Gen Chem - Molecules

What occurs during ionic bonding?

One atom gives one (or more) electron(s) to another creating positive and negative ions that are attracted to each other

What occurs during covalent bonding?

Atoms share electrons

What is the organisation of electrons in the covalent bond in H2?

• Both atoms ‘share’ their electron

• Shared electrons (-ve charge) exist mostly between the two hydrogen nuclei (+ve charge)

• Both nuclei are attracted to this which forms the bond

• Mutual attraction of nuclei to electron density located between them

• Bonded is a lower energy arrangement than two separate individual atoms so bond forming is a favourable process

• Mutual attraction to the shared electron density (opposite charges) causes nuclei to get closer but nuclei will repel each other too (like charges)

• Bond length is the point where the forces balance

How is atomic distances related to energies?

• Two atoms that are too far apart to interact are considered to be at zero energy

• Energy becomes negative as a bond is formed

• If atoms are pushed closer together, repulsion dominates and energy rises sharply

How can bonds be overcome chemically?

• Can separate a bond with energy (heat, radiation or a chemical reaction)

• All require energy to be put in to overcome the bond dissociation energy

What is the bond dissociation energy?

Difference in energy between to non-interacting atoms and those two atoms forming a bond

What denotes a strong bond?

A large bond dissociation energy

How is bond length determined?

• All bonds vibrate like springs

• Depends on temperature

• Bond length is defined as the mean distance between the atoms

How do atoms form a covalent bond in the Lewis model?

• Electrons are always shared in pairs

• These count towards both atoms’ outer-shell electron count

• Unshared electrons are ‘lone pairs’

How are bond orders determined in the Lewis Model?

• Bond order = the count of the bonds

• Single bond = bond order 1

• Double bond = bond order 2

• Triple bond = bond order 3

What are the deficiencies of the Lewis model?

• Nitrogen monoxide is a stable molecule

• O has 6 valence electrons and N has 5

• No way to share that gives both an outer shell

• PF5 is a naturally synthesised molecule

• P has 5 valence electrons

• P will have 10 valence electrons in PF5

• Causes hypervalency

• Lewis predicted all 12 valence electrons in O2 exist as bonding pairs and lone pairs

• Experiments show O2 is paramagnetic which only occurs with unpaired electrons

How are polar bonds formed?

• When electrons are unevenly shared within a covalent bond

• Polarity of bonds is often important in organic mechanisms

How does electronegativity affect bonds?

• In H2 and other homonuclear diatomics the electron density of the two atoms is symmetrical

• If the atoms were different one atom of the pair would likely attract more of the electron density of the bond than the other

• At an extreme would cause all bonding electrons to be on one atom and there to be an ionic bond rather than covalent

What is the Pauling Electronegativity Scale used for?

• Describes how much atoms attract electron charge to themselves

• Higher the number the greater the attraction

• Electronegativity can be used to predict the polarity of bonds

• F is an electronegative element

• Cs is an electropositive element

What are the contrasting theories for bonding?

• Valence bond theory

• Molecular orbital theory

What are the two approximations of molecular orbital theory?

• Orbital approximation

• Linear combination of atomic orbitals

What does the orbital approximation involve?

• Overall electronic state of a molecule is a product of one wavefunction for each electron

• Single-electron wavefunctions are molecular orbitals and can span the entire molecule, not just individual atoms

What does the linear combination of atomic orbitals involve?

• Simplifies the challenge of constructing molecular orbitals

• Molecular orbitals can be constructed by adding and subtracting atomic orbitals

How are molecular orbitals formed in H2?

• Single electron occupies a 1s orbital

• Adding the wavefunction (in-phase combination) produces a molecular orbital with a lot of electron density between the atoms - spans the whole molecule

• Wavefunctions for atomic orbitals for 1s show the location of the nuclei

• Molecular orbital is a in-phase combination - increased probability of finding electrons between nuclei and considered to be a bonding orbital as it holds the atoms together

• Subtracting these wavefunctions (out-of-phase combination) produces a molecular orbital with one ‘positive’ part and one ‘negative’ part

• Out of phase combination has zero electrons between the nuclei - probability form a node between the nuclei

• Actively pulling the molecule apart - anti-bonding orbital

How do you determine the number of molecular orbitals possible in one molecule?

The number of possible molecular orbitals is always equal to the number of atomic orbitals from which they are formed

What does the molecular energy diagram for H2 look like?

• Energy of a bonding orbital is lower than the energy level of the independent atomic orbitals

• Agrees with H + H → H2

• Energy of the antibonding orbital is higher

• Each molecular orbital has capacity for two electrons and lower energy orbitals are filled first

• Both electrons in H2 occupy a bonding orbital which holds the molecule together

What does the molecular orbital diagram look like for He2 and what does this mean for the stability of the molecule?

