CHEM NOTES (Y1) (L2-L20)
Here are notes addressing the learning objectives you provided, drawing on the information from the sources:
1. Describe the structure of an atom in terms of orbitals
An atom consists of a small, positively charged nucleus containing protons and neutrons, surrounded by
electrons that travel around the nucleus. The number of protons determines the element. Electrons are located
in specific regions of space called atomic orbitals. Each electron in an atom is described by four quantum
numbers: the principal quantum number (n) indicating the energy level or shell, the azimuthal quantum
number (l) defining the shape of the orbital (s, p, d, etc.), the magnetic quantum number (ml) specifying the
orientation of the orbital in space, and the spin magnetic quantum number (ms) describing the electron spin.
According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of four
quantum numbers. s orbitals (l=0) are spherical, p orbitals (l=1) are dumbbell-shaped and oriented along x, y,
and z axes, and d orbitals (l=2) have more complex shapes. Orbital lobes represent regions of high electron
probability, and nodal planes are regions with zero electron probability.
2. Compare the electronic properties of different elements
The electronic configuration of an element describes how its electrons are arranged in atomic orbitals and is
determined by the Aufbau principle (filling lower energy orbitals first), the n+l rule (lower n+l filled first, then
lower n), and Hund's first rule (degenerate orbitals filled singly before pairing). Different elements have
different numbers of protons and therefore different numbers of electrons, leading to unique electronic
configurations. Valence electrons are those in the outermost shell and are primarily involved in chemical
bonding. Elements strive to achieve a stable, filled outer electron shell. Electronegativity (χ), the ability of an
atom to attract electrons in a bond, varies between elements and influences the type of chemical bond formed.
Effective nuclear charge (Zeff), the net positive charge experienced by valence electrons, also differs among
elements and affects atomic size and ionisation energy. For example, oxygen has a higher effective charge than
carbon, leading to a smaller atomic radius and higher electronegativity. The energy of atomic orbitals is also
influenced by the effective nuclear charge; a higher effective charge lowers the orbital energy.
3. Sketch simple MO diagrams and assess the resulting bonding interactions
Molecular orbitals (MOs) are formed by the linear combination of atomic orbitals (LCAO) of two or
more atoms. Atomic orbitals can combine in-phase (additive) to form lower-energy bonding MOs or out-of-
phase (subtractive) to form higher-energy antibonding MOs. The number of MOs formed equals the number
of AOs combined. To sketch an MO diagram:
• Draw the energy levels of the atomic orbitals of the participating atoms.
• Combine AOs of similar energy and appropriate symmetry to form bonding and antibonding MOs,
placing them at lower and higher energy levels, respectively.
• Fill the MOs with the total number of valence electrons from the participating atoms, following the
Pauli Exclusion Principle and Hund's rule.
The bond order is calculated as ½ * (number of electrons in bonding MOs - number of electrons in
antibonding MOs). A positive bond order indicates a stable bond, a higher bond order suggests a stronger
and shorter bond, and a bond order of zero implies that a stable bond does not form. The filling of bonding
MOs contributes to the stability of the molecule, while the filling of antibonding MOs counteracts this stability.
4. Compare and contrast different chemical bonds
Chemical bonds hold atoms together in molecules. The main types of chemical bonds discussed are:
• Ionic Bonds: Formed by the transfer of electrons between atoms with a large difference in
electronegativity (typically > 1.7), resulting in positively and negatively charged ions held together by
electrostatic attraction. These are typically strong, stable bonds. Example: NaCl.
, • Covalent Bonds: Formed by the sharing of electrons between atoms with a small difference in
electronegativity (typically < 0.7).
o Nonpolar Covalent Bonds: Electrons are shared equally. Example: C-C bond in ethane.
o Polar Covalent Bonds: Electrons are shared unequally due to a moderate difference in
electronegativity (typically 0.7 - 1.7), resulting in partial charges on the atoms. Example: O-H
bond in water. Covalent bonds can be single (σ bond, formed by head-on overlap), double (σ + π
bond, π bond from side-by-side overlap of p orbitals), or triple (σ + 2π bonds). Double and
triple bonds are stronger and shorter than single bonds and restrict rotation.
