Topic 5b: Covalent bonds

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Bonding is a spectrum, and where on the spectrum a molecule falls is determined by the difference in electronegativity between its constituent atoms.

The opposite ends of the chemical bonding spectrum: ionic and covalent

ionic bonds: electrostatic forces between cations and anions

Covalent bonds: the sharing of electron density between two atoms

Nonpolar covalent bond: the electron density is distributed equally among both atoms

polar covalent bond: as a result of differing electronegativities, the electron density is unequally distributed between the atoms. These bonds possess partial ionic character due to bond polarity

Molecular polarity: occurs when an entire molecule, which can be made of several covalent bonds, has a net polarity.

This means one end has a negative charge and the other end has an overall positive charge.

Thus, polar molecules can act as electric dipoles, which can interact with electric fields surrounding the molecule.

Dipole moment: gives the strength of the dipole’s interaction with an electric field; Calculated by multiplying the magnitude of separated charge and the bond length.

Bonding is a spectrum, upon which there are bonds that are ionic with covalent character. A common example of this phenomenon are carbon-metal bonds, such as methyllithium (H3CLi), where it can be thought of as a metal anion and the lithium counterion, however this is not exactly accurate or a purely ionic bond because there is shared electron density between the atoms.

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Lewis dot diagrams, which display the valence electrons that each atom has, can be useful in determining how many bonds an atom is able to form.

The formal charge of the atom or molecule should be written in the Lewis dot diagram.

In order to calculate the formal charge, count all of the lone pair electrons and half of each pair present in the diagram, and then compare this total to the number of valence electrons that a neutral atom would normally carry.

There are resonance forms, some of which are more prevalent or contribute more significantly to the structure of the molecule. The rules of resonance forms are as follows:

each atom should have an octet

structures should have minimal charge separation (there should not be multiple charges present within the structure)

the atoms themselves should not move in resonance structures

overall charge should remain the same, even if the charges may change position from atom to another.

any charge separation that is within the structure should be guided by electronegativity trends

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σ bonds: The strongest type of covalent bond that can form when 2 orbitals (S, P, D. and/or F) directly overlap between the nuclei of two atoms.

π bonds: occur when there's an overlap between unhybridized P orbitals of two adjacent atoms. The overlap does not occur between the nuclei of the atoms.

Determining whether bonds are Sigma or Pi:

single bonds are assumed to be Sigma

in double bonds, 1 bond is Sigma and one is π.

In triple bonds, one bond is Sigma and the other two are Pi.

Double and triple bonds result in greater rigidity of molecules, shorter bond lengths between atoms, and higher bond energies than single bonds.

Conceptualization of atomic geometry:

hybridization of atomic orbitals: the mixing of atomic orbitals to form new hybrid orbitals.

These orbitals have a specific geometrical arrangement in space that matches the observed bonding geometry of molecules.

Denoted as sp#, where the # is how many p orbitals are required to explain the observed bonding.

Molecular shape: one of the key factors determining the functionality and reactivity of a given molecul

How to determine hybridization of an atom: https://www.masterorganicchemistry.com/2018/01/16/a-hybridization-shortcut/#a1

VSEPR theory (valence shell electron pair repulsion): states that atoms will arrange themselves in such a way that their electron clouds experience the least amount of repulsion from each other

Describes 5 shapes of simple molecules:

Linear: typically seen in AX2 triatomic molecules that have two bonding orbitals that are 180 degrees apart. Examples: BeCl2, CO2. Have sp hybridization.

Trigonal planar: typically seen in AX3 molecules, have bonding angles of 120 degrees. Have sp2 hybridization.

tetrahedral: typically seen in AX4 molecules, where a central atom is bonded to four other atoms. The angle between any two bonds will be 109.5 degrees. Have sp3 hybridization.

Trigonal bipyramidal: type of geometry where the five atoms bonded to the central atom are not all equivalent. In this case two different positions are defined—axial and equatorial. The axial position is more crowded and atoms in an axial position will have a 90 degree bond angle between one another. Atoms in an equatorial position will have a 120 degree bond angle between one another. Have sp3d hybridization.

octahedral: typically seen in AX6 molecule, all the atoms are geometrically equivalent with bond angles of 90 degrees. Have sp3d2 hybridization.

Deviation from ideal VSEPR bond angles is possible due to electron repulsion.

​​​​​​​bond angle strain: deviation from ideal bond angles that with significant deformity has a destabilizing effect

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Elements in the same group of the periodic table are similar to one another and tend to have similar structural formulas.

