Biochemistry Overview: Lecture Series on Thermodynamics

Biochemistry Overview

• The lecture series extends chapter 5 of the textbook.
• The textbook's chapter 5 is considered a basic introduction.
• More detailed content is covered in the lecture series.
• Students should independently study chapter 5 after the lecture.
• Chapter 4 is prerequisite knowledge.
• Chapters 1-3 will be covered later.

Lecture 1: Thermodynamics

• Lecture topics: Energy, entropy, enthalpy, and work.

Lecture 1.2: Interactions

• Topics include: Covalent bond, hydrogen bond, ionic bond, polarity, van der Waals interactions, and hydrophobic effect.

Atoms

• Atoms are composed of a positively charged nucleus surrounded by orbiting electrons.
• Bohr model (outdated) and orbitals are depicted.

Electronegativity

• Some atoms have a strong tendency to attract electrons (high electronegativity).
• Other atoms weakly attract electrons and tend to lose them (low electronegativity).
• Electronegativity correlates with position in the periodic table.
• A periodic table of electronegativity values is shown.

Electronegativity of elements

• Noble elements (inert) are stable.
• Other elements seek to achieve stable noble configurations.
• Elements on the left side of the periodic table generally lose electrons.
• Elements on the right side of the periodic table generally gain electrons.
• Most elements do neither; instead they share electrons

Covalent Bond

• Two atoms share an electron pair.
• This sharing requires atoms with similar electronegativities.

Carbon

• Carbon is a crucial element in biochemistry.
• Most biochemical compounds are built around a carbon backbone.
• Carbon's four valence electrons allow it to form four covalent bonds.
• This capability creates branched or ring-shaped structures with side chains.

C-C bond (Atomic Force Microscopy)

• Atomic Force Microscopy (AFM) can visualize covalent bonds, including between carbon atoms.
• Examples diagrams of covalent bonds are shown

Ionic Bond

• An ionic bond forms when a large difference in electronegativity exists between interacting atoms.
• One electron transfers from one atom to another.
• This results in a positive and a negative ion.
• Ionic compounds in water completely dissociate into their ions.
• Example: NaCl (sodium chloride or table salt)

Example: NaF

• Na donates an electron to F, producing ions (Na+ and F-)
• These ions create a salt crystal structure.
• In water, NaF will dissociate completely into ions

Ionic interactions (salt bridges)

• Oppositely charged ions strongly attract one another.
• Interactions of this type keep salts solid and are widespread in biomolecules.
• Charged atoms and groups are present in DNA, RNA, phospholipids, and proteins.
• Interactions of this type are called "electrostatic" due to permanency of charges.
• Charges can attract (+,-) or repel (++,--) each other

Polarity

• When atoms with intermediate electronegativity differences interact, they form covalent bonds.
• Electrons in the bond are unevenly distributed, creating a polar molecule.
• A partial charge separation occurs in the molecule.

Polarity in molecules

• A polar bond does not guarantee a polar molecule.
• For molecules like Carbon Dioxide (CO2), the polarity of the individual bonds cancel out
• Polarity requires partial charge separation due to asymmetry

Water

• Water is a polar molecule.
• The oxygen atom carries a partial negative charge.
• The hydrogen atoms carry a partial positive charge.

Hydrogen bond

• Partially positive hydrogen atoms in water are attracted to partially negative oxygen atoms in neighboring water molecules.
• Hydrogen bonding creates networks of water molecules.
• H-bonds also form between different molecules containing N-H, O-H, or F-H bonds, because these elements are highly electronegative.

Hydrogen bond (cont.)

• O-H and N-H groups are common in biomolecules
• These groups can form intermolecular H-bonds with other water molecules or other biomolecules.
• Intramolecular H-bonds can occur within a single molecule.

Strength of a hydrogen bond

• Strength depends on the distance and angle between interacting atoms or molecules.
• Short, straight bonds are strongest
• Typical enthalpies of binding depend on specific group to group interactions

Non-hydrogen dipole interactions

• Dipoles (partial charges) exert attractive or repulsive forces on other dipoles depending on orientation
• Dipole interactions are common in molecules like the C=O group.

Induced dipoles

• An ion or dipole can induce a dipole in a neighboring group, even if the neighboring group is usually nonpolar.
• Similarities in electronegativity value prevent strong bonds, such as the C-H bond.

Induced dipoles (cont)

• Atoms or groups that are normally nonpolar become polar due to neighboring dipoles or ions.
• The resulting interactions are called "induced dipoles"
• These induced dipoles create attractive interactions, referred to as "Debye forces" and are weaker than permanent dipoles

Induced dipole (cont)

• Induction of dipoles, in larger molecules with delocalized electrons is readily achieved
• Often seen in ring structures

Two induced dipoles?

• London dispersion forces are the weakest attractive forces between two nonpolar groups .
• These arise from the transient shifting of electrons, which creates temporary dipoles in the neighboring molecules.

Van der Waals interactions

• This Dutch Nobel laureate demonstrated that non-ionic, non-covalent interactions have a distance dependence.
• Atoms repel at very close distances, but attract at intermediary distances.

Van der Waals interactions (cont.)

• van der Waals forces are a collection of weak, distance-dependent attractive forces between atoms and molecules.
• Covalent and ionic bonds are not considered Van der Waals forces,

Strength of Interactions

• A table summarizing bond type, dissociation energies in kcal/mol and kJ/mol.

Hydrophobic effect

• The hydrophobic effect is not an interaction.
• It is the lack of interaction between non-polar molecules and polar molecules (water).

Hydrophobic effect (cont.)

• Polar and nonpolar molecules do not mix; nonpolar molecules are repelled by water.
• The hydrophobic effect occurs due to the greater entropy of bulk water when non-polar molecules are expelled from the water.
• This increased freedom of water molecule movement increases entropy.

Hydrophobic effect (cont.)

• Transferring non-polar molecules from water to a non-polar solvent results in a favorable change in Gibbs Free energy (∆G).

This quiz focuses on the key concepts of thermodynamics discussed in the Biochemistry lecture series, specifically extending Chapter 5 of the textbook. Students will explore energy, entropy, enthalpy, and various interactions such as covalent and hydrogen bonds. Mastery of these topics builds a foundation for further biochemistry studies.