5. Explain how delocalisation stabilises molecular structures
Delocalised electrons are electrons that are not confined to a bond between two specific atoms but are spread
over three or more atoms in a molecule. This often occurs in molecules with conjugated π systems (alternating
single and multiple bonds) or where resonance structures can be drawn. In these cases, the π electrons occupy
molecular orbitals that extend over several atoms, rather than being localised in a single π bond. This
sharing of electron density over a larger area lowers the overall energy of the molecule, making it more
stable than if the electrons were localised. The extra stability gained by delocalisation is called resonance energy
or delocalisation energy. Examples include benzene, where the π electrons are delocalised over the entire ring,
and the peptide bond, which has partial double bond character due to delocalisation of the nitrogen lone pair
into the carbonyl group.
6. Describe and explain the structure of simple molecules
The structure of simple molecules is determined by the types of atoms present, the way they are bonded
together (connectivity), and their spatial arrangement (geometry). Lewis structures can represent the
connectivity of atoms and the presence of lone pairs. The Valence Shell Electron Pair Repulsion (VSEPR)
theory predicts the geometry around a central atom by minimising the repulsion between electron pairs (both
bonding and non-bonding) in the valence shell. The number of electron pairs dictates the basic electron pair
geometry (e.g., linear, trigonal planar, tetrahedral). The molecular geometry (arrangement of atoms) may differ
from the electron pair geometry if lone pairs are present, as lone pairs exert stronger repulsion, affecting bond
angles. Hybridisation of atomic orbitals (e.g., sp, sp2, sp3) can also be used to explain molecular geometries
and bonding. The three-dimensional structure is further defined by bond lengths (distance between nuclei),
bond angles (angle between three bonded atoms), and dihedral angles (angle between two planes of three
atoms).
7. Apply MO theory to explain the properties of key molecular structures encountered in biomolecules
such as peptide bonds or aromatic rings
• Peptide Bond: The peptide bond (-CO-NH-) linking amino acids in proteins exhibits partial double
bond character due to the delocalisation of the lone pair of electrons on the nitrogen atom into the π
system of the carbonyl group. MO theory explains this by showing the formation of π molecular orbitals
that extend over the O-C-N atoms, resulting in increased rigidity and a planar geometry around the
peptide bond. This also contributes to the restricted rotation around the C-N bond and influences
protein structure.
• Aromatic Rings (e.g., Benzene): Benzene, a key component of aromatic amino acids, has a planar
hexagonal ring of six carbon atoms, each sp2 hybridised with a p orbital perpendicular to the ring. MO
theory describes the combination of these six p orbitals to form delocalised π molecular orbitals that
extend above and below the plane of the ring. This delocalisation of π electrons results in significant
stabilisation (aromaticity) and explains the characteristic chemical reactivity and spectroscopic
properties of aromatic compounds. The delocalised electron cloud can participate in π-π interactions
and π-cation interactions.
8. Integrate specific terminology in your descriptions of molecules and chemical bonding
When describing molecules and chemical bonding, it is important to use precise terminology, including terms
such as: atoms, nucleus, protons, neutrons, electrons, orbitals (s, p, d), quantum numbers (n, l, ml, ms),
,electron configuration, valence electrons, electronegativity, ionisation energy, effective nuclear charge,
chemical bond (ionic, covalent, polar), sigma (σ) bond, pi (π) bond, molecular orbital (bonding,
antibonding), bond order, hybridisation (sp, sp2, sp3), VSEPR theory, molecular geometry (linear,
trigonal planar, tetrahedral, etc.), bond length, bond angle, dihedral angle, lone pair, delocalisation,
resonance, resonance structure, functional group (alcohol, ether, amine, aldehyde, ketone, carboxylic
acid, etc.), nucleophile, electrophile.. Consistent and correct use of these terms is essential for clear
communication in chemistry.
9. Catalogue the different interactions between atoms in biomolecules
Biomolecules are stabilised and interact through various types of interactions:
• Covalent Bonds: Strong bonds that hold atoms together within molecules (e.g., peptide bonds, disulfide
bonds).