Condensed formulas rely on the following rules:

the formula is read from left to right

each carbon atom in the formula is connected to the next carbon atom in the formula

everything to the right of a carbon is connected to that carbon if it is before the next carbon

parentheses indicate whole groups attached to a carbon

normal valencies must be satisfied (carbon makes maximum of 4 bonds, nitrogen 3, oxygen 2, hydrogen one, halogens 1)

To convert the condensed formula into structural formula, first lay out the carbon backbone. Then you would add on the substituents appropriately.

line notation is even simpler than condensed formula. The rules for line notation are:

every unlabeled vertex or terminus is a carbon atom

any unused valency of carbon is filled by hydrogen

it is important to include the lone pair or any charges on atoms in line notation, as these can affect the chemistry of the species.

For instance, lone pairs are integral to hydrogen bonding.

This is why most H-bond donors are polar; The main atoms with which hydrogen bonds are nitrogen, fluorine, oxygen, and sulfur.

Because these atoms draw most of the electron density towards themselves, hydrogen gains a partial positive charge which then allows it to hydrogen bond with atoms containing a lone pair.

Aside from water, the other common molecule in which hydrogen bonding is integral is DNA; The nitrogenous bases are held together with H bonds.

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Resonance: describes the delocalization of electrons within molecules.

electrons do not have fixed positions but have probabilities of being found in certain spaces (orbitals). Resonance forms illustrate areas of higher electron densities.

Resonance structures compare and contrast two or more possible Lewis structures that can represent a particular molecule.

molecules with multiple resonance structures are more stable than those with fewer.

Some resonance structures are more favorable than others

Resonance hybrid: the net sum of valid resonance structures, which represents the overall delocalization of electrons within the molecule

Rules for writing resonance forms:

all of the structures for the same molecule or atom should have the same number of electrons

all structures must follow the rules of writing Lewis structures

the hybridization of the structure must remain the same

the skeleton of the structure cannot be changed; Only the electrons can move

all structures must have the same amount of lone pairs

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Isomers: molecules with the same molecular formula but different arrangement of the atoms in space.

Structural (constitutional) isomer: two or more organic compounds have the same molecular formulas but different structures.

Conformational isomers: occur when there's rotation about a Sigma bond that results in difference in conformation

Stereoisomers: differ in the spatial arrangement of atoms, rather than the order of atomic connectivity.

Geometric (cis-trans) stereoisomers:

cis isomer: 2 substituents that are the same are on the same side of the molecule

trans isomer: 2 substituents that are the same are on opposite ends of the molecule

Optical stereoisomers:

enantiomers: nonsuperimposable mirror images. Enantiomers are identical in almost all physical properties, except in the way that they rotate polarized light (optical rotation).

Enantiomers that rotate light clockwise (to the right) are dextrorotary isomers, while enantiomers that rotate light counterclockwise (to the left) are levorotary isomers.

You cannot tell just by looking at a molecule its optical activity.

If a sample is optically inactive, it either does not contain chiral molecules, or it is a racemic mixture (has 50/50 composition of D-isomers and L-isomers).

Enantiomeric excess: a quantitative measure that describes scalemic (enantiomerically enriched) mixtures. To calculate this, you would subtract the percentage difference between the two components (ex: 55/45 mixture has a 10% ee).

Optically pure samples are always optically active, however not all optically active samples are optically pure (they could be a scalemic mixture).

Diastereomers: stereoisomers which are not mirror images of each other. They are achiral

Chiral: objects that are non-superimposable mirror images of each other; Do not have a plane of symmetry.

The molecular basis for chirality is the presence of one or more chiral centers in chiral molecules.

A molecule cannot be chiral if any two or more substituent groups of the molecule are the same.

all chiral molecules have a characteristic specific rotation

Comparing stereoisomers:

enantiomers are molecules where EACH and EVERY chiral center is inverted

diastereomers are molecules where AT LEAST ONE but NOT ALL chiral centers are inverted

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R and S nomenclature:

R is right, S is left in Latin

There is NO correlation between R/S nomenclature and optical activity (D/L)

The lowest priority group should be facing away from the page

Priority is assigned based on mass.