• Has two extra electrons than H2

• Cannot go into the same bonding orbital

• Extra electrons go into the antibonding orbital

• Causes the same amount of bonding and antibonding orbitals so no overall bonding so He2 does not exist

How is the bond order determined in MO theory?

• Can account for both bonding and antibonding orbitals in bond order determination

• Bond order = (number of electrons in bonding orbitals — number of electrons in antibonding orbitals)/2

• Molecules are only stable if bond order >0

Why are the core electrons ignored in MO theory?

• Closer to the nucleus than the valence electrons so have little overlap to the other atom - low degree of interaction

• Bonding and antibonding orbitals are theoretically created but are only slightly different in energy from normal orbitals

• Both core bonding and antibonding orbitals are filled which cancel out in terms of energy

How do the molecular orbitals of Li2 vary from H2?

• Li2 has more diffuse 2s orbitals with weaker overlap than the 1s orbitals in H2

• Reduction in energy of the bonding orbital is therefore less

What are the three different p orbitals?

• px

• py

• pz - one assumed to point towards a bonding atom

How can pz orbitals overlap between two atoms?

• Bonding can be in-phase or out of phase

• In-phase bonding produces one large bonding orbital to be formed between the nuclei - sigma g molecular orbital

• Out-of-phase bonding produces two small antibonding orbitals to be formed between the nuclei - sigma u antibonding orbital

How can px and py orbitals experience bonding/antibonding interactions?

• ‘side-on’ overlap of the orbitals

• Known as pi bonding and pi antibonding orbitals

• pi bonding orbitals combine above and below the nuclei to form two large molecular orbitals

• pi antibonding orbitals form 4 sections with nodal planes between the nuclei and opposite phases

What are the energy levels of sigma and pi orbitals compared to atomic energy levels?

• pi bonding and antibonding orbitals are found between the sigma bonding and antibonding orbital energy levels

• pi bonding and antibonding orbitals have a weaker side on overlap

• Both pi bonding and antibonding effects are less than sigma bonding and antibonding effects

• Energy levels of the pi bonding and antibonding orbitals are closer to the atomic orbital energy levels

What are the HOMO and LUMO in a molecular orbital energy diagram?

• HOMO - highest occupied molecular orbital

• LUMO - lowest unoccupied molecular orbital

How does ionising a molecule affect its bond order?

• Removing an electron from a bonding orbital causes the bond order to go down by ½

• Removing an electron from an antibonding orbital causes the bond order to go up by ½

How does s-p mixing occur in molecular orbital theory?

• Orbitals with the same symmetry labels can interact (sigma g - sigma g)

• The lower energy orbital is pushed further down in energy (stabilised) and the higher energy orbital is push higher up in energy (destabilised)

• The closer in energy the sigma molecular orbitals are the greater the effect

How does the extent of s-p mixing change across a period?

• Energy gap between 2s and 2p increases across the period

• Resulting sigma molecular orbitals become further separated in energy

• Magnitude of s-p mixing decreases across the group

What are homonuclear diatomics?

Molecules made up of two atoms of the same element

What are heteronuclear diatomics?

Molecules made up of two atoms of different elements

How can polarity be explained using MO theory?

• Homonuclear diatomics have the same energy levels for all contributing atomic orbitals

• Heteronuclear diatomics have different energy levels for contributing atomic orbitals

• A molecular orbital closer in energy to one atomic orbital will have a greater contribution from that orbital and electrons within the molecular orbital are more located on that atom

What are the limitations of s-p mixing in terms of orbital symmetry?

• s orbitals can interact with pz orbitals

• s orbitals cannot interact with either px or py orbitals - produce pi orbitals (non-bonding)

How does VSEPR theory predict the shapes of molecules with different numbers of bonding and lone pairs?

• Shape of a molecule can be worked out be considering valence electron pairs on the central atom - both bonding and lone pairs

• Negative electron pairs repel each other and move as far apart as possible

• Lone pairs of electrons will repel more than bonding pairs

How does VSEPR theory account for double and triple bonds within the molecule?

• Double and triple bonds still occupy one position around the central atom

• These bonds will repel slightly more than single bonds

How do you construct a molecular orbital diagram for a molecule with three atoms?

• Confirm the shape of the molecule using VSEPR theory

• Analyse the combinations of the outer Xn orbitals - in-phase and out-of-phase combinations

• Mix these with the atomic orbitals of the central atom

What change needs to be made when considering atoms with many electrons (eg Xe)?

• Construct a partial molecular orbital scheme

• Only considers the interactions that produce the key molecular orbitals (including the HOMO and LUMO)

How are conjugated systems produced across molecules?

• Most significant in conjugated pi systems

• Pi systems - made from pz orbitals not used in the pi bonding that forms the backbone of the molecule

What is electromagnetic radiation?

