Everything except stuff taught by AW and AB and workshop 5 and the tutorial questions

What is valency?

• Atoms want a complete outer shell of electrons

• For example, carbon has 4 electrons in the outer shell, so for a full outer shell, it needs 4 extra electrons. This is done by forming covalent bonds with other atoms and sharing the bond electrons. The valency of carbon is 4.

• Valency of an atom = The number of bonds needed to complete the shell

 

sp³

• Tetrahedral

• 109.5

• All bonds are sigma

sp²

• Trigonal planar

• 120

• Bonds are single to the hydrogens, and a double bonds to carbons (sigma and pi combination)

sp

• Linear

• 180

• Sigma bonds to the hydrogens and 2 pi bonds between the carbons along with a sigma bond, making a triple bond

What’s the hybridisation in these molecules?

Sigma bonds

• Strongest type of covalent chemical bond

• Forms when atomic orbitals overlap in a head on arrangement (like two 2 s orbitals, one s orbital and one p orbital)

• σ

How does sigma bonding affect the shape and flexibility of drug molecules?

Sigma bonds = single bond from orbital overlap means free rotation

How does pi bonding affect the shape and flexibility of drug molecules?

Multiple bond from overlap of p-orbitals means fixed orientation

Double bonds in structures means there is…

restricted rotation

Isotopes

Atoms of the same element with the same number of protons (atomic number) but different number of neutrons (different atomic mass)

What is the kinetic isotope effect?

• The change in the rate of a chemical reaction that occurs when one atom in the reactants is replaced by one of its isotopes

• Shows how the mass difference between isotopes affects how fast a reactions happens

What are the uses of the kinetic isotope effect?

• Determining which bonds are broken in the RDS

• Studying enzyme catalysed reactions

• Investigating reaction transition states and bond strength differences

Name the three types of radiation

Alpha, beta, gamma

The decay of unstable atomic nuclei:

¹⁸F decays to ¹⁸O, with a half life of 109.7 minutes. What is the decay constant for ¹⁸F?

What’s the relationship between the decay constant and the half life?

The decay constant (λ) and the half-life (t₁/₂) of a radioactive isotope are inversely related — when one increases, the other decreases.

What is radiolabelling?

A technique in which a radioactive isotope is incorporated into a molecule so that the molecule can be tracked or detected by measuring the radiation it emits.

Do atoms want a have complete outer valence shell?

Yes

Atoms with nearly empty or nearly complete shells tend to ionise easily to form salts. Is there energy gain in salt formation?

Yes, the energy gain from the electrostatic interaction of the charged species.

This is an example of salt formation - ionic bonding

What is salt formation also known as?

Ionic bonding

Example of Ionic Bonding

NaCl

What is a covalent bond?

A chemical bond that involves the sharing of pairs of electrons between atoms

Example of Covalent Bonding

H₂

• The 1s orbitals of both atoms merge into a single bond orbital that contains both electrons

What is the bond dissociation energy?

• When two atoms bond, their electrons and nuclei arrange into a lower-energy state, releasing energy.

• The bond dissociation energy is the energy required to break this bond and separate the atoms.

• It is a measure of bond strength:

  • Higher bond dissociation energy = stronger bond.

• For a pair of atoms, greater orbital overlap → stronger bond → higher bond dissociation energy.

Name 2 types of chemical covalent bonds:

• Sigma (σ) bonds

• Pi (π) bonds

Are sigma (σ) bonds strong?

• Strongest type of covalent chemical bond

How do sigma (σ) bonds form?

• Forms when atomic orbitals overlap in head on arrangement

Sigma (σ) bond in H₂

The 1s atomic orbitals overlapping to form the new sigma (σ) bonds, also known as a single bond

How do pi (π) bonds form?

Formed from overlap of two orbital lobes on one atom with two orbital lobes on another – in a lateral sense (side to side).

What do pi (π) bonds form with?

An existing sigma bond - thus they are double or triple bonds

Why does NH₃ have bond angles of 107° instead of 109.5° (tetrahedral)

The lone pair exerts a slightly greater repulsion than the sigma bonds, compressing the bond angle from 109.5° to 107°

Why is the bond angle for H₂O 104.5°?

There are 2 lone pairs on the O

What is a hydrogen bond?

A partially electrostatic attraction between a H which is bound to a more electronegative atom such as N, O, or F, and another adjacent atom bearing lone pair of electrons

Most electronegative atoms

N, O, F

Example of Hydrogen Bonding

How are hydrogen bonds typically described?