If order of priority is clockwise, then R. If order of priority is counterclockwise, then S

E and Z nomenclature:

E is opposite and Z is together

Used when the groups on either side of a double bond are all different from one another

Priority is assigned based on atomic mass

If highest priority groups are opposite, then E. If highest priority groups are on same side of axis, then Z

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Carbons with only single bonds tend to be tetrahedral (sp3), carbons that are part of a double bond tend to be trigonal (sp2), and carbons with triple ones are always diagonal (sp).

The addition of differing substituents decreases symmetry. If a tetrahedral center is surrounded by four different groups, then the center is asymmetric.

Exceptions:

carbanions: carbons with a negative charge, have lone pairs on the carbon giving them the negative charge. You would predict sp3 geometry about a methyl anion, and most carbon ions do exhibit a pyramidal shape.

Carbocations: carbon centers with only three bonds and a positive formal charge, example methyl cation. These exhibit planar, sp2 geometry because there are only three substituents to account for.

Radicals: have P orbitals that are only half filled (have only 1 electron, instead of 2 like a conventional lone pair), therefore requiring less space than a completely filled orbital, but more than an empty one; typically are depicted with planar geometry.

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Depictions of molecular structure:

Sawhorse (dash-wedge) projection: Substituents coming out of the plane towards the observer are depicted with wedges; Substituents going away from the observer are shown with dashes. Substituents drawn with a plain line that is neither a dash nor wedge, are assumed to lie in the plane of the paper.

staggered conformation: the plain bonds have a “Z” formation (see below)

eclipsed conformation: the plain bonds have a "U” formation (see below)

when the molecule is flipped, the wedges and dashes become the opposite (if there is a dash, it becomes a wedge and vice versa). This does not happen if the molecule is simply turned or rotated in-plane.

Newman projection: shows a carbon- carbon bond come on where the small dot represents the front carbon and the large circle behind it represents the back carbon.

in a staggered conformation, there are 2 possible substituent relationships:

Gauche: describes substituents that are side by side

anti-periplanar: describe substituents that are as far away as possible from one another

in an eclipsed conformation, there is 1 possible relationship between substituents --> eclipsed

Conformational analysis: Analysis of the relative energies of different molecular conformations and substituent relationships within those conformations

Eclipsed conformations are higher energy than staggered conformations, as a result of steric interactions from the close proximity of substituents.

Large groups in a gauche arrangement in staggered conformations have higher energies then those same large groups being antiperiplanar to each other.

H- H bonds can be ignored in this analysis because the atoms are relatively small.

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Cycloalkanes are less flexible than linear alkanes, and therefore have a limited number of conformations. The most stable conformation of a cycloalkane is the chair conformation.

Substituents can either be axial or equatorial in a chair conformation; with a chair flip, the substituents take on the opposite attitude.

Stable conformations arise from the minimization of the ring strain:

Bond angle strain: arises from compression of sp3 bond angle to accommodate a cyclic arrangement

Torsional strain: comes from torque on individual bonds with eclipsed substituents

Ring strain decreases as the # of carbons in the ring increases

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Structural representations can give varying levels of information. Constitutional is the lowest level of detail (how the atoms are connected to one another), then stereochemical, and then conformational (the shape a molecule adopts, like tertiary structure of a protein).

when converting from a Newman projection to sawhorse, imagine the molecule as a door that you are shutting. So in Newman, you are looking straight down a C-C bond; imagine your push this bond into the page, such that you are closing a door and it becomes a sawhorse projection, and you are now seeing that same bond face-to-face, horizontally instead of sighting down its length.

This technique can also be used for converting from Fischer to sawhorse projections. Imagine you are closing the door from the substituents on the L hand side and pushing and pushing until the door closes. This is essentially a flip of the molecule, which means we would change from the wedges to dashes and vice versa. It is important to note the wedges or dashes on the Fischer so that they are converted correctly into the sawhorse.

Conversion of chair conformations into dash-wedge formations:

foldaway ironing board: imagine the back C-C bond is a hinge. Push upwards on the bond with the substituents, so that you imagine you are closing the legs of an ironing board. This results in the molecule rotating up so that the substituent containing bond is now at the top (see below)

drop-leaf table: imagine there is a bond which acts as a hinge. You push on the bond with the substituents downwards, imagining a latch opening, releasing the flat part downwards, resulting in a rotation of the molecule so that the substituents are now on the bottom instead of the top of the molecule. imagine there is a bond which acts as a hinge. You push on the bond with the substituents downwards, imagining a latch opening, releasing the flat part downwards, resulting in a rotation of the molecule so that the substituents are now on the bottom instead of the top of the molecule.