• Form of energy

• Consists of oscillating electric and magnetic fields

• Travels through space at the speed of light - 2.998 × 10^8 m s^-1

• Exhibits wave-particle duality - can be categorised by frequency/wavelength and quantised particles (photons)

What are the three equations needed for EM calculations?

• E = hv

• c = lamda x v

• E = hvL

How does UV-Vis spectroscopy work?

• Measures the absorption of energy caused by the excitation of energy

• Electrons are excited from low energy molecular orbitals to higher energy molecular orbitals

• Known as electronic transitions

• Uses EM radiation in the range of 200-1000nm

• Dispersive elements separate the wavelengths of light

• Measures transmittance and absorbance

Is UV-Vis spectroscopy quantitative or qualitative?

• Both

• Used quantitatively to determine the concentration of a sample

• Used qualitatively to determine the identity of a species to a known spectrum or as a method to follow the progress of a reaction

How is a UV-Vis spectrum presented graphically?

• x-axis: wavelength - units of nm

• y-axis: absorbance - no units

• Position of peaks on the x-axis depends on their energy of transition - energy gap between orbitals determines the wavelengths absorbed by the sample

• Peak intensity depends on the type of electronic transition and the concentration of the sample

How is UV-Vis spectroscopy used to determine the concentration of a sample?

• Beer-Lambert Law: A = e x c x l

• A: absorbance (no units)

• e: molecular extinction coefficient (dm^3 mol^-1 cm-1)

• c: concentration (mol dm^-3)

• l: pathlength (cm)

What is the molecular extinction coefficient?

• Unique to the molecule

• Unique to the electronic transition/wavelength - each peak will have a different extinction coefficient

What is the pathlength?

• Distance the electromagnetic radiation passes through

• Normally dictated by the depth of the cuvette

How do electronic transitions occur between molecular orbitals?

• Electrons are excited from the highest occupied molecular orbital (HOMO) to the lowest unoccupied molecular orbital (LUMO)

• The further apart the HOMO and LUMO are in terms of energy, the shorter the wavelength required to excite the electrons

Which orbitals will electrons be able to transition between?

• Between bonding and antibonding orbitals

• Electrons are always promoted from a bonding to antibonding orbital

Which transitions do not take place for electrons?

• sigma to sigma*

• pi to sigma*

• n to sigma*

Why do x - sigma* transitions not occur?

• Too high energy - the wavelengths needed are <200nm (far-UV region)

• Too much background noise to see specific peaks

• Latter two (pi/non-bonding to sigma*) transitions are symmetry forbidden - have low absorbance intensities

Which electronic transitions are useful between molecular orbitals?

pi to pi* and n to pi*

How do pi to pi* transitions occur and why are they useful?

• Electrons in pi bonds require less energy for excitation

• Isolated C=C bond gives a strong absorbance peak

• Delocalisation of pi-electrons in conjugated molecules reduces the energy gap between the HOMO and LUMO (including aromatic molecules)

• Longer conjugated systems cause longer wavelengths to be needed - HOMO and LUMO continue getting closer in energy

• Symmetry allowed transition (p-orbitals)

How do n to pi* transitions occur and why are they less useful than pi to pi* transitions?

• Transition associated with lone pairs (O, N, or S)

• Officially a symmetry forbidden transition (px/py to pz) - peaks are often weak in intensity

• Non-bonding electrons in heteroatoms can become involved in resonance and extend conjugation

Which three amino acid side chains allow UV-Vis spectroscopy to be used?

• Tryptophan

• Tyrosine

• Phenylalanine

Which of the three conjugated amino acids is not considered when using UV-Vis spectroscopy?

• Phenylalanine

• Highest energy requirement so smallest wavelength and lowest extinction coefficient

What residue is used in UV-Vis spectroscopy rather than phenylalanine?

• Cystine residues (formed by a disulphide bond between two cysteines)

• Lone pairs of electrons on S atoms allow for pi to pi* transitions

• Proteins can be denatured using denaturants (do not alone break disulphide bonds)

• Protein chromophores resemble those of the isolated amino acid model compounds when denatured

What does the extinction coefficient of a protein depend on?

• Number of each amino acid chromophore

• Solvent exposure of those chromophores

How can protein concentrations be determined experimentally utilising UV-Vis spectroscopy?

Use the Edelhoc method - compares the absorbance at 280nm between native and denatured forms of the protein

What is the method for completing the Edelhoc method?

• Record a UV-Vis absorption spectrum of the native protein dissolved in biological buffer

• Record a UV-Vis absorption spectrum of denatured protein in the presence of a denaturant (6M GdnHCl or 8M urea)

• Determine the number of Tryptophan, Tyrosine and Cystine residues present

• Calculate the denatured extinction coefficient for model compounds in either denaturant

• Calculate the native extinction coefficient