• Non-covalent Interactions: Weaker, but collectively important for structure and function.
o Hydrogen Bonds: Occur between a hydrogen atom bonded to an electronegative atom (donor)
and another electronegative atom with a lone pair (acceptor). Crucial for protein folding and
DNA base pairing.
o Electrostatic Interactions (Ionic Bonds/Salt Bridges): Attractions between oppositely
charged groups. Salt bridges can also involve hydrogen bonds.
o Van der Waals Forces: Weak, short-range interactions. Include:
§ London Dispersion Forces: Present between all molecules due to temporary
fluctuations in electron density.
§ Dipole-Dipole Interactions: Occur between molecules with permanent dipoles.
o Hydrophobic Effect: The tendency of nonpolar molecules or nonpolar parts of molecules to
aggregate in an aqueous environment to minimise unfavorable interactions with water and
increase the entropy of water. Drives protein folding and membrane formation.
o π-π Interactions: Interactions between the electron clouds of aromatic rings.
o π-Cation Interactions: Attraction between an aromatic ring and a positively charged ion.
10. Explain the difference between strong and weak acids in terms of dissociation properties
Acids are proton (H+) donors.
• Strong acids dissociate completely or very nearly completely into their ions when dissolved in water.
For example, HCl → H+ + Cl-. The concentration of the undissociated acid is negligible in solution.
• Weak acids only partially dissociate into their ions in water, establishing an equilibrium between the
undissociated acid and its conjugate base and H+ ions. For example, a weak acid HA dissociates
according to HA <=> H+ + A-. At equilibrium, a significant amount of the weak acid remains
undissociated. The extent of dissociation depends on the acid's strength, quantified by its acid
dissociation constant (Ka) or pKa.
11. Define and explain pKa and its importance in biological buffering
pKa is the negative logarithm (base 10) of the acid dissociation constant (Ka): pKa = -log10(Ka). Ka is a
measure of the strength of a weak acid; a larger Ka (smaller pKa) indicates a stronger acid. In biological
buffering, pKa is crucial because a buffer solution, which resists changes in pH, is most effective when the pH
of the solution is close to the pKa of the weak acid component of the buffer. At pH = pKa, the concentrations
of the weak acid and its conjugate base are equal, providing the maximum capacity to neutralise added acid or
base. Biological systems maintain narrow pH ranges essential for enzyme activity and cellular processes.
Physiological buffers, such as the bicarbonate and phosphate systems in the blood and intracellular fluids, have
pKa values near the physiological pH (around 7.4) to function effectively in maintaining pH homeostasis.
Amino acid side chains in proteins also contribute to buffering, and their effectiveness depends on their
respective pKa values relative to the physiological pH.
12. Use the Henderson-Hasselbalch equation
, The Henderson-Hasselbalch equation relates the pH of a solution containing a weak acid and its conjugate
base to the pKa of the acid and the relative concentrations of the acid and base: pH = pKa + log10([A-
]/[HA]) where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid.
This equation is used to:
• Calculate the pH of a buffer solution.
• Determine the ratio of conjugate base to weak acid required to achieve a desired pH.
• Understand how changes in the concentrations of the buffer components affect the pH. For example, if
the concentration of the conjugate base [A-] is ten times that of the weak acid [HA], then pH = pKa +
log10(10) = pKa + 1. If [A-] = [HA], then pH = pKa + log10(1) = pKa.
13. List a number of physiological buffers
Physiological buffers help maintain stable pH in biological systems. Examples include:
• Bicarbonate buffer system (H2CO3/HCO3-) in blood plasma.
• Phosphate buffer system (H2PO4-/HPO4^2-) in intracellular fluids and urine.
• Proteins, including haemoglobin in red blood cells, which have amino acid residues with pKa values
near physiological pH.
14. Explain how proteins can be used as physiological buffers
Proteins contain numerous amino acid residues, many of which have ionizable side chains (e.g., carboxyl,
amino, imidazole) that can act as weak acids or bases. Each ionizable group has a characteristic pKa value.
