there are CH3 groups at the points where the lines stop.
This comes in handy when drawing large molecules such as C10H22. Instead of laboriously drawing the
molecule out we can draw –
Question 5.6 What are the molecular formulae of the following compounds?
(i)
(ii)
(iii)
One last thing on shorthand is the use of “R”. When chemists write a molecule as
the “R” simply signifies an organic group which they are not particularly interested in! That is to say that the
“R” group will remain unchanged during a chemical reaction.
Rings Alkanes, alkenes and alkynes are all linear molecules, where the carbon atoms are in a chain. There is however
another way of joining carbons together and that is in rings. Consider for example the following molecule.
Here we have six carbon atoms joined in a ring. The naming of these compounds is still based upon the
number of carbon atoms so this molecule is called cyclohexane. You will notice that the molecular formula for
cyclohexane is C6H12. For all cycloalkanes the carbon:hydrogen ratio is Cn:H2n. This is exactly the same ratio as
for alkenes. However, cycloalkanes are saturated molecules and they will have a very different chemistry from
alkenes with the same molecular formula, so beware!
How to draw organic structures • Organic compounds contain so many C and H atoms that we
develop highly simplified ways of represenJng the structures • We were already starJng to do this with Lewis dot structures:
replacing the two dots of an electron pair with a single line • Organic compound shorthand conJnues this trend: you’ll see any
combinaJon of the following Jme-‐saving methods used: – Don’t show any dots: show bonding pairs as lines; lone pairs are
assumed to be there – Don’t show the atomic symbol for H bonded to C – Don’t show C (“skeletal”) – Don’t draw out funcJonal groups; instead
• Write chemical formula • Write “R”
Remember electronegaJvity?
Differences in electronegaIvity can have an impact on the distribuIon of shared electrons.
Polar (covalent) bonds I
• In organic chemistry we’re mainly concerned with covalent bonds • However, not all covalent bonds are created equal: electrons may
not be shared equally between atoms
• If the electronegaJvity difference between two atoms in a bond is great, then the more electronegaJve atom “has more” of the electrons
• The δ+ and δ-‐ symbols indicate the polarity of the bond
7
ion-dipole intermolecular forces: ion (polar) ↔ polar
H O H
+ -
+
Na+
Polar (covalent) bonds II
• The polarity of the bond is quanJfied by a vector quanJty called a dipole moment µ – Describes the overall distribuJon of electrons in the bond – Is a product of charge magnitude (scalar q) and bond length (vector r) – In units of debyes (D) – µ = 0 for a nonpolar bond
• Polar molecules have permanent µ > 0
• In molecules with >1 polar bond, the net µ for the molecule is the vector sum of all µ for each bond
µ > 0
Diversify your bonds
Covalent Increasing Ionic Character
Polar Covalent
Ionic
Hydrogen bonds
• Dipole-‐dipole interacJons are those between permanent dipoles of two polar molecules – Only found in molecules with net dipole µ > 0
• Hydrogen bonds are a specific type of dipole-‐dipole interacJon
• Happens when an H (usually bonded to an O, N or F) is electrostaJcally aVracted to a lone pair on another molecule
19
• Hydrogen bonding typically occurs when a hydrogen atom bonded to O, N, or F, is electrostatically attracted to a lone pair of electrons on an O, N, or F atom in another molecule.
Hydrogen Bonding
Recap: learning goals for today Hopefully, you now feel able to: • Use VSPER to predict molecular geometries • Understand and apply basic chemical nomenclature to name chemical compounds • Use molecular orbitals to explain observed bond angles and magneJc properJes of
compounds • Draw organic structures based on chemical formula • Understand different schemaJc representaJons 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
If not, please ask your quesIons during secIon, on Piazza, or come to office hours! (Thursdays at noon by appointment)