Hydrogen bonds are typically described as an electrostatic dipole-dipole interaction. The δ+ of the proton is attracted to the lone pair of electrons

Hydrogen bonds are generally regarded as primarily electrostatic, however they have some covalent nature. When we say hydrogen bonds have some covalent nature, it means what?

• A hydrogen bond is not purely electrostatic (attraction between charges).

• There is partial sharing of electrons between the hydrogen atom and the electronegative atom it is hydrogen-bonded to (e.g. O, N, or F).

How do hydrogen bonds have a “covalent nature”?

• It has direction (not like ionic bonding)

• Produces bond distances that are shorter than would be expected from the sum of Van der Waals radii

In hydrogen bonds, the more electronegative the donor…

….the more covalent character is observed

Strength of Hydrogen Bonds

What does hydrogen bonding do in small molecules?

• Increases the melting point

• Increases the boiling point

• Increases the solubility

• Increases the viscosity

What are Van der Waal forces?

• Weak intermolecular attractions that act between all atoms and molecules.

• aka London forces

• Weakest type of intermolecular force

• Present in all molecules

• Strength increases with:

  • More electrons / larger atoms

  • Greater surface area

Where do Van der Waal forces arise from?

• Electrically neutral molecules may exhibit permanent electric dipoles

• A permanent dipole molecule and a non permanent dipole molecule

• Molecules with no permanent dipoles

1) Van der Waals forces

• Some electrically neutral molecules may exhibit permanent electric dipoles.

• For instance consider water – a bent molecule with a bond angle of around 105°.

• The oxygen atom is always somewhat negatively charged and the hydrogen atoms (side) somewhat positive

• These dipoles  have a tendency to align, and this results in a net attractive force.

• Note – this is a lot smaller than any H-bonding interaction present (e.g. in water).

2) Van der Waals forces

• Molecules with a permanent dipole may temporarily distort the electric charge in a nearby molecule (polar or non-polar).

• The extra attraction is between the permanent dipole and the ‘induced’ dipole on the nearby molecule.

3) Van der Waals forces

• Thirdly in molecules with no permanent dipole, temporarily, dipoles result from the random electron motion within the atoms.

• At any given time the centre of positive charge arising from the nucleus and the centre of negative charge arising from the electrons are unlikely to coincide.

• This leads to instantaneous, but short lived dipoles, even though over time the average polarisation is zero.

• The resulting instantaneous dipoles are too short lived to align with other molecules to give an attractive force, however they can induce polarization in adjacent molecules.

• These specific interactions (forces) are known as London forces, or dispersion forces.

What is pi stacking (π- π interactions)?

Pi stacking refers to attractive, noncovalent interactions between aromatic rings (which contain pi (π) bonds).

What are the 3 recognised arrangements of pi stacking (π- π interactions)?

Hydrophobic interactions

• Water features a relatively strong network of hydrogen bonds.

• In solution this network is dynamic (constantly changing).

• Non-polar molecules, such as hydrocarbon chains, cannot form hydrogen bonds, and thus attempts to mix result in breaking up of the h-bond network.

What is an amphiphile?

Molecules with both hydrophobic and hydrophilic domains

What will amphiphilic molecules form in water?

Micelles

• Hydrophilic heads outside

• Hydrophobic tails inside

Why are micelles biologically important?

• Proteins obviously contain residues (amino acids)

• Amino acids can have strongly hydrophobic elements, like glycine, alanine, valine, etc

• When the protein folds into its 3D shape, it is common to have the hydrophobic core compose of these residues

• Charged and polar residues can interact with the surrounding water molecules

• Minimizing the number of hydrophobic solvation sphere is the principle driving force behind the folding process

Primary structure of Protein

Refers to the sequence of amino acids in the polypeptide chain. The chain is held together by peptide bonds (formed from the acid group of one and the amine group of the other).

Secondary structure of Protein

Refers to highly regular sub-structures on the polypeptide backbone chain. Two main types are the α-helix and the β-sheet, and they are defined by patterns of hydrogen bonds between the main-chain peptide groups.

Tertiary structure of Protein

Refers to the overall three dimensional shape of the protein molecule. The α-helices and the β-sheets are folded into a compact globular structure. This is primarily driven by the non-specific hydrophobic interactions (the burial of the hydrophobic residues in the core, away from water). However further stability is also introduced by some specific tertiary interactions such as salt bridges, hydrogen bonds, the tight packing of side chains (van der Waals) and disulphide bonds.

Quaternary structure of Protein

Refers to assemblies of two or more individual polypeptide chains into one single functional unit (stabilized in a similar way to the tertiary structure).