When the pH of the surrounding environment changes, these side chains can donate or accept protons (H+),
helping to minimise the change in pH. For example, if pH rises (OH- increases), acidic side chains can release
H+ to neutralise the OH-. If pH drops (H+ increases), basic side chains can accept the excess H+. The overall
buffering capacity of a protein depends on the number and types of ionizable residues and their pKa values
relative to the physiological pH.
15. Describe the function of Haemoglobin as a buffer for Hydrogen ions produced in metabolism and
explain what allows this to happen
Haemoglobin, the oxygen-carrying protein in red blood cells, also acts as a significant buffer in the blood.
During metabolism, carbon dioxide (CO2) is produced, which dissolves in the blood and forms carbonic acid
(H2CO3), which then dissociates into H+ and bicarbonate (HCO3-). The increase in H+ lowers the blood
pH. Haemoglobin buffers this H+ by binding to it, particularly in its deoxy form (when oxygen is released to
tissues). The imidazole side chain of histidine residues in haemoglobin has a pKa value close to the
physiological pH of blood. This allows histidine to effectively accept or donate protons within this pH range,
neutralising the excess H+ ions and preventing a large drop in blood pH. The binding of H+ to haemoglobin is
also allosterically linked to the release of oxygen (Bohr effect), further enhancing the efficiency of oxygen
delivery and pH regulation.
16. Describe and visualise organic structures following agreed conventions
Organic molecules are based on a hydrocarbon framework with functional groups containing heteroatoms (O,
N, S, halogens). Standard conventions for drawing organic structures include:
• Line-angle (skeletal) formulas: Carbon atoms are implied at the ends of lines and vertices, not
explicitly drawn with a 'C'.
• Hydrogen atoms attached to carbon are usually not shown, their presence is inferred by carbon's
valency of four.
• Heteroatoms (and any hydrogens attached to them) are always shown with their chemical
symbols.
• Bonds are represented by lines: single bonds by one line, double bonds by two, and triple bonds by
three.
Here are notes addressing the learning objectives you provided, drawing on the information from the sources:
1. Describe the structure of an atom in terms of orbitals
An atom consists of a small, positively charged nucleus containing protons and neutrons, surrounded by
electrons that travel around the nucleus. The number of protons determines the element. Electrons are located
in specific regions of space called atomic orbitals. Each electron in an atom is described by four quantum
numbers: the principal quantum number (n) indicating the energy level or shell, the azimuthal quantum
number (l) defining the shape of the orbital (s, p, d, etc.), the magnetic quantum number (ml) specifying the
orientation of the orbital in space, and the spin magnetic quantum number (ms) describing the electron spin.
According to the Pauli Exclusion Principle, no two electrons in an atom can have the same set of four
quantum numbers. s orbitals (l=0) are spherical, p orbitals (l=1) are dumbbell-shaped and oriented along x, y,
and z axes, and d orbitals (l=2) have more complex shapes. Orbital lobes represent regions of high electron
probability, and nodal planes are regions with zero electron probability.
2. Compare the electronic properties of different elements
The electronic configuration of an element describes how its electrons are arranged in atomic orbitals and is
determined by the Aufbau principle (filling lower energy orbitals first), the n+l rule (lower n+l filled first, then
lower n), and Hund's first rule (degenerate orbitals filled singly before pairing). Different elements have
different numbers of protons and therefore different numbers of electrons, leading to unique electronic
configurations. Valence electrons are those in the outermost shell and are primarily involved in chemical
bonding. Elements strive to achieve a stable, filled outer electron shell. Electronegativity (χ), the ability of an
atom to attract electrons in a bond, varies between elements and influences the type of chemical bond formed.
Effective nuclear charge (Zeff), the net positive charge experienced by valence electrons, also differs among
elements and affects atomic size and ionisation energy. For example, oxygen has a higher effective charge than
carbon, leading to a smaller atomic radius and higher electronegativity. The energy of atomic orbitals is also
influenced by the effective nuclear charge; a higher effective charge lowers the orbital energy.