The primary structure of a protein is reported starting at which end?

Amino terminal (N) end

Protein secondary structure - β-sheets

Protein secondary structure – α-helix

The carbonyl of residue n interacts with the amide proton of residue n+4.

Tertiary and Quaternary structure

• The hydrophobic effect is the primary driving force in protein folding – the hydrophobic residues are forced into core of the forming globular structure, while the hydrophilic residues remain at the surface.

• The gain in energy resulting from van der Waals forces in the hydrophobic regions is a big contributor to the stability of the folded protein.

• Additional factors are π-stacking, salt bridges and disulphide bonds.

Definition of Electronegativity

A chemical property that describes the tendency of an atom to attract a shared pair of electrons (or electron density) towards itself

Electronegativity is determined by what?

• Nuclear charge (more protons = more pull)

• Number and location of electrons in atomic shells

How does electronegativity change across a periodic table?

Increases from left to right of the periodic table, and from bottom to top

• F is the most electronegative

What explains the reduction in electronegativity when you descend a group?

There is extra shielding of the shells of electrons

What effect does electronegativity have on charge distribution of a molecule?

When there are difference in electronegativity, atoms will become slightly positive (δ+) and slightly negative (δ-)

This leads to induction and inductive effect

The stronger the electronegativity….

….the stronger the inductive effect

What is the inductive effect?

The permanent shifting of electron density along σ-bonds due to electronegativity differences.

What is a -I group?

An electron-withdrawing group that pulls electron density away through σ-bonds

X =

• Br

• Cl

• NO₂

• OH

• SH

• SR

• NH₂

• NHR

• NR₂

• CN

• COOH

• CHO

• C(O)R

What is a +I group?

An electron-donating group that pushes electron density through σ-bonds (e.g. alkyl groups, metals).

X =

• R (alkyl or aryl)

• Metals (e.g. Li, Mg)

How does distance affect the inductive effect?

The inductive effect decreases rapidly as distance from the substituent increases.

What does pKa measure?

Acid strength.

Lower pKa = stronger acid = more willing to donate H⁺.

How do electron-withdrawing groups affect pKa of acids?

They stabilise the conjugate base → lower the pKa → increase acidity.

How do electron-donating groups affect pKa of acids?

They destabilise the conjugate base → raise the pKa → decrease acidity.

What is resonance?

The delocalisation of electrons over several atoms, represented by multiple canonical forms.

Why does resonance stabilise molecules/ions?

Because charge is spread over a larger area, lowering energy and increasing stability.

Why are benzene bonds all the same length?

• Because π electrons are delocalised over the ring (true structure is a resonance hybrid).

Carboxylate resonance forms

Why are resonance forms useful?

• Resonance forms are useful for showing electron movement and reactivity, even though none is fully correct.

• Curly arrows are used to show the movement of bond electrons between resonance forms.

Definition of the Mesomeric effect

Electron donation or withdrawal through resonance involving π systems and lone pairs.

• Double bond adjacent to a lone pair of electrons

+M

When electron density is pushed into the π-bond (orbital).

-M

Where a π-orbital overlaps with an adjacent p-orbital that is low in electron density.

How do +M groups affect phenol acidity?

They destabilise the phenoxide ion → increase pKa → make phenol less acidic.

How do –M groups affect phenol acidity?

They stabilise the phenoxide ion → decrease pKa → make phenol more acidic.

What does pKa mean for bases?

It refers to the protonated form of the base; higher pKa = stronger base.

What happens when pH = pKa for a base?

About 50% of the base is protonated.

Why are alkyl amines more basic than ammonia?

Alkyl groups donate electron density (+I), making nitrogen more able to accept H⁺.

Why is phenylamine much less basic than cyclohexylamine?

The lone pair in phenylamine is delocalised into the aromatic ring, reducing availability to bind H⁺.

Why is pyrrole very weakly basic?

Its lone pair is part of the aromatic system and cannot accept a proton.

Why are amides extremely weak bases?

The lone pair on nitrogen is delocalised into the carbonyl group by resonance.

Why are carbonyl groups polar?

Oxygen is more electronegative than carbon, giving O δ⁻ and C δ⁺.

How do substituents affect carbonyl reactivity?

• Electron-withdrawing groups increase δ⁺ on carbon → more reactive

• Electron-donating groups decrease δ⁺ → less reactive

Alcohol → Aldehyde → Carboxylic acid

Ester

Ether

Ketone

α - lactone

β - lactone

γ - lactone

δ - lactone

Carboxylic acid → Carboxylate