3. Sketch simple MO diagrams and assess the resulting bonding interactions
Molecular orbitals (MOs) are formed by the linear combination of atomic orbitals (LCAO) of two or
more atoms. Atomic orbitals can combine in-phase (additive) to form lower-energy bonding MOs or out-of-
phase (subtractive) to form higher-energy antibonding MOs. The number of MOs formed equals the number
of AOs combined. To sketch an MO diagram:
• Draw the energy levels of the atomic orbitals of the participating atoms.
• Combine AOs of similar energy and appropriate symmetry to form bonding and antibonding MOs,
placing them at lower and higher energy levels, respectively.
• Fill the MOs with the total number of valence electrons from the participating atoms, following the
Pauli Exclusion Principle and Hund's rule.
The bond order is calculated as ½ * (number of electrons in bonding MOs - number of electrons in
antibonding MOs). A positive bond order indicates a stable bond, a higher bond order suggests a stronger
and shorter bond, and a bond order of zero implies that a stable bond does not form. The filling of bonding
MOs contributes to the stability of the molecule, while the filling of antibonding MOs counteracts this stability.
4. Compare and contrast different chemical bonds
Chemical bonds hold atoms together in molecules. The main types of chemical bonds discussed are:
• Ionic Bonds: Formed by the transfer of electrons between atoms with a large difference in
electronegativity (typically > 1.7), resulting in positively and negatively charged ions held together by
electrostatic attraction. These are typically strong, stable bonds. Example: NaCl.
, • Covalent Bonds: Formed by the sharing of electrons between atoms with a small difference in
electronegativity (typically < 0.7).
o Nonpolar Covalent Bonds: Electrons are shared equally. Example: C-C bond in ethane.
o Polar Covalent Bonds: Electrons are shared unequally due to a moderate difference in
electronegativity (typically 0.7 - 1.7), resulting in partial charges on the atoms. Example: O-H
bond in water. Covalent bonds can be single (σ bond, formed by head-on overlap), double (σ + π
bond, π bond from side-by-side overlap of p orbitals), or triple (σ + 2π bonds). Double and
triple bonds are stronger and shorter than single bonds and restrict rotation.
5. Explain how delocalisation stabilises molecular structures
Delocalised electrons are electrons that are not confined to a bond between two specific atoms but are spread
over three or more atoms in a molecule. This often occurs in molecules with conjugated π systems (alternating
single and multiple bonds) or where resonance structures can be drawn. In these cases, the π electrons occupy
molecular orbitals that extend over several atoms, rather than being localised in a single π bond. This
sharing of electron density over a larger area lowers the overall energy of the molecule, making it more
stable than if the electrons were localised. The extra stability gained by delocalisation is called resonance energy
or delocalisation energy. Examples include benzene, where the π electrons are delocalised over the entire ring,
and the peptide bond, which has partial double bond character due to delocalisation of the nitrogen lone pair
into the carbonyl group.
6. Describe and explain the structure of simple molecules
The structure of simple molecules is determined by the types of atoms present, the way they are bonded
together (connectivity), and their spatial arrangement (geometry). Lewis structures can represent the
connectivity of atoms and the presence of lone pairs. The Valence Shell Electron Pair Repulsion (VSEPR)
theory predicts the geometry around a central atom by minimising the repulsion between electron pairs (both
bonding and non-bonding) in the valence shell. The number of electron pairs dictates the basic electron pair
geometry (e.g., linear, trigonal planar, tetrahedral). The molecular geometry (arrangement of atoms) may differ
from the electron pair geometry if lone pairs are present, as lone pairs exert stronger repulsion, affecting bond
angles. Hybridisation of atomic orbitals (e.g., sp, sp2, sp3) can also be used to explain molecular geometries
and bonding. The three-dimensional structure is further defined by bond lengths (distance between nuclei),
bond angles (angle between three bonded atoms), and dihedral angles (angle between two planes of three
atoms).
7. Apply MO theory to explain the properties of key molecular structures encountered in biomolecules
such as peptide bonds or aromatic rings
• Peptide Bond: The peptide bond (-CO-NH-) linking amino acids in proteins exhibits partial double
bond character due to the delocalisation of the lone pair of electrons on the nitrogen atom into the π
system of the carbonyl group. MO theory explains this by showing the formation of π molecular orbitals
that extend over the O-C-N atoms, resulting in increased rigidity and a planar geometry around the
peptide bond. This also contributes to the restricted rotation around the C-N bond and influences
protein structure.
• Aromatic Rings (e.g., Benzene): Benzene, a key component of aromatic amino acids, has a planar
hexagonal ring of six carbon atoms, each sp2 hybridised with a p orbital perpendicular to the ring. MO
theory describes the combination of these six p orbitals to form delocalised π molecular orbitals that
extend above and below the plane of the ring. This delocalisation of π electrons results in significant
stabilisation (aromaticity) and explains the characteristic chemical reactivity and spectroscopic
properties of aromatic compounds. The delocalised electron cloud can participate in π-π interactions
and π-cation interactions.
8. Integrate specific terminology in your descriptions of molecules and chemical bonding
When describing molecules and chemical bonding, it is important to use precise terminology, including terms
such as: atoms, nucleus, protons, neutrons, electrons, orbitals (s, p, d), quantum numbers (n, l, ml, ms),
,electron configuration, valence electrons, electronegativity, ionisation energy, effective nuclear charge,
chemical bond (ionic, covalent, polar), sigma (σ) bond, pi (π) bond, molecular orbital (bonding,
antibonding), bond order, hybridisation (sp, sp2, sp3), VSEPR theory, molecular geometry (linear,
trigonal planar, tetrahedral, etc.), bond length, bond angle, dihedral angle, lone pair, delocalisation,
resonance, resonance structure, functional group (alcohol, ether, amine, aldehyde, ketone, carboxylic
acid, etc.), nucleophile, electrophile.. Consistent and correct use of these terms is essential for clear
communication in chemistry.
9. Catalogue the different interactions between atoms in biomolecules
Biomolecules are stabilised and interact through various types of interactions:
• Covalent Bonds: Strong bonds that hold atoms together within molecules (e.g., peptide bonds, disulfide
bonds).
• Non-covalent Interactions: Weaker, but collectively important for structure and function.
o Hydrogen Bonds: Occur between a hydrogen atom bonded to an electronegative atom (donor)
and another electronegative atom with a lone pair (acceptor). Crucial for protein folding and
DNA base pairing.
o Electrostatic Interactions (Ionic Bonds/Salt Bridges): Attractions between oppositely
charged groups. Salt bridges can also involve hydrogen bonds.
o Van der Waals Forces: Weak, short-range interactions. Include:
§ London Dispersion Forces: Present between all molecules due to temporary
fluctuations in electron density.
§ Dipole-Dipole Interactions: Occur between molecules with permanent dipoles.
o Hydrophobic Effect: The tendency of nonpolar molecules or nonpolar parts of molecules to
aggregate in an aqueous environment to minimise unfavorable interactions with water and
increase the entropy of water. Drives protein folding and membrane formation.
o π-π Interactions: Interactions between the electron clouds of aromatic rings.
o π-Cation Interactions: Attraction between an aromatic ring and a positively charged ion.
10. Explain the difference between strong and weak acids in terms of dissociation properties
Acids are proton (H+) donors.
• Strong acids dissociate completely or very nearly completely into their ions when dissolved in water.
For example, HCl → H+ + Cl-. The concentration of the undissociated acid is negligible in solution.
• Weak acids only partially dissociate into their ions in water, establishing an equilibrium between the
undissociated acid and its conjugate base and H+ ions. For example, a weak acid HA dissociates
according to HA <=> H+ + A-. At equilibrium, a significant amount of the weak acid remains
undissociated. The extent of dissociation depends on the acid's strength, quantified by its acid
dissociation constant (Ka) or pKa.
11. Define and explain pKa and its importance in biological buffering
pKa is the negative logarithm (base 10) of the acid dissociation constant (Ka): pKa = -log10(Ka). Ka is a
measure of the strength of a weak acid; a larger Ka (smaller pKa) indicates a stronger acid. In biological
buffering, pKa is crucial because a buffer solution, which resists changes in pH, is most effective when the pH
of the solution is close to the pKa of the weak acid component of the buffer. At pH = pKa, the concentrations
of the weak acid and its conjugate base are equal, providing the maximum capacity to neutralise added acid or
base. Biological systems maintain narrow pH ranges essential for enzyme activity and cellular processes.
Physiological buffers, such as the bicarbonate and phosphate systems in the blood and intracellular fluids, have
pKa values near the physiological pH (around 7.4) to function effectively in maintaining pH homeostasis.
Amino acid side chains in proteins also contribute to buffering, and their effectiveness depends on their
respective pKa values relative to the physiological pH.
12. Use the Henderson-Hasselbalch equation
, The Henderson-Hasselbalch equation relates the pH of a solution containing a weak acid and its conjugate
base to the pKa of the acid and the relative concentrations of the acid and base: pH = pKa + log10([A-
]/[HA]) where [A-] is the concentration of the conjugate base and [HA] is the concentration of the weak acid.
This equation is used to:
• Calculate the pH of a buffer solution.
• Determine the ratio of conjugate base to weak acid required to achieve a desired pH.
• Understand how changes in the concentrations of the buffer components affect the pH. For example, if
the concentration of the conjugate base [A-] is ten times that of the weak acid [HA], then pH = pKa +
log10(10) = pKa + 1. If [A-] = [HA], then pH = pKa + log10(1) = pKa.
13. List a number of physiological buffers
Physiological buffers help maintain stable pH in biological systems. Examples include:
• Bicarbonate buffer system (H2CO3/HCO3-) in blood plasma.
• Phosphate buffer system (H2PO4-/HPO4^2-) in intracellular fluids and urine.
• Proteins, including haemoglobin in red blood cells, which have amino acid residues with pKa values
near physiological pH.
14. Explain how proteins can be used as physiological buffers
Proteins contain numerous amino acid residues, many of which have ionizable side chains (e.g., carboxyl,
amino, imidazole) that can act as weak acids or bases. Each ionizable group has a characteristic pKa value.
When the pH of the surrounding environment changes, these side chains can donate or accept protons (H+),
helping to minimise the change in pH. For example, if pH rises (OH- increases), acidic side chains can release
H+ to neutralise the OH-. If pH drops (H+ increases), basic side chains can accept the excess H+. The overall
buffering capacity of a protein depends on the number and types of ionizable residues and their pKa values
relative to the physiological pH.
15. Describe the function of Haemoglobin as a buffer for Hydrogen ions produced in metabolism and
explain what allows this to happen
Haemoglobin, the oxygen-carrying protein in red blood cells, also acts as a significant buffer in the blood.
During metabolism, carbon dioxide (CO2) is produced, which dissolves in the blood and forms carbonic acid
(H2CO3), which then dissociates into H+ and bicarbonate (HCO3-). The increase in H+ lowers the blood
pH. Haemoglobin buffers this H+ by binding to it, particularly in its deoxy form (when oxygen is released to
tissues). The imidazole side chain of histidine residues in haemoglobin has a pKa value close to the
physiological pH of blood. This allows histidine to effectively accept or donate protons within this pH range,
neutralising the excess H+ ions and preventing a large drop in blood pH. The binding of H+ to haemoglobin is
also allosterically linked to the release of oxygen (Bohr effect), further enhancing the efficiency of oxygen
delivery and pH regulation.
16. Describe and visualise organic structures following agreed conventions
Organic molecules are based on a hydrocarbon framework with functional groups containing heteroatoms (O,
N, S, halogens). Standard conventions for drawing organic structures include:
• Line-angle (skeletal) formulas: Carbon atoms are implied at the ends of lines and vertices, not
explicitly drawn with a 'C'.
• Hydrogen atoms attached to carbon are usually not shown, their presence is inferred by carbon's
valency of four.
• Heteroatoms (and any hydrogens attached to them) are always shown with their chemical
symbols.
• Bonds are represented by lines: single bonds by one line, double bonds by two, and triple bonds